inorganic chem

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I NORGC?N I C

CHEMISTRY

LABORATORY

EXPERIMENTS

WILS BEHGSTHOM, M.S.

and

MCIRLYS HOWELLS, Ph.D.

St. Paul Technical Institute

235 Marshall Avenue

St. Paul, MN 55102

Produced by Independent School District # 625 with the Minnesota Technical Assistance Program, University of Minnesota and the Environmental Protection Agency

PREFACE

In 1966, with the help of local industry, the Chemical Technology Program at St. Paul Technical Institute was developed to train qualified technicians to work in industrial chemistry laboratories. The emphasis of the training became chemistry principle applications in a laboratory situtation to reinforce chemistry theory.

laboratory, the students are required to perform a wide variety of experiments and are trained to safely handle many chemicals. These experiences prepare them for the job related activities they will encounter.

imperative that we begin training the students in chemical awareness and their right to know of the hazards related to materials they encounter on the job and in their training. Also, the cost of disposal of waste chemicals and storage requirements make it necessary for each training facility to evaluate their situation with regard of maintenance cost their program. This manual was developed to reduce the amount of hazardous materials and unnecessary exposure of the students and staff while still allowing for the maximum learning experience. This was done by microscaling quantities of chemicals used and by substituting other reagents where possible.

and evaluated during the year of writing, this, unfortunately, was not the case. There are a number of experiments that did not get a complete appraisal. Also, some of the tested laboratory experiments did not get a sufficient number of trials to verify their reliabilty at this time. We will continue to edit and revise over the next year and will be grateful for any suggestions that will aid in the producing of a highly useful manual.

With a special emphasis on the analytical aspects in the

-- . The OSHA requirements in the industrial setting made it

Although it was intended that all of the experiments be tested

We would like to express our appreciation to Cindy McComas for her help, guidance and support during the past year.

We would like to thank M. Renee Aanson for her help in editing and formatting of the manual this past year.

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TABLE OF CONTENTS

Laboratory Safety Practices

Analytical Balance Weighing Instructions Basic Laboratory Techniques

Good Laboratory Work Habits

Locker Check-In, Laboratory Safety, Laboratory Note- book and Data Recording, Metric, English and SI System of Measurement, Use of Laboratory Balances, Use of Bunsen Burner, Requirements for Report

Procedure (Solid, Liquid), Requirements for Report

Procedure, Requirements for Report


Procedure (Moisture Content, VCM, Ash & Fixed Carbon) Requirements for Report

Percentage of Water in a Hydrate Procedure, Requirements for Report

Empirical Formula of a Compound Procedure, Requirements for Report

Percentage of Potassium Chlorate in a Mixture Procedure, Requirements for Report

Inorganic Nomenclature Oxidation Number, Nomenclature for Binary Compounds, Nomenclature of Compounds Containing Polyatomic Anions, Nomenclature of Acid Salts, Writing Formulas Table of Metallic Elements and Polyatomic Ions Table of Nonmetallic Elements and Polyatomic Ions

Valence-Bond Theory, Valence Shell Electron Pair Repulsion Theory, Polarity Table of Likeness of Valence Bond to VSEPR Theory







Density Determination

Study Assignment for Symbols, Formulas and Equations Physical Separation of Mixtures

Paper Chromatography

r- Charcoal Analysis

Molecular Geometry

Limiting Reactant

Percent of Oxygen in Air

Gram Molecular Weight of Carbon Dioxide

Equivalent Weight of an Unknown Metal

Relationship of Temperature and Pressure to Change in

viscosity

Physical State

1

6 8

12

15 19

22

25

27

29

32

35

42 43 46

50 52

55

57

6 0

63

67

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TABLE OF CONTENTS

Surface Tension 70 Procedure, Requirements for Report












Calorimeter Constant, Heat of Neutralization, Heat of Solution, Procedure, Requirements for Report

Factors that affect reaction rates, Procedure, Requirements for Report.







Strengths of Electrolytes 73

Standardization of a Basic Solution 77

Vinegar Analysis 82

Acidity of Fruit Juices 84

c- - Standardization of an Acid Solution 87

Antacid Analysis 91

Analysis of Active Ingredient in Commercial Tablets 94

Iron - Copper Sulfate Reaction 97

Redox Titration with Potassium Permanganate 99

Colloids 103

Colligative Properties 106

Calorimetry 111

Rates of Reaction 117

Determination of a Rate Constant 120

Determination of the Order of a Rate Equation 123

Determination by Spectrophotometric Study the Rate of Bromination of Acetone 12s

Equilibrium-LeChatelier's Principle 128

Spectrophotometric Determination of an Equilibrium Constant 130

Determination of the Ksp of a Slightly Soluble Salt 133

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TABLE OF CONTENTS

Heterogeneous Equilibrium - Determination of the Ksp of a

Hydrolysis

Dissociation Constants

Determination of pKa of an Organic Acid

Buffers

Electrochemistry - Voltaic Cell - Nernst Equation Electrolytic Cell - Determination of Avogadro’s Number Electrolytic Cell - Coulometric Production of Cupric Ion Isolation of Caffeine from a Soft Drink

Preparation of Aspirin

Slightly Soluble Silver Salt Procedure, Requirements for Report


Requirements for Report








r-

135

137

139

141

142

144

14 6

148

150

152

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1

LABORATORY SAFETY PRACTICES

The laboratory is a safe working place when precautions and proper techniques are employed. Your safety practices are as important to your employer or your school as they are to you they are to you since your personal safety, the safety of your fellow workers/students, and the protection of property and equipment are important to them. Most precautions are just common-sense practices. These include the following:

1.

2 .

3.

- * - - 4 .

5 .

6 .

7 .

8 .

9 .

10.

11.

12.

13.

Wear safety glasses at all times while in the laboratory. It is the law in this state.

Wear shoes at all times.

Eating, drinking and smoking are strictly prohibited in the laboratory at all times.

Know where to find and how to use safety and first aid equipment.

Consider all chemicals to be hazardous unless you are instructed otherwise.

If chemicals come in contact with your skin or eyes, wash immediately with large amounts of water and then report it to your laboratory instructor.

Never taste anything. Never directly smell the source of any vapor or gas; instead, by means of your cupped hand, bring a small sample to your nose.

Any reactions involving skin irritating or hazardous chemicals, or unpleasant odors, are to be performed in the fume hood.

Never point a test tube that you are heating at yourself or your neighbor since it may splatter out of the opening.

No unauthorized experiments are to be performed in the laboratory.

Clean up all broken glassware immediately.

Always pour acid into water, NOT water into acid, because the heat of solution will cause the water to boil and the acid to splatter. For the same reason, always pour concentrated reagent solutions into diluted reagent solutions.

Avoid rubbing your eyes unless you are sure your hands are clean.

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14. When inserting glass tubing or thermometers into rubber stoppers, lubricate the tubing and the hole in the stopper with glycerin. Wrap the tubing in a towel and grasp the tubing as close to the end being inserted as possible. Slide the tubing into the rubber stopper with a twisting motion. DO NOT PUSH. Keep your hands as close together as possible to reduce the leverage. Finally, remove excess lubricant by wiping with the towel.

15. Notify your instructor immediately in case of an accident.

16. Many common organic reagents such as alcohols, acetone and ether are highly flammable. Do not use them any where near an open flame.

17. Comply with all special precautions mentioned in your experiments and with all special directions emphasized by your instructor.

18. Never use any chemical found in an unlabelled container.

1 9 . Read the reagent bottle label twice to be certain that it is the P- - chemical you want. The label of the reagent will list content’s

purity and safety hazards. If there is no indication of safety hazards of the chemical, treat it as though it is flammable, volatile, and poisonous until you have check the Material Safety Data Sheet, MSDS, for the chemical.

2 0 . NEVER WORK ALONE IN THE LABORATORY.

GOOD LABORATORY WORK HABITS

1. Read the assignment before coming to the laboratory.

2 . Unless instructed to do otherwise, work independently.

3. Record your results directly into your notebook. DO NOT recopy from another piece of paper.

4 . Do not clutter your working area with excess chemicals and/or equipment to avoid accidents.

5. Excess liquid reagents and solutions should be disposed of by pouring them in the sink and washing them away with large amounts of tap water, unless the experiment and instructor outlines a special disposal method.

6. Place excess solids in designated waste container. Never return excess reagents to the original b o t t l e .

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7. Do not place reagent-bottle stoppers and/or caps on the benchtop; hold them in your hand. Your instructor will demonstrate this technique. Replace the stopper on the same bottle, never a different one.

8. Leave the reagent bottles on the shelf where you found them.

9. Use only the amount of reagent recommended in the experiment

10. Always use distilled water in these experiments.

11. Do not borrow apparatus for other lockers. If you need extra equipment, obtain it from your instructor.

12. When weighing, do not place chemicals directly on the balance pan.

13. Do not weigh hot or warm objects. Objects should be at room temperture.

14. DO NOT place hot objects on the benchtop. Place them on a wire gauze.

15. When finished with your experiment, return all equipment to its proper storage place and clean all glassware with detergent and tap water. The glassware should be rinsed several times with sm-all quantites of distilled water and returned to your locker or appropriate storage place.

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--

DOUBLE BLAM SPATULA

SCOOP SrYLE SPATULA t

BUCWER FWNEL

WATCH G U S S

FORCEPS

5

RING STAH) W I T H RING

Analytical Balance Weighing Instructions

I mition8 nf the

. l fib 0 = Amrt parition

Plrtiak.(wr mition

1 = F u l l - n l r r Position

0 Woight-contml knob: 0.1 f

@ Woight-contml knob: 1 g

@ Woight-contml knob: 10 g

@ Roadoutpano1

@ T i n knobt

@ Irordjustmontknob

@ Digit8lcountor knob4

@ ~ o i g ~ ~ i n g p a n

@ hstment twor

*Only on Modrls H20 and H1E.

tA l l &Is except HE.

SOnb on Modrls H10 and H20.

1. lkton mighing, k w n thrt: I The balance is rmsted. The arrestment kvar

A There is nothing on tho pan, thc pan is cban, and

e. The wnightoontrol knobs, digitalcounter knob,

should k in a vertical position (at 0).

the wbighingchamber doom are closed.

and t8m knob am st at zero.

II. Setting tho olro point:

The zero point must be rechecked before each wbigh- i ng.

a. Tum the arrestment lever to the kft to the full- release position (at I). -

FI8.1- nl" EII0.d.

A When the numbed rok stops moving, turn the zem-adjustmrvt lamb until the "00" line of the numbed sak is p.rfsctly centered in the indi- ator slit (or on Models H8 and H18 until the "00" line of the numbend Kale is perfectly aligned with the zero line of the vemier Kale).

c Amst the balance by tuming the arrestment lever to 0.

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111. Waighing tlw mmpk:

a. Place sample on pan with forceps. Use a container or weighing paper.

b. Close the mighingchamber doom. c Tum the arrestment b a r to the partialnleasa posi-

tion (see diagram).

d. Tum the l0-g weightcontrol knob clockwise until the filling guida a m chang, direction. When th. a m change direction, dial back one step.

c Repeat wi th th. nmaining mightcontrol knobs.

1. Turn the arrestment lever to the arrest position and after a slight pause, tum the lever to the fullnlease position.

g, If your balance model has a vemier readout (H8 and H18), rwd the result when the numbered scale stops mi ng.

If your balance model has a digital readout (H10 and H201, tum the digitalcounter knob counterclockwise (after the numbered scale stops moving) until th. next lower scale division line is perfectly centend in the indicator slit or aligned with t h pointer.

IV. Comploting tha woighing

a. Record the result

b. A m s t the balance.

c R e m sample from pan with forceps.

d. Retum all knobs to zero. a. Clean the weighing chamber.

BALANCE READOUTS for Modo18 Hb, H10, HlO, H2O

M E T T L E R INSTRUMB NT C O R P O R A T I O N 20 Nassau Stmet, Princeton, New Jarsoy OM40

Reiz in t 2 w i n i s s i o s irznted by id .25t ler in 1985

BASIC LABORATORY TECHNIQUES

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Locker Check In

A student will be assigned a drawer and locker the first laboratory period and will be given an inventory list of glassware and equipment which should be found in the drawer. Refer to the preceding illustrated pages to identify items which are unfamiliar to you. If all items are not clean and ready for use, wash with soap and water, rinse well with tap water, and finally rinse several times with distilled water before storing them in the locker. Check with the instructor for replacement of those items which are not found in the locker.

Laboratory Safety

-* - - Your laboratory instructor will discuss safety and laboratory

procedures with you. A safety film will be shown on common hazards and accident situations found in a laboratory. A required reading assignment in the Chemical Technician's Ready Reference Handbook of is pages 13 - 36. The instructor will request that you sign a sheet of standard laboratory and safety procedures signifying that you understand and will abide by these guidelines. The instructor will indicate the location of fire and safety equipment in laboratory and demonstrate their operations. Students will become familiar with Right To Know Legislation and Material Safety Data Sheets.

Laboratory Notebook and Data Recording

Students will be required to maintain a laboratory notebook of all laboratory experiences. Required reading assignment in Chemical Technician's Ready Reference Handbook (CTRRH) is pages 6 - 13. The organization of the notebook will be left to the individual, however, the instructor will make initial suggestions.

The notebook should be a bound book, not loose leaf pages. The first page should show your name, the chemistry course name and number, and the date on which the course started. This should be followed on the next page by a "Table of Contents" which will list each experiment and the page on which it begins. Depending on the size of the notebook it may be appropriate to leave several pages before starting the first experiment entry. be described separately. Use the following form when it applies:

Each experiment should

1. Page number 2. Title of experiment 3. Object of experiment 4 . Equations of reactions (if any) 5. Table listing principal reactants and products, their formulas,

and important physical properties and hazardous properties from the Material Safety Data Sheet available in the laboratory

6. Procedure and sketch of experimental apparatus 74. Data and observations 8. Calculations (if any) 9. Conclusions (and interpretations where they are appropriate).

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For each experiment items 1 - 6 in the format outline should be in the notebook before work is started in the laboratory. Data tables can also be organized prior to starting if you k n o w how the information is to be collected. Write only on the right hand page of the notebook. Do not write information on odd Pieces of Pap er to C O P Y into your no tebook la ter. Keep your writing legible, accurate, neat, and clean. Use ink, preferably black, so that the record is permanent. Ball point pens are better since they are less likely to run if liquid is spilled on the notebook. Never erase or use correction fluids to eliminate mistakes. Instead, draw a line through unwanted entries. If notebook entries for one experiment are continued on some page other than adjacent page in the book, indicate which page to turn to and, at the new page, indicate the number of the page from which it is continued. Include all the information necessary for another person to duplicate your experiment, however, be as brief and concised as possible. Enter data and observations exactly. The essence of a notebook is accuracy. If an experiment is ended by an accidental spill or breakage, indicate this and start it over again. The notebook must be a permanent record of what you did and what happened.

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e The student will construct a table of metric and SI units and

conversion factors for physical quantities of length, volume, mass, pressure, temperature, and energy. This information may be obtained from a textbook or other suitable reference. (In the CTRRH this information can be found on pages 787 -794.) The student will construct a second table containing the common metric prefixes, their abbreviations, and their designated power of ten.

The following is an example problem of how the conversion information can be used. A 145'1b student, weighs how many kilograms? Student must recognize that a pound, lb, to kilogram, kg, is necessary to make the conversion. Since lb is the measured unit the lb must appear

145 l b X 1 ker = 6 5 . 7 5 9 6 kg = 65.8 kg 2 . 2 0 5 lbs

in the denominator of the conversion factor for lb units to cancel. Another way to preceive the problem is that the unit you wish to convert to must be in the numerator of the factor. Please note the answer is rounded off to contain the same number of significant digits as contained in the original measurement. When just one conversion factor is not available to make the conversion from the one system to another, a series of factors are used to make the conversion. For example, 10 gallons of gasoline are equivalent to how many milliliters? The gallons are first converted to qts, qts

10 gal X 4.0 ats X 1 liter X 1000 ml = 3 7 , 8 4 2 ml 1 gal 1 . 0 5 7 qts 1 liter 38,000 ml

to liters, liters to ml. Each time the unit which must cancel is placed in the denominator of the factor. Notice the answer is rounded off to contain the same number of significant digits as the measured unit.

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The following problems are required to be completed and included in the laboratory report submitted by the student.

1. 2 . 3 . 4 . 5 . 6 . 7 .

9 . a .

Convert Convert Convert Convert Convert Convert Convert Convert Convert

175 lbs to kg. 16 fluid ounces to ml. 800 m to miles. 4.00 atm to torr. 4 5 . 3 kcal to joules. 4 5 2 ml to kl. 267 kg to mg 3/8 inch to mm 65 miles/hour to km/hour

Remember to round off answers to the appropriate number of significant digits.

The student will use two types of balances when measuring mass -. in the laboratory. The triple beam balance is capable of measuring

to the nearest hundredth of a gram, 0.01 g, and is normally employed when approximate weighings are required for a procedure. The analytical balance is capable of measuring to the nearest milligram or tenth of a milligram] 0.001 g or 0.0001 g , and is employed when that level of accuracy is necessary in the procedure. The instructor will demonstrate a weighing method on both types of balances and the student is required to read pages 543 - 555 in CTRRH which show the various models and outlines general procedure of use.

Students will be required to weigh and record weights of the following items on both types of balances.

1. 50 ml beaker 2. evaporating dish 3 . 125 ml flask 4 . a small test tube

It is recommended that the student weigh these items on the triple beam balance first to make the weight determination on the analytical balance more efficiently. When the student is confident about the operation of the analytical balance, he/she should obtain a preweighed, numbered metal tag from the instructor and demonstrate the weighing process. DO NOT handle the tag with your fingers, use your forceps since fingers will deposit oils and moisture on the tag and change its weight. Whenever using an analytical balance, it is recommended that the appropriate tool, tongs or holder, be use to avoid touching the item. This minimizes the error in the weighing process. Record the tag’s weight and check your results with the instructor.

Use of Bunsen Bur ney

The laboratory burner is one of the most often used pieces of equipment in your locker. Learning its proper adjustment procedure is very important. The natural gas used in most laboratories is composed primarily of methane, C H 4 . If sufficient oxygen is

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supplied, methane burns with a blue, nonluminous flame producing carbon dioxide and water. With insufficient oxygen, carbon particles are produced which, heated to incandescence, produce a yellow, luminous flame.

Attach the burner’s tubing to the gas outlet on the l a b bench. Turn off the burner’s gas control valve, located on t h e base of the burner, by screwing the valve in towards the base. Turn on fully the gas outlet on the lab bench. Close the air holes located at the base of the barrel, and open the burner’s gas valve slightly. Bring a lighted match or striker up the outside of the barrel until the escaping gas ignites. After it ignites, adjust the gas valve until the flame is pale blue and has two or more distinct cones. Opening the air control holes produces a slight buzzing sound characteristic of a burner’s hottest flame. The addition of too much air may extinguish the flame. When the best adjustment is reached, three distinct cones are visible: a bright blue inner, intermediate luminous and a violet outer area. The intermediate luminous area represents the hottest part of the flame and this may be demonstrated

to the barrel. Refer to CTRRH for illustration of burner diagram. c* . . by using crucible tongs to hold a wire screen in the flame parallel


At the conclusion of each experiment, the student will be required to summarize the information observed and collected in report form. A format similar to the notebook outline is required. The following information must be included:

1. Name of student and date 2. Title of experiment 3 . Object of experiment 4 . Equations of reactions 5 . Data tables and observations 6. Sample calculations and results table 7 . Conclusions

Answer all the questions listed at the end of each experiment. This information is submitted to the instructor for grading in much the same way a report would be submitted to a supervisor in an industrial laboratory. The instructor may request to look at your laboratory notebook periodically to verify the accuracy of your records and adherence to the notebook guidelines.

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DENSITY DETERMINATION

All pure substances have characteristic properties which can be measured and used to identify them. A common example of this is the boiling point of water is 1 O O O C . However, many other pure substances may have the same boiling point as water so other physical properties must be measured or observed to identify a clear, colorless liquid as water. One property which is very unique to pure substances is its density. Density is determined by measuring the mass of a substance and dividing by the volume that the substance

Density = the mass of the substance volume of the substance

occupies. Common units for mass are grams or pounds. Common units for volume are cubic centimeters (cm3), milliliters (ml), or cubic feet (ft3). The density of water at 40C is 1.000 g/cm3 which is equal to 1.000 g/ml; therefore] the units milliliters and cubic centimeters are usually interchangeable. In the US standard system,

-- the density of water at 4 o C is 62.4 lb/fts. Since scientists prefer easy conversion from metric system to US system, they developed specific gravity which is a ratio of the substance's density divided by

Specific Gravity = the density of the subst- the density of water at 40C

density of water at 4oC. The specific gravity is a unitless number, but in the metric system it has the same numerical value as the density. Specific gravity is the value reported for a substance] regardless of which system of measurement is used to determine its value. The density of iron is 7.86 g/cm3 or its specific gravity is 7.86. Using the specific gravity times the density of water in the US

7.86 X 62.4 lb/ft3 = 793 lb/ftj

system, the density of iron can be calculated. Engineers and manufacturers generally work with the SI system of measurement. This is a new version of the metric system adopted by International Union of Pure and Applied Chemistry] IUPAC, to provide a more logical mattrix for all basic measurements globally. In this experiment] you will determine the densities of a solid and a liquid.

The volume of a solid can be measured by two methods:

1) measure with a metric ruler relevant dimensions necessary for appropriate mathematical volume formula based on the shape of object;

graduated cylinder. 2) measure the volume of the object by water displacement in a

The volume of a liquid can be determined by measuring the volume of liquid in a graduated cylinder or by using a graduated pipet to deliver a measured volume to a preweighed flask.

The mass of both solid and liquid must be determined. The student is required to read pages 193 - 200 in the CTRRH so that he/she will become familiar with other methods of density determination.

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Procedure

A. Solid

Obtain a numbered metal cylinder from the instructor and determine its mass to the nearest 0.01 g. Using a small metric ruler measure the diameter and length of the cylinder. A cylinder’s volume formula is 3.14 (pi) x radius ( 1 / 2 of the diameter) squared x length. Fill a 50 or 100 ml graduated cylinder with water to a convenient level so that the volume may be determined. Since water adheres to the glass wall of the cylinder, the water level will not appear flat. This is a meniscus. To properly read the volume, use a small, white piece of paper with a wide dark mark. Place the paper behind the cylinder well below the top of the water and slowly move up the barrel of the cylinder until the transparent meniscus appears to darken. At all times read the volume using the meniscus for clear liquids. Volume readings should always be taken at eye level. This technique is illustrated on pages 581 - 582 in CTRRH. Record the volume of water. Carefully add the metal cylinder to the water and measure the new volume as previously described. Pay close attention to the accuracy with which the volumes can be determined by both methods and calculate the densities with the appropriate number of significant digits.

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B. Liquid

Obtain unknown liquid from instructor in a 100 ml beaker. Weigh an empty, clean, 10 ml graduated cylinder to the nearest 0.01 g. From the beaker pour liquid into graduated cylinder with the aid of your stirring rod just short of the 10 ml mark (page 71 in CTRRH). This is accomplished by placing the rod over the lip of the beaker so that it extends 2 inches or more beyond the lip. The rod is held in place by exerting pressure with index finger while holding the beaker. To transfer liquid touch the rod to the side of the graduated cylinder and slowly pour. Use medicine dropper to reach the 10 ml mark. Carefully weigh the cylinder to the nearest 0.01 g and record. When measurement has been completed, liquid in the graduated cylinder may be disposed of in the lab sink with running tap water. Calculate the density with the appropriate significant digits.

The remaining liquid in the beaker will be used for a second trial using a graduated pipet to measure volume and the analytical balance to measure mass. Weigh a clean, dry stoppered flask to the nearest 0.001 g. Use the rubber bulb to draw some of the liquid into a clean, dry pipet. Remove the bulb and use your index finger to close the opening of the pipet. Hold the pipet horizontally and slowly roll the liquid in the pipet to wet the inside surface. Allow liquid to drain from the pipet into the lab sink. Examine the inside for beads of liquid and if found clean the pipet thoroughly. It is a necessity that a pipet deliver an exact volume and if small amounts of liquid adhere to the inside, that is volume which is not measured.

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Pipets should always drain cleanly. Rinse the inside of the pipet with another small quantity of liquid and allow it to drain. Draw the unknown liquid into to the pipet above the 10 ml mark. Using the pressure of the index finger on the opening slowly bring the meniscus to the 10 ml mark. Read the volume at eye level. Do not worry if the meniscus is not exactly at the ten ml mark since the pipet is graduated it will allow an exact volume reading. Move the pipet to the preweighed flask and allow the liquid to drain into flask by lifting your index finger. As the level of draining liquid approaches the zero line on the pipet, place index finger over opening to control flow rate and eventually stop it. Remember that flow must be stopped above the zero line since the area from zero to the tip is not measureable. Read the volume at eye level and record. Stopper your flask and measure the mass to the nearest 0.001 g. Calculate the density of unknown liquid. Dispose of excess unknown liquid in lab sink with large amounts of tap water.


Follow the guidelines established in Basic Laboratory Techniques -._ section for writing your report. In the conclusion section, make

comparisons between the two different methods used for density determinations relative to their accuracy and which measurement determined the number of significant digits in the calculated density. Report average density of both solid and liquid with the appropriate number of significant digits.

1. If you had an irregularly shaped solid, how could you determine its volume? Explain.

2 . If air bubbles adhere to the surface of the metal cylinder when submerged for measurement, how does this affect the density determination? Explain.

3. If several drops of unknown liquid remain on the pipet’s inner wall, how will this affect the reported density? Explain.

4 . If draw a sketch of a cross section fo a graduated cylinder illustrating the proper method for reading volume.

5 . Using a suitable reference source, briefly describe two other methods by which density may be determined.

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STUDY ASSIGNMENT FOR SYMBOLS, FORMULAS, AND EQUATIONS

The chemist communicates in a technical language that is specific to the substances with which he/she deals. The language is a notation abbreviated to save time and space similar to the shorthand notations used by stenographers. However, if one becomes familiar with these terms and notations, the language is readily understandable. A chemist's basic language is symbols for elements, formulas for compounds, and equations for reactions.

In this assignment, students will study each of these and learn to recognize their significance. Students will be supplied with a periodic table of elements, and will use the CRC Handbook of Chemistry and Physics or textbook to find necessary information to complete the exercises.

A symbol represents the name and also one atom of an element. Since Latin and Greek were the languages used by early chemists, some of the elements were given Latin names, whereas the more recently discovered elements have English names. For example, Au represents aurum (gold), Cu represents cuprum (copper), Fe represents ferrum (iron), but Co represents cobalt and Se, selenium. The Periodic Table lists the elements, their atomic numbers (the whole number), and atomic weights. The atomic weights are the average masses of all naturally occurring isotopes of the element and, consequently, many of the elements do not have whole numbers for their atomic weights. Those elements having whole number masses have been synthetically prepared. Elements to the right of the "stepped" line beginning at boron are nonmetals and to the left are metals.

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A formula is a combination of symbols representing the name and also one molecule or formula unit of a compound, a chemical combination of elements. Subscripts indicate the number of atoms of each element in a compound. For example, H2O is the formula for water, a chemical combination of 2 atoms of hydrogen, H, and 1 atom of oxygen, 0; COz is the formula for carbon dioxide, a chemical combination of 1 carbon atom and two oxygen atoms whereas CO is the formula for carbon monoxide, a chemical combination of 1 carbon atom and 1 oxygen atom. A change in subscript represents a change in chemical composition of a compound.

A chemical equation represents a description of a chemical reaction before and after it takes place. It uses symbols and formulas to describe chemical substances. The reactants appear on the left side of the arrow, meaning "to produce", and products appear on the right side. For example, the reaction between

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carbon and oxygen to produce carbon dioxide simplifies to

c + 0 2 --- > c02.

The plus sign, +, means "and", but does not imply a summation. The chemical equation implys that one carbon atom and one oxygen molecule react to produce one molecule of carbon dioxide.

A chemical equation not only simplifies a statement, but it also indicates the number of atoms involved in a reaction and how they rearrange to form products. In the above equation, note that 1 C atom and 2 0 atoms are rearranged to form a product containing 1 C atom, 2 0 atoms. We say that atoms are conserved in a chemical reaction; they are only rearranged. If each atom has its own mass, then mass also is conserved.

A correct equation must have the same number of each type atom appearing on both sides of it. To achieve this mass balance, coefficients are placed before the symbols or formulas in the equation. For example, potassium metal and chlorine gas

-. . . react to produce potassium chloride:

K ( s ) + C 1 2 ( 6 ) --- > K C 1 ( s ) (unbalanced)

The symbols for potassium and chlorine and the formula for potassium chloride are correctly written. To indicate 2 atoms of-chlorine in the product, a coefficient of 2 is placed before K C 1 , or

K ( s ) + C l 2 (g) - - - > 2 K C 1 ( S I (unbalanced)

However, the equation now shows 2 potassium atoms on the right; a coefficient of 2 in front of the K on the left balances the equation:

2 K Cs) + C l 2 (g) --- > 2 K C 1 (s)

The equation is now mass balanced or obeys the Law of Conservation of Mass.

Section A

Use the Periodic Table (Sargent-Welch) supplied by the instructor to answer the following questions.

1. Identify the elements'on the Periodic Table which are gases and liquids. Construct a table containing the symbol of the element, its name, its state at room temperature and atmospheric pressure, the smallest particle of the element that occurs naturally.

2. All naturally occurring gases, except the noble gas, are are diatomic. All halogens, Group VIIA, are diatomic. Write the formulas for these elements in their naturally occurring state.

17

3. Match the following names for the elements with their symbol. List the standard English name for the element.

Symbol Latin Name English Name Choices

Na K Fe cu Ag Sn Sb W Au

aurum argentum cuprum f errum hydragyrum kalium natrium plumbum stannum stibium wolf ram

4 . Of the 106 known elements, few make a significant contribution to the composition of the earth’s crust. Using the CRC Handbook, construct a table of the ten most abundant elements

symbols, atomic weights, and classify them as metal or nonmetal c- - in the earth’s crust, their percentages, atomic numbers,

Section B

1. Write the formulas for the following compounds which are encountered commonly in everyday living. Using the CRC Handbook or chemical dictionary, look up each compound’s physical state, color, and appearance.

a. table salt 1 Na atom , 1 C1 atom b. table sugar (sucrose) 12 C atoms, 22 H atoms, 11 0 atoms c. ammonia 1 N atom, 3 H atoms d . water 2 H atoms, 1 0 atom e. baking soda 1 Na atom, 1 H atom, 1 C atom, and

3 0 atoms f. limestone 1 Ca atom, 1 C atom, 3 0 atoms g. Vitamin C 6 C atoms, 8 H atoms, 6 0 atoms h. vinegar 2 C atoms, 4 H atoms, 2 0 atoms

2. The instructor will make available to you approximately 50 samples of compounds, labelled with name and formula, in large test tubes. These ionic compounds, normally formed by the reaction of a metal with a nonmetal, are generally called salts. Complete the following generalizations based on your observations.

a. Blue colored salts often contain b. Green colored salts often contain - c. Most salts exhibit a --- color. d. Most cobalt salts are in color. e. Salts containing MnO4- exhibit a

g. Salts containing (203-2 exhibit a

color. color. color

f. Salts containing Crz07-2 exhibit a -

but the exceptions are -- and __-_____--

18

Section C

1. Write a chemical equation for the following statements of reactions and balance the equation.

a. Methane, C H 4 , and o x y g e n to produce carbon dioxide and

b. Hydrogen and oxygen to produce water. c. Calcium carbonate, CaC03, and hydrochloric acid, HC1, to

water.

produce calcium chloride, CaC12, water, and carbon dioxide.

M g z N a . d. Magnesium and nitrogen to produce magnesium nitride,

e. Potassium and sulfur to produce potassium sulfide, K 2 S . f. Aluminum and hydrochloric acid to produce aluminum

chloride, A1C13 , and hydrogen.

19

PHYSICAL SEPARATION OF MIXTURES

Matter can be classified into one of two general categories, pure substances or mixtures. Pure substances are elements or chemical combinations of elements called compounds. The elements hydrogen and ‘oxygen can be chemically combined to form the compound water. Mixtures are physical combinations of two or more pure substances such that each substance maintains its own chemical identity. For example, in a salt-water mixture, each component retains the chemical properties as found in their pure state; the water is H20 molecules and the salt is sodium ions, Na+, and chloride ions, C1- .

Mixtures are either homogeneous or heterogeneous. Homogeneous mixtures consist of two or more substances forming one phase with uniform properties throughout. The salt-water mixture is an example and this is referred to as a solution. Air is an example of a homogeneous mixture of gases. Heterogeneous mixtures consist of two or more distinct phases or regions each having its characteristic properties. Examples of heterogeneous mixtures are concrete, composed of cement and stone, ice water, composed of solid water and liquid water, and Italian salad dressing, composed of oil, water and spices.

c- -

Most substances found in nature or prepared in the laboratory are impure, thus representing mixtures. The method chosen to separate these mixtures depends on the differences in the chemical and/or physical properties of the mixture’s components. Some of the common physical methods used to separate components are:

1.

2.

3

4 .

5.

6.

7 .

centrifugation - separating solids from liquids or liquids of different specific gravities by fast rotation chromatography - separating components on the basis of their individual selective adsorption on a stationary phase as the mixture passes through or over it crystallization - forming a crystalline solid from solution by decreasing its solubility distillation - purifying a liquid by boiling it and condensing and collecting the vapors extraction - removing a substance from a mixture by adding a a liquid in which the substance is more soluble filtration - separating a liquid from a solid using a porous material through which only the liquid will pass sublimation - heating a solid substance directly to vapor and recondesing it as a solid ( not all solids sublime)

In this experiment, gravity filtration will be used to separate a sand-salt mixture. Three techniques are associated with all filtration processes: decantation, washing, and transfer. When a solid settles to the container’s bottom in a liquid and does not appear to be suspended, it can be separated from the liquid by pouring off the liquid so that no solid is carried along. This technique is decantation. The solid is then washed with a small amount of water to remove any remaining soluble material, salt.

20

The solid is washings are with the aid The stirring

allowed to settle and liquid is then decanted. Several recommended. The transfer of the solid is accomplished of a wash bottle and stirring rod with a rubber policeman. rod is used in the same manner as liquid transfer

described previously. The beaker and rod are held at inclined position above the filter with the rod touching the filter. A water stream from the wash bottle is used to wash the solid into the filter See pages 266 - 272 in CTRRH.

Procedure

In a 250 ml beaker, weigh a 2 - 3 g sample of sand-salt mixture to the nearest 0.01 g. Obtain a ring stand and ring to support the funnel found in your locker. If funnel is too small to be supported by the ring, place your clay triangle on the ring and then insert the funnel in the triangle. Obtain a piece of filter paper and fold it precisely in half. Fold it in half again but do not make a perfect right angle(800). You should not have perfect quarters. Take the edge of the smaller section and tear off the corner. The tear enables a close seal to be made across the paper’s folded portion when placed in the funnel. Weigh the filter paper to the nearest 0.01 g . Place filter paper snugly in the funnel so that the edge with the tear is in contact with the funnel wall. Moisten the filter with distilled water and press it against the funnel’s wall to form a good seal.

Add 20 ml of water to the sand-salt mixture and stir with a glass rod for several minutes. Incline the beaker so that sand collects directly below the lip of the beaker. While the sand is settling, place a 100 ml beaker beneath the funnel so that the tip of the funnel is just touching the side of the beaker. Decant the liquid using proper techniques. Add 10 ml of water to the beaker and repeat the stirring, settling, and decanting processes. Add 10 ml of water to the beaker and repeat these three manipulations. Transfer the sand to the filter paper using the wash bottle as previously described. Allow the liquid to drain from the filter and place it on a marked watch g lass for drying in the drying oven. When the paper is dry, allow it to cool and weigh it again to determine the mass of the sand. Calculate the percentage of sand in the mixture by dividing the mass of sand by original sample mass and multiply by 100.

Remove the funnel from the ring and place a wire gauze on the ring to support a preweighed evaporating dish. Pour liquid from the filtering process into the evaporating dish and heat it with a burner. DO NOT attempt to overfill the evaporating dish. When all the liquid has been transferred to the dish, rinse 100 ml beaker several times with wash bottle. Heat the dish until the volume of water has been reduced to approximately 5 ml. Place in drying oven to complete the drying of the salt. Allow the dish to cool and weigh to determine the mass of salt. Calculate the percentage of salt in the mixture in the same manner used for the percentage of sand. Do the percentages total loo%? If not, give possible reasons for the discrepancy. Repeat the entire procedure for a second trial making any adjustments you think are necessary to improve your results.

2 1


Follow guidelines established in Basic Laboratory Techniques experiment for writing your report. Compare the first and second trial results and report average percentages of sand and salt.

Quest ions

1.

2 . 3. 4 .

-* - 5 .

Using a suitable reference source, define the following words: a) adsorption b ) phase c ) soluble d) solution e) solvent f ) suspension What substance is extracted in this experiment? Why is the sand washed with 10 ml portions of distilled water? If the washings in question 2 are not done, how would this affect the percentage of sand? Explain. Explain how you would separate a sugar-sand mixture. Would the procedure in this experiment have to be modified in any way?

22

PAPER CHROMATOGRAPHY

Physical methods used for identification depend upon the differences in physical properties of substances present in mixtures. Chromatography is the name applied to a series of processes ir which the components of a mixture are distributed between a station& y phase and a mobile phase. Normally, the stationary phase is a solid such as paper, starch, alumina, or silica and the mobile phase is a liquid such as water, common organic solvents, or solvent mixtures. The word chromatography, literally means color writing, and was coined by a Russian botanist, Twsett, when he separated color pigments into the individual dyes using this method in 1906. Color is not necessary to achieve separation by this method. Colorless compounds can be reacted with other reagents to make them visible.

Paper chromatography is the simplest and most efficient technique available to introduce this unique separation method. The basis for paper chromatography is the fact that porous paper, cellulose, has an enormous surface area to which molecules or ions of substances are attracted (adsorbed) and then released (desorbed) into the solvent as an aqueous solution passes over the paper. Separation of components occur, that is, they will travel at different speeds in a moving solvent because the varying attractions between these components and the paper. This phenomenon of separation is known as partitioning.

-. ~

In paper chromatography, a small line of the mixture to be separated is placed at one end of a strip of paper, and solvent is allowed to move up the paper through the line by capillary action. The solvent and the various components of the mixture travel at different speeds along the paper. The identity of the components can be deduced by comparing a chromatogram of the unknown mixture with chromatograms of mixtures with known composition (standards). An additional aid in identification of substances is its Rf value, which is defined as the ratio of the distance traveled by the substance to

R € = distance traveled by component distance traveled by solvent

the distance traveled by the solvent. The Rf value of a compound is a characteristic of the compcmnd and solvent used and serves to identify the constituents of a mixture. Your aim in this experiment will be to determine the number of substances and their Rt values present as dyes in commercial ink samples, such as Crayola marking pens. Any commercial ink sample can be separated into component colors provided the appropriate solvent is used.

23

Procedure

Fill a 400 ml beaker with water to a height of 1 cm. Obtain six strips of chromatographic paper. With a pencil, draw a light line about 1 cm from one end of the strip and parallel to the short dimension of the

~-

Sample

L ine B

Line A

strip (Line A ) . Draw a second parallel line, Line B, 1 cm from Line A . Line A will represent the level to which the strip should be submerged into the solvent, water. Using the marking pen, apply a thin ink line over Line B on the paper strip. This may require some practice on filter paper to ensure a thin straight line. Prepare separate paper strips for each marking pen color you w.sh to analyze. Attach a paper clip to the end of the strip opposite the lines and mark this end, using a pencil, with the name of the color marking pen. CAUTION: Do not handle the paper with your fingers if either they or the paper, or both are wet. Allow the ink color to air dry after application to the paper strip. Hang the paper on a stick and adjust paper clip so that strip will be submerged to Line A when the stick is placed across the top of the beaker. Place the strip in the beaker as described previously and observe what takes place. When the solvent front has almost reached the top of the paper strip, remove it from the beaker and mark the position with a pencil. This should be done quickly because some solvents evaporate rapidly. Allow the chromatogram to air dry. Measure the distance from Line B to each separated dye in the sample and record in your notebook. Measure the distance from Line B to the marked

solvent front and record in your notebook. Proceed in the same manner with the other strips that you prepared using the various colored marking pens. Save all chromatograms.

Repeat the procedure to separate the same inks into component colors with a 5 0 : 5 0 methanol-water solvent mixture. Repeat the procedure with methanol as the solvent. time remaining in the lab period, samples of other commercial inks may be done.

If the student has extra

Tabulate information found on the individual chromatograms. The following items should be listed in the table; sample, solvent, component’s distance traveled, solvent’s distance traveled, and component’s Rr value. ratio explained in the introduction to the experiment.

The last entry can be calculated using the

Beauiremonts f o r ReDort

Follow established guidelines for the report. Attach to the report, a representative chromatogram from each solvent system used. For a sample, discuss the similarities and/or differences of Rf values with respect to the three solvent systems used. Comment one components which seem to be the same in the samples.

24 Glues t ions

1. Distinguish between absorption and adsorption. 2. What is an eluting solution and how is it used in

chromatography? 3. Sketch what a chromatogram would look like if it had two

components, one with a Rf = 0.65 and another with a Rf = 0.80. 4 . What factors influence the Rf values? Would it be possible to

compare the Rf values for the dyes in the ink samples to those in the candy samples?

5. What forces cause the solvent to move along the paper?

25

CHARCOAL ANALYSIS

Fossil fuels represent a large portion of the world’s energy supply. They also serve as an economical source of raw materials for the manufacture of products such as plastics. synthetic fabrics, and medicines. A very important property of a fuel is its heat of combustion which is the amount of heat (calories, joules, or BTU’s) released per amount of fuel (barrels, cubic feet, or tons). The energy rating of a fossil fuel gives us information about its chemical composition. Although charcoal is not a fossil fuel, but amorphous carbon from the incomplete combustion of animal or vegetahie matter, its chemical composition can be analyzed in the same fashion as fossil fuels. The average person is more familiar with charcoal as the energy source for cooking on outdoor grills.

In this experiment, the ash, fixed carbon, moisture, and volatile combustible matter (VCM) for charcoal is determined. The analysis approximates charcoal’s composition because values vary

c- - between charcoal samples. A student wishing to make a BEST BUY comparison would be required to do a large number of samples from different bags of the same brand of charcoal. The collected information could then be compared to a corresponding number of samples from other brands.

Procedure

A . Moisture Content

Support a clean, empty crucible and lid on a clay triangle using a ring stand. Heat the crucible with a hot burner flame for several minutes to burn off any impurities. Handle the crucible and lid with crucible tongs for the remainder of the experiment. Move the crucible and lid to the carrying tray, allow to cool, and weigh to the nearest 0.001 g. Add approximately 1 g of charcoal to the crucible and weigh again. The lid may be weighed separately since its use is not required in the first section of the procedure. Put the crucible in a 100 ml beaker marked for easy identification and place in a drying oven for 1 hour. Allow crucible to cool and reweigh. The weight loss represents the moisture content of the charcoal.

B. Volatile Combustible Matter, VCM

Now support the crucib.le and dried charcoal in a clay triangle and cover with the lid leaving a slight opening at one side. Initially, heat the apparatus gently, gradually increasing to a hot flame for several minutes. Use tongs to close the lid completely and allow the sample to cool to room temperature. Before weighing the crucible, examine the upper portion of the crucible and the lid for deposits and discolorations. If deposits appear on the lid, heat it directly in the flame. If deposits appear on the crucible, carefully heat the upper portion of the crucible without the lid and avoid heating the charcoal sample. Weigh the crucible, lid and remaining sample, when cool, to the nearest 0.001 g. The weight loss is the VCM evolved during the heating process.

26

C. Ash and Fixed Carbon

Support the crucible in the clay triangle and partially cover it with the lid. Heat the sample intensely and if the sample ignites, use the tongs to cover the crucible until flame is extinguished. Continue intense heating until no black residue remains on the lid or in the crucible. This will require that you periodically use the tongs to rotate the crucible's position in the flame. The residue in the crucible should appear to be a light gray-brown color when the combustion has been completed. Cool the crucible, lid, and contents and weigh. The residue remaining in the crucible is the ash content. The weight l o s s after the VCM removal is the fixed carbon content. Repeat the entire procedure f o r a second trial.


Tabulate the data collected in the experiment and calculate the percentages of ash, fixed carbon, moisture and VCM from data. Pay attention to significant digits in the data and reported numbers. Comment on the consistency or inconsistency of values from the two trials.

Questions

1. Of the four determinations in the experiment, which is proportional to the heat content or would be equivalent to an energy rating?

2. Identify in each section of the procedure whether a physical or chemical change is being used to analyze the charcoal.

3 . For those physical changes you identified in the previous question, refer to the Physical Separations experiment to relate them to a particular separation process.

4 . If the moisture is not removed in the drying oven, how does this affect the reported VCM (high or low)? Explain.

5 . If the crucible's lid is not used in Section B of the procedure, what reported values are affected and how? Explain.

27

PERCENTAGE OF WATER IN A HYDRATE

Many pure substances combine with water in a fixed mole ratio to yield compounds called hydrates. For example, zinc sulfate combines with water to form crystalline ZnSO4-7HzO which is a stable compound at normal atmospheric conditions. All pure samples of this hydrate show the same percentage of water by analysis. Thus, this hydrated compound obeys the law of constant composition. Upon heating a sample of such a hydrate, it may lose all its water of hydration and revert to the anhydrous salt. Substances which have adsorbed water on the surface do not show constant composition and therefore are not hydrates. An example of this would be common table salt, NaC1, which becomes very sticky on humid, summer days. In these cases, the percentage of water is not constant for all samples of a particular compound, and the water is not chemically bonded as part of the crystal structure.

The purpose of this experiment is to determine the percentage of wate'r in an unknown hydrate. Water is removed from the hydrate by heating an accurately weighed hydrate sample until the residue has reached a constant weight. The percentage of water in the sample is calculated by using the weight of water lost and the initial hydrate sample weight multiplied by 100.

Procedure

Clean and dry a porcelain crucible and cover. Place the empty, covered crucible on a clay triangle supported by a ring on a ring stand. Heat the crucible and cover in the hottest flame of the Bunsen burner for 5 minutes. A dull red glow should be observed on the crucible and cover. This will require that you manipulate the burner to heat uniformly. This will insured that all volatiles and combustible materials are removed prior to the analysis procedure and that a constant weight for the crucible and cover may be recorded. The crucible and cover must be allowed to cool to laboratory temperature for approximately 15 minutes. Using crucible tongs transfer the crucible and cover to a carrying tray and weigh them to the nearest 0.001 g. Add 1 to 1.5 g of zinc sulfate heptahydrate to the crucible and weigh the covered crucible to the nearest 0.001g.

Place the covered crucible on the clay triangle with the cover slightly opened. Heat the crucible gently for a few minutes to avoid loss of material from spattering during initial heating. Remember

contents for approximately 15 minutes gradually increasing the temperature of the flame. Allow the crucible to cool on the triangle after removing the flame until it has reached room temperature. Transfer it to the carrying tray and weigh the covered crucible to the nearest 0.001 g. Reheat the crucible and contents for about 5 minutes and, after cooling, weigh it again. Repeat this heating, cooling and weighing sequence until two consecutive weighings are within 0.005 g. This sequence of heating, cooling and weighing to obtain consistent results is known as heating to a constant weight.

$hat water boils a t 1OOoC. Continue to heat the crucible and

28

Calculate the actual percentage of water in the zinc sulfate using the gram formula weight of the hydrate and the weight of water indicated in the formula multiplied by 100. Calculate the experimental percentage of water in your hydrate based on water lost as a result of heating divided by the hydrate sample weight multiplied by 100. To evaluate your results and determine how well you did the analysis, you can calculate relative error using the following formula.

Actual Value - Experimental Value Percent of Actual Value X 100 = Relative Error

An absolute relative error of 3% or less is excellent. If the percent of error is greater than 3%, think about factors that you observed during the procedure that may have affected your results Modify your actions accordingly for the unknown hydrate sample.

Ohtain from your instructor an unknown hydrate sample. Repeat the procedure which you used on the known hydrate. Do a minimum of two trials and report the average percentage of water in your unknown hydrate. DO NOT forget to record your sample number and/or letters.


Data tables should be constructed for all hydrate trials so the information is readily found in the report. One sample calculation must be demonstrated for a completed trial. Record all observations and adjustments made to improve your technique.

Questions

1. Using a suitable reference source, define the following words: anhydrous deliquescent desiccant desiccator efflorescent hydrate hygroscopic

2 . Using the CRC Handbook of Chemistry and Physics find: a) five hydrates, each with a different number of waters and

b) two crystalline substances with no waters of hydration c) two substances known to be deliquescent d) two substances known to be efflorescent e ) all of the hydrates of.sodium carbonate

cover, how does this affect the reported percent of water? Explain.

4 . .If the salt decomposes to yield volatile materials, is the reported percent of water in the sample too high or too low? Explain.

5. Anhydrous calcium chloride is used inside dessicators to remove moisture. Explain how CaC12 removes the water vapor.

record their names and formulas.

3 . If the volatile impurities are not burned off the crucible and

29

EMPIRICAL FORMULA OF A COMPOUND

In a given chemical compound, the elements are always combined in the same proportion by mass. A compound’s empirical formula specifies the lowest whole number ratio of the atoms in the compound. Under a specific set of reaction conditions for two elements, only one compound can be formed. When those conditions are changed, however, it is sometimes possible for other compounds to be formed. An example of this is the reaction of hydrogen and oxygen to form two different compounds, water ( H 2 O ) and hydrogen peroxide ( H 2 0 2 1 . When water is formed, the ratio is always two hydrogen atoms to one oxygen atom and in hydrogen peroxide a ratio of two hydrogen atoms to two oxygen atoms. This illustrates the Law of Multiple Proportions which state that whenever two elements form more than one compound, the different masses of one element that combine with the same mass of the other element are in the ratio of small whole numbers. In water, 1 gram of hydrogen combines with 8 grams of oxygen. In hydrogen peroxide, 1 gram of hydrogen combines with 16 grams of oxygen. Also note that the empirical and molecular formula are the same formula for water whereas the empirical formula of hydrogen peroxide is half of its molecular formula because an empirical formula is always the lowest whole number ratio of atoms in the compound.

The following information is required to determine the empirical formula of a compound. First, the composition by weight of each element is determined. Second, the number of moles of each element is calculated by using their gram atomic weights. Third, the empirical formula is expressed as the smallest whole number mole ratio of the elements. When we determine an empirical formula, the compound’s weight composition is established by either analysis or synthesis. In analysis, a known weight of the compound is decomposed and each elements weight is measured. For instance, water can be electrolyzed to produce hydrogen and oxygen. Each gas can be collected separately and measured. In synthesis, the compound’s weight resulting from the chemical combination of the known weights of two elements is measured.

This experiment synthesizes a magnesium-oxygen compound and/or a copper or nickel-sulfur compound. Using the metal’s initial weight and the compound’s final weight, the weight of the nonmetal (oxygen or sulfur) is derived. Thus the empirical formula may be determined by calculating the simplest whole number mole ratio of the elements. The magnesium-oxygen synthesis may be done at the laboratory bench but the copperhickel-sulfur synthesis MUST be done in a fume hood because a side reaction of forming sulfur dioxide, S02, a gas which is a health hazard as an irritant to eyes, skin and respiratory system.

30

P r o c e d u r e

Section A

Support a clean crucible and lid on a clay triangle and heat with a hot, blue, nonluminous flame until it glows red. This will require that the burner be manipulated to heat all surfaces. While the crucible and lid are cooling to room temperature, obtain a 0 . 2 to 0.3 g sample of Mg ribbon and polish the surface with steel wool. Cut the ribbon into several small pieces which will conveniently fit in the crucible and divide the pieces so that two trials may be completed. When the crucible is cool, using crucible tongs, transfer it to a carrying tray and weigh it to the nearest 0.001 g. Add several pieces of Mg to the crucible and reweigh it. The Mg sample should weigh 0.lg or more. Place the crucible containing Mg ribbon on the clay triangle and cover it with the lid leaving a slight opening on one side. Heat slowly. If too much air reaches the Mg, rapid oxidation occurs and it burns very brightly. If this happens, immediately cover the crucible with the lid and then adjust it for a slight opening again. Continue heating until no change is apparent in the Mg residue. Remove the lid and heat the open crucible.

Remove the flame and allow the crucible to cool for several minutes. Add a few drops of water to the residue in the crucible to decompose any magnesium nitride formed during combustion. You may detect a slight odor of ammonia. Dry the residue using a low flame and allow the crucible to cool. Transfer it to a carrying tray and weigh to the nearest 0.001g. Repeat the sequence of steps in this parargraph until the second and third weighings are within - + 0 . 0 0 8 g .

Section B

Support a clean crucible and lid on a clay triangle and heat with a hot, blue, luminous flame until it glows red. This will require burner manipulation to heat all surfaces. While the crucible is cooling, obtain 0.5 to 1.0 g of copper wire and coil it around a pencil using forceps. Once it has been coiled, compress the coil so that it will occupy a minimum of space in the bottom of the crucible. Using crucible tongs, transfer the crucible and lid to a carrying tray and weigh it to the nearest 0.001 g. Add the copper sample to the crucible and reweigh it to the nearest 0.001g. (At the instructor’s option, nickel may replace copper in this procedure.)

In a FUME ROOD, support the crucible on a clay triangle and add enough powdered sulfur, approximately 0 . 4 g, to cover the copper wire. Place the lid on the crucible and adjust it to a very slight opening on one side. Heat the crucible slowly with a low flame. When the light blue sulfur flames around the crucible’s edges are no longer visible, move flame around the crucible and lid to ensure the complete combustion of the excess sulfur. Cover the crucible completely and heat it with hot flame for at least 5 minutes. Allow the crucible, lid and contents to cool to room temperature and weigh it to the nearest 0.001g.

I

31

Return the crucible to clay triangle in the fume hood and add approximately 0.1 g of powdered sulfur and reheat it as before until all of the sulfur has burned. The second weighing of the crucible, lid and contents should be within 20 .008 g. If it is not, repeat steps in this paragraph to a constant weight value. A second trial of the experiment should be run to verify the results of first trial.

Jteauirements f or Re P ort

Data and results tables should be constructed. One sample calculation must be demonstrated for a completed trial. Record your observations and adjustments made to improve your results. Using the CRC Handbook of Chemistry and Physics for the actual formula of the synthesized compound, calculate the relative error on the average of your two trials.

Quest ions

1. Using a suitable reference source, define the following words: a) elemental analysis b) decomposition 6 ) formula d) synthesis e) gram equivalent weight f) oxidation

and formulas of the compounds prepared in the experiment. In the case of copper there will be two compounds listed. Search for a chemical reaction by which the other compound of copper and sulfur may be prepared. Give your source of information.

3 . If the lid is not used to control the burning of the Mg, what can happen as a result of the rapid oxidation? How would it affect the reported empirical formula?

cover, the surface color changes from blue-black to orange or orange-red. What chemical reaction is taking place. Justify your answer with specific chemical facts.

5 . When nitrogen and oxygen combine at various reaction conditions, the following compounds can be formed: N 2 0 , NO, N O 2 , N 2 0 3 , and N 2 0 5 . Calculate the amount of nitrogen that will combine with 8 g of oxygen for each compound. You will be calculating the gram equivalent weight of nitrogen in each compound which should verify the Law of Multiple Proportions.

2. Using the CRC Handbook of Chemistry and Physics, find the names

4 . If the copper sulfide compound is heated intensely without the

32

PERCENTAGE OF POTASSIUM CHLORATE IN A MIXTURE

t

Chemical reactions provide information regarding the relationships between the products and reactants. The stoichiometry of a reaction is a description of the relative quantities by moles of reactants and products as indicated by the coefficients in the balanced equation. In a balanced equation, the sum of the atoms of each element appearing in the reactants must equal their sum among the products. A balanced equation is a shorthand notation and quantitative description of a chemical reaction. For example, hydrogen reacts with oxygen to form water, but this gives no information regarding the proportions of reactants and products. Therefore, it must be translated to an equation form.

H2 + 0 2 - - > H2 0 unbalanced

ZHZ + 0 2 - - > 2Hz 0 balanced

The balanced equation may be read as two molecules of hydrogen react witb 1 molecule of oxygen to form two molecules of water. The most common interpretation of the balanced equation is in terms of moles, that is, 2 moles of hydrogen react with 1 mole of oxygen to form two moles of water.

This experiment is an example of quantitative analysis, more specifically, gravimetric quantitative analysis. When potassium chlorate, KC103, is heated, it decomposes to produce oxygen, 0 2 ,

MnO2

heat KC103 ------ > KC1 + 02 unbalanced

and potassium chloride, KC1. Manganese dioxide, MnOa, is added to the potassium chlorate as a catalyst. The role of a catalyst in any reaction is to initiate the chemical reaction and to reduce the energy require for the reaction to take place. A mixture of potassium chlorate and sodium chloride is heated until all of the oxygen has been released, leaving a residue consisting of K C 1 and unreacted NaC1. The weight loss of oxygen is calculated by subtracting residue weight from original mixture weight. The weight of potassium chlorate in the mixture is calculated from the weight of oxygen lost using the balanced equation. Thus, the percentage of potassium chlorate in the original mixture can be computed. To increase the accuracy and develop good technique for the unknown sample, the student will do two trials on known, prepared mixtures. By observation and percent relative error, the student can determine when he/she is ready for an unknown sample.

3 3

Procedure

CAUTION: Great care must be taken to avoid introduction of any organic material or other reducing agents into the potassium chlorate mixture. A serious explosion is almost certain to result from such carelessness. Goggles must be worn at all times in the laboratory. A heat resistant test tube, such as Pyrex or Kimax, must be used for the experiment.

Wash several heat resistant test tubes, rinsing each very well with distilled water, and dry in IlOOC drying oven. In a clean, dry, small beaker, 50 or 100 ml size, obtain approximately 4.0 g of potassium chlorate and cover with a clean, dry watch glass until you are ready to start the weighing process. In a clean, dry, small beaker obtain approximately 2.0 g of sodium chloride and cover with a clean, dry watch glass. Since both compounds are white crystalline solid, be sure to mark the beakers to the identify their contents. In small test tube obtain a very small amount of manganese dioxide, MnOz.

Weigh an empty, clean test tube to the nearest 1 mg. Add half of*the potassium chlorate in the beaker to the test tube and reweigh. Add half of the sodium chloride in the beaker to the test tube and reweigh. Add a tiny amount of the manganese dioxide, which is a catalyst, to the test tube and reweigh. Be sure to record all weighings. Shake the test tube to mix the catalyst uniformly throughout the contents of the test tube.

Support the test tube in a clamp at a 450 angle on a ring stand. Be sure to choose a clamp whose jaws are not rubber coated to avoid melting during the heating process. With a hot flame on the burner, start to heat the upper portion of the sample. The inclined position of the sample should allow you to control the rate of decomposition. If rapid decomposition occurs, the sample will appear to climb the walls of the test tube making good results more difficult to achieve. After a clear molten liquid has been obtained in the bottom of the test tube, proceed to heat the upper part of the tube to ensure complete decomposition. This entire process should require 10 to 15 minutes. Allow the test tube to cool to room temperature and weigh to nearest 1 mg. Repeat the heating and weighing processes until consecutive weighings agree within 0.01 g. Calculate the actual percent of potassium chlorate in your prepared sample. Calculate your experimental percentage of potassium chlorate on the basis of oxygen weight loss from the balanced equation. Calculate the relative error for your trial. If relative error is 3% or less, the student has the option of immediately doing the unknown without a required second trial.

Secure an unknown sample from instructor and weigh to the nearest 0.001g. Quantitatively transfer the sample to an empty, clean, weighed test tube and reweigh. Add a small amount of catalyst, MnOz, and reweigh. known, prepared sample. Calculate and report the percent of potassium chlorate in unknown sample on the basis of the oxygen weight loss.

Repeat heating and weighing techniques employed on the

34


Data and results tables should be constructed for known and unknown samples. Sample calculations should be included for known and unknown samples. Record observations and adjustments made during procedure to improve results.

Quest ions

1. Using a suitable reference source, define the following words: a) catalyst b ) chemical equation c) chemical reaction d ) quantitative e) stoichiometry

2 . Write the complete balanced equation for the reaction which takes place when potassium chlorate is heated and gas is evolved.

3. A mixture of 30.0% KC103 and 70.0% NaCl was heated to constant weight. There was a loss of 0 . 7 2 0 g of oxygen. Calculate the weight of KC103 in the mixture and the original mixture weight.

bften appears during the decomposition. Since MnO2 is black, use reference sources to explain this observation and cite your source of information.

matches. Explain the role that potassium chlorate plays in the chemistry of matches.

4 . When MnOz is added to KC103, a characteristic purple color

5. Potassium chlorate is sometimes used in the manufacture of

I

35

INORGANIC NOMENCLATURE

In order to communicate with chemists internationally, some standardization of the technical language is needed. In the past common names for many compounds evolved and are still universally understood, for example, water, lime, and ammonia. With new compounds being prepared daily, a random system is no longer viable. The Committee on Publications of the International Union of Pure and Applied Chemistry, IUPAC, has generated a system of rules for naming compounds. The American version of this system is found in the CRC Handbook of Chemistry and Physics. In this study assignment, you will learn some rules for naming and writing formulas for inorganic compounds. You are already familiar with some symbols for the elements and the names for several common compounds. For example, NaCl is sodium chloride. Continued practice and work in writing formulas and naming compounds will increase your knowledge of the chemist's vocabulary. The following information is supplied to help you learn the basics of the language in a systematic manner.

OXIDATION NUMBER

When elements combine to produce a compound, each element is assigned an "apparent" charge. This apparent charge, the charge an atom would have if both electrons in each bond were assigned to the more electronegative element, may be either positive or negative. It is called the oxidation number or state of the element in the compound.

Oxidation numbers are very handy bookkeeping devices for keeping track of what happens to electrons when various elements combine to form compounds. By remembering a few generalizations concerning oxidation numbers, the correct chemical formulas for a large number of compounds can be written and it becomes unnecessary to memorize disconnected chemical formulas. As in any system, a consistent s e t of rules must be followed.

Any element in the free state, not combined with another element, has an oxidation number of zero, regardless of the complexity of the molecule in which it occurs. Each atom in Ne, 0 2 , PI, and Se has an oxidation number of zero.

The oxidation number of any monoatomic ion is equal to the charge on the ion. The ions Ca+2, Fe+3, and Cl- have oxidation numbers of +2,+3, and -1 respectively.

Oxygen is assigned an oxidation number of - 2 , except in peroxides when it is -1 such as in H 2 O z and in the molecule OF2 when it is +2. The oxidation number of oxygen is - 2 in MgO, Fe203, BaO, and K z O .

36

4 .

5.

6 .

7 .

a.

Hydrogen has an oxidation number of +1, except in metal hydrides such as NaH when it is -1. Its oxidation number is + 1 in H20, NaHSO4 , and H2SO4.

Some elements exhibit only one oxidation state in certain types of compounds:

a.

b.

C .

a .

e.

f .

For the

The elements of Group IA (Li, Na, K, Rb, Cs, and Fr) always have an oxidation number of +1 in compounds.

The elements of Group IIA (Be, Mg, Ca, Sr, Ba, and Ra) always have an oxidation number of +2 in compounds.

Boron and aluminum always possess an oxidation number of +3 in compounds.

In binary compounds with metals, the nonmetallic elements of Group VIA (0, S, Se, and Te) usually exhibit an oxidation number of -2.

In all binary compounds with metals, the elements of Group VIIA (F, C1, Br, I, and At) have an oxidation number of -1.

In assigning oxidation numbers in compounds, the elements closest to the most electronegative element, fluorine, in the periodic table are always given the negative oxidation number. In the compound N203, oxygen has a -2 charge.

neutral compounds, the sum of the oxidation numbers of all atoms in the compound must equal zero.

Example: Liz Se

2 Li atoms, +1 for each (Rule 5a) = +2 1 Se at om. - 2 for e ach (Rule 5 d) = -2 Sum of oxidation numbers = o

For charged species, ions, the sum of the oxidation numbers of the elements must equal the charge on that species.

Example: SO4-2

3 0 atoms, -2 for each (Rule 3 ) = -6 1 C atom. must be +4 = +Q Sum of Oxidation numbers = -2, the ion charge.

Some chemical elements show more than one oxidation number depending on the compound. The rules above may be used to determine their values. Consider the compounds PbC12 and PbC14. Since the chlorine atom has an oxidation number of -1 when combined with a metal (Rule 5e), lead's oxidation number in PbC12 must be +2, whereas it is +4 in PbC14.

I

37

Consider the compounds H 3 P 0 4 and H 3 P O 3 . According to Rule 3, each oxygen atom is - 2 ; but since there is no rule for the phosphorus atom, it must be determined indirectly.

For H3P04 there are:

3 H atoms, +1 for each = +3 4 0 atoms, -2 for each = -8 1 P atom which, according to Rule 7, must have an

oxidation number equal to + 5 , so that +3 + ( - 8 ) + ( + 5 ) = 0 ;

and for H 3 P 0 3 , there are:

3 H atoms, +1 for each = + 3 3 0 atoms, -2 for each = -6 1 P atom which, according to Rule 7, must have an

oxidation number equal to + 3 . so that + 3 + (-6) + (+3) = 0 .

Students will apply these rules to a set of exercises in the study assignment to become familiar with their usage.

NOMENCLATURE FOR BINARY COMPOUNDS

1 Metal and Nonmetal - Salts. Compounds of two elements are named directly from the elements involved. In naming a binary salt, it is customary to list the more metallic or more electropositive element first. The root of the second element is then listed with the suffix -ide added to it. NaCl is called sodium chloride, A1203 is aluminum oxide. The names of a few exceptions having the -ide ending, but yet not truly binary compounds, are those that contain the N H 4 + , O H - , and CN- ions. NhC1 is ammonium chloride; NaOH is sodium hydroxide; and KCN is potassium cyanide.

When a metal cation exhibits more than one oxidation number, the nomenclature becomes somewhat complicated. Two systems are used for expressing different oxidation numbers of a metal. The old system applies a different suffix to the Latin root name for the metal. The -ous ending designates the lower oxidation number, whereas the -ic ending indicates the higher one. The 6econd system, sometimes referred to as the "Stock" system, uses Roman numerals after the English name to indicate the metal's oxidation number. Some examples which illustrate the use of the two systems are iron, copper and lead salts.

Formula "Old System" Stock System

FeClz ferrous chloride iron(I1) chloride FeCla ferric chloride iron(II1) chloride cu I cuprous iodide copper(1) iodide CUI2 cupric iodide copper(I1) iodide PbBrz plumbous bromide lead(1I) bromide PbBr4 plumbic bromide lead(1V) bromide

In naming compounds formed between metals and nonmetals, the Stock system is the preferred one in chemistry today, however many chemical manufacturing companies still use old names on reagent bottles.

2 . Two Nonmeta I s . Often more than two compounds form from the chemical combination of two nonmetals. Nitrogen and oxygen, for example, combine to form NzO, NO, NOz , N203 , N z O 4 , and N z 0 5 . It is obvious that these are all nitrogen oxides, but to identify them individually, Greek prefixes are used to designate the number of atoms present in the molecule. Below is a list of common prefixes.

Prefix Heanin4

mono- di- tri- tetra- penta- hexa- hepta- octa- nona- deca-

one two three four five six seven eight nine ten

The mono prefix is seldom used, since one is generally understood unless otherwise specified. The nitrogen oxides mentioned above are named dinitrogen oxide, nitrogen oxide, nitrogen dioxide, dinitrogen trioxide, dinitrogen tetroxide and dinitrogen pentoxide, respectively. Occasionally, part of the prefix is omitted from the spelling to simplify the pronunciation. The Stock system is sometimes used in naming a compound formed between two nonmetals. In this system the list of nitrogen oxides above would have the names, nitrogen(1) oxide, nitrogen(I1) oxide, nitrogen(1V) oxide, nitrogen(II1) oxide, nitrogen(1V) oxide, and nitrogen(V) oxide respectively. As you can see, the Stock system does not differentiate between molecular formulas, such as NO2 and N 2 0 4 , since both are called nitrogen(1V) oxide. It is suggested that students use prefixes to avoid confusion in naming and formula writing.

39

3 . Binary Acids. A binary acid is a water solution of a compound formed between hydrogen and a more electronegative nonmetal. To name the acid, the prefix hydro- and the suffix -ic is added to the root name for the nonmetal.

HC1 hydrochloric acid when dissolved in water, hydrogen chloride as the gas.

HBr hydrobromic acid when dissolved in water. H2 s hydrosulfuric acid when dissolved in water..

NOHENCLATURE OF COMPOUNDS CONTAINING POLYATOMIC ANIONS

1. Metal Cation and Polvatomic Anion Containina Oxvnen - Salts. Most polyatomic anions consist of one element which is a nonmetal and oxygen; the entire group of atoms carries a negative charge. When a polyatomic anion combines with a metal to form a salt, the anion is named by using the element’s root with the suffix -ate. The sulfate ion is SO4-2, the nitrate ion N03- , and the chromate ion, CrO4-2. If the element in the polyatomic anion has two different oxidation numbers the ion having the element with the higher oxidation number carries the -ate suffix, whereas the ion containing the element with the lower oxidation number carries the -ite suffix. For example, SO3-2 in which sulfur is +4 is called sulfite and SO4-2 in which sulfur is +6 is called sulfate. Some examples of typical salts are properly named in the following list.

Na3 PO4 P is + 5 sodium phosphate Na3 PO3 P is + 3 sodium phosphite

LiN03 N is +5 lithium nitrate LiNOz N is + 3 lithium nitrite

K2 Se04 Se is +6 potassium selenate K2 SeO3 Se is +4 potassium selenite

2. Oxyacids. An oxyacid is a compound consisting of hydrogen and a polyatomic anion (where one atom is oxygen) dissolved in water. It is named by adding the suffix -ic to the root of the element other than oxygen in the polyatomic anion. Sulfuric acid is H2SO4, nitric acid is HN03, and chromic acid is HzCrO4. If an element in the polyatomic anion occurs with two different oxidation numbers in an acid, the acid having the element with the higher oxidation number uses the -ic suffix, whereas the acid having the element with the lower oxidation number uses the -ous suffix. Note the similarity between the old metal naming system and the oxyacid naming system. Sulfurous acid is HzSOs and nitrous acid is HNO2.

40

A comparison of the nomenclature between the salts and acids containing polyatomic anions shows consistency to the rules for naming these compounds.

H3As04 arsenic acid K3 As04 potassium arsenate H3AsOs arsenous acid K3As03 potassium arsenite

HzTeO4 telluric acid Kz TeO4 potassium tellurate H2 TeO3 tellurous acid K2 TeO3 potassium tellurite

Note that the-ous suffix for the acid becomes the -ite suffix for the corresponding salt and the -ic suffix for the acid becomes the -ate suffix in a salt form.

Because of the existence of certain oxyacids in which oxidation numbers of +7 and +1 can be assigned to the unique element in the polyatomic anion, prefixes are added to the base acid names of the two common oxidation states. A representative group of acids illustrating this are the halogen containing oxyacids and their corresponding salts.

HC104 perchloric acid LiClOi lithium perchlorate HC103 chloric acid LiClOa lithium chlorate HClO2 chlorous acid LiClOz lithium chlorite HClO hypochlorous acid LiClO lithium hypochlorite

NOMENCLATURE OF ACID SALTS

Acid salts are salts in which a metal replaces less than the total number of hydrogens in the parent acid. The presence of hydrogen is indicated by inserting its name into that of the salt.

KHSOi potassium hydrogen sulfate, one K+ replaces one

MgHP04 magnesium hydrogen phosphate, Mg+2 replaces

NaH2P03 sodium dihydrogen phosphite, Na+ replaces one

NaHC03 sodium hydrogen carbonate, Na+ replaces one

H+ in HzSO4

two H+ in H3PO4

H+ in HaPOa

H+ in HzCO3

The older system of naming acid salts substituted the prefix bi for the existence of a single hydrogen in the polyatomic anion. For example, NaHCOs was called sodium bicarbonate and NaHS was named sodium bisulfide.

WRITING FORMULAS

In writing formulas it is mandatory that the sum of the elements’ oxidation numbers in the compound always equals zero.

Example - Write the formula for strontium fluoride. From the Table of Common Names and Oxidation Numbers you find that strontium exists as Sr+2 and fluoride as F-. Therefore, in order for the sum of the oxidation numbers to be zero, the formula must be SrFz.

41

Example - Write the formula for aluminum sulfide. Aluminum is Al+3 and sulfide is S - 2 . The formula must be A l z S 3 .

Example - Write the formula for potassium dichromate. Potassium is K+ and dichromate is C r z O r - 2 . The formul must be KzCrz07.

Example - Write the formula for tin(1V) carbonate. Tin(1V) is Sn+4 and carbonate is CO3-2. The formula must be Sn(C0a)z.

42

A . METALLIC ELEMENTS AND POLYATOMIC IONS

Oxidation Number = + 1

NH4+ Ammonium

Cu+

H+ Hydrogen

Hgz+z Mercury(I),Mercurous

Copper ( I ) , Cuprous


Ba+2 Barium

Cd+2 Cadmium

Cat2 Calcium

Cr+2 Chromium(II),Chromous

Co+2 Cobalt(II),Cobaltous

Cu+2 Copper( 11), Cupric

Fe+2 Iron(I1) ,Ferrous

Pb+2 Lead(II),Plumbous


Al+3 Aluminum

Lit

K+

Ag+

Na+

Mg+ 2

Mn+ 2

Hg+ 2

Ni+ 2

Pd+ 2

Sr+ 2

Sn+ 2

Zn+ 2

co+ 3

As+3 Arsenic(III),Arsenious Fe+3

Cr+3 Chromium(III),Chromic Sc+3

idation Number = + 4

c+ 4 Carbon Si+ 4

Pb+4 Lead(IV),Plumbic Sn+ 4

Mn+4 Manganese( IV) zr+ 4

Oxida tion N umber = + s As+5 Arsenic(V),Arsenic v5 +

Lithium

Pot ass ium

Silver

Sodium

Magnesium

Manganese(I1)

Mercury(II),Mercuric

Nickel(I1)

Palladium(I1)

Strontium

Tin(II),Stannous

Zinc

Cobalt(III),Cobaltic

Iron(III),Ferric

Scandium

Silicon

Tin(IV),Stannic

Zirconium

Vanadium

I

4 3

Table I (continued)

B. NONMETALLIC ELEMENTS AND POLYATOMIC IONS

Oxidation Number = - 1

CH3 COO- Acetate or

c2 H3 02 - Acetate

Br- Bromide

‘2103 - Chlorate

c1- Chloride

CN- Cyanide

F- Fluoride

OH- Hydroxide

Oxidation Number = - 2

CO3 - 2 Carbonate

CrO4 - 2 Chromate

Cr2 07 - 2 Dichromate

s302-2 Dithiosulf ate

Si03 - 2 Silicate

0- 2 Oxide

- - Oxidation Number - BO3 - 3 Borate

As04 - 3 Arsenate

As033 - Arsenite

N- 3 Nitride

H- Hydril

(210-

I-

NO3 -

NO2 -

C104 -

I04 -

MnO4 -

e

Hypochlorite

Iodide

Nitrate

Nitrite

Perchlorate

Periodate

Permanganate

C204-2 Oxalate

02-2 Peroxide

SO4 - 2 Sulfate

SO3-2 Sulfite

s- 2 Sulfide

Sz 03- 2 Thiosulfate

P- 3 Phosphide

PO4 - 3 Phosphate

PO3 - 3 Phosphite

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44

ASSIGNMENT

Obtain several sheets of paper from your instructor and complete the following exercises.

1. Determine the oxidation number of the underlined element in each of the following compounds or ions.

a. GO m. EC13'

b. GO2 h. QF2 n. SO2

c. LF7 i. NBr3 0. 5Cl2

d. b O 2 p . H2Q2

2 . Refer to Table I and locate all cations that have more than one oxidation number. List them with names according to t h e "old" system and the Stock system.

3 . Name the following compounds.

a. b.

d. e. f. 8 . h.

C.

CsF SrBra Ca ( BrO3 ) 2 Mg2 N2 S i O z P2 0 5 KMnO4 Na2 CrO4

i . j. k. 1. m. n. 0 .

P.

Liz S NaCz H3 02 ClF3 AgC104 Mg2 SO3 HC 1 P2 s3 Na3 PO4

9. r. 5 . t. U. V. W . X.

v2 0 5 NaCN Cd(0H)z Ba(N02 )2 Na2 02 NH4 C10 NiCz 0 4 a1p03

4 . Name the following compounds by both "-ous, -ic" and Stock system.

a. b.

d . e. f. g . h.

C.

CrS COP CrCls CuCN PbO Hgz Iz cuco3 As2 Os

i. Sn(NOa)4 j. CoSiO3 k. SnFz 1. Mn(0H)z m. As2S3 n. HgSz03 0. FeB03 p. Co(N02)z

Q . r. S. t. U. V . W. X .

Pb( C204 ) z FeSOi mn02 cuc104

Hg3 N2 PbCrO4 As2 (SO3 13

c0p04

45

5. Name the following as acids.

a . HBr b. HClO c. H2S d. H2S04

i. H3BOa e. H3POa f. H3AsO4 j. HCN g . HMnO4 k. HNO2 h. H2CrO4 1. HzCOa

6. Name the following acid salts. Use the older method where possible.

a. KHSO3 d. KH2AsOa g. NH4H2PO4 b. NaHSO4 e. NaHS h. K2HAs04 c. Ca(HC03)2 f. NaHC03 i. LizHPOa

7. Write formulas for each of the following compounds.

a. b.

d. e. f. g. h. i. j. k. 1. m. n.

C.

Ferrous sulfate Potassium permanganate Calcium carbonate Iron(II1) oxide Cupric hydroxide Aluminum sulfide Mercury(I1) chloride Cadmium sulfide Copper(1) chloride Ammonium cyanide Sodium chromate Nickel nitrate Manganese(I1) oxide Manganese(1V) oxide

0. P. 9. r. S. t. U. V .

W. X .

Y. 2 . aa. bb.

Lead(I1) carbonate Stannous chloride Sodium nitrite Barium acetate Silver thiosulfate Nitrogen triiodide Potassium iodide Sodium silicate Calcium hypochlorite Potassium chlorate Cuprous iodate Sodium dihydrogen phosphite Ammonium oxalate Potassium dichromate

8. Write formulas for each of the following compounds.

a. b.

d. e. f.

C.

g. h. i.

Hydrogen fluoride Sulfuric acid Iodine pentafluoride Hydrobromic acid Hypobromous acid Hydrogen sulfide Sulfur trioxide Phosphorus pentafluoride Potassium borate

j. Perchloric acid k. Nitrous acid 1. Silicon dioxide m. Acetic acid n. Cobalt(II1) thiosulfate 0 . Nitric acid p . Hydrocyanic acid q. Carbon monoxide r. Chlorous acid

46

MOLECULAR GEOMETRY

A molecule’s structure is necessary to explain its chemical and physical properties. For example, the fact that water is a liquid at room temperature, dissolves innumerable salts and sugars, becomes denser than ice, boils at a relatively high temperature, and has a low vapor pressure results from its molecule being bent rather than linear.

Because structure is so important, chemists have developed theories ‘to explain molecular shapes. In 1916, G . N . Lewis proposed a theory accounting for the significance of valence electrons in bonding. He proposed the “Octet Rule” in which atoms form bonds by losing, gaining, or sharing enough electrons in order to have the same number of valence electrons (eight) as the nearest noble gas in the Periodic Table. The bond formed is ionic or covalent depending on whether the electrons are transferred or shared. The Octet Rule is valid for most compounds formed between second and third period representative elements. The Lewis dot structure for HzO shows that by sharing electrons between the oxygen and hydrogen, all three atoms achieve the same number of valence electrons as the nearest noble gas. Hydrogen atoms are isoelectronic with the helium atom and the oxygen atom is isoelectronic with the neon atom. Lewis dot structures account only for the valence electrons on each atom and only produce two-dimensional structures. They do not explain how electrons are shared, nor do they predict a compound’s actual structure.

H:O:H m .

VALENCE-BOND THEORY

The basic hypothesis of Valence-Bond Theory is that a covalent bond forms when adjacent atomic orbitals overlap sharing pairs of electrons. When these overlapping, adjacent atomic orbitals point directly at one another, the bond is a sigma-bond. For example, a sigma-bond forms between a hydrogen atom and a chlorine atom making the HC1 molecule if the hydrogen atom’s 1 s atomic orbital overlaps a chlorine atom’s 3p atomic orbital. The result is a pair of shared valence electrons between two atoms and sigma-bonding.

47

In 1931, Linus Pauling proposed that the orientation of orbitals involved in sigma-bonding determines the shape of a polyatomic molecule or ion. A detailed look at the bonding and structure of methane, C H 4 , supports this.

The carbon atom’s electronic configuration, Is2 2s2 2p2, indicates that only two free electrons are available for bonding.

2p- - - 11

2s - However, for methane all four C-H bonds are equivalent, all H-C-H bond angles are 109.50, and the structure is tetrahedral in shape. To account for the CH4 properties, Valence Bond Theory states that a mixing (hybridization) of the 2s and the three 2p atomic orbitals (valence orbitals) on the carbon atom provides four equivalent atomic orbitals, each containing a single electron that may be used in bond formation with the hydrogen atoms’s atomic orbitals. See the figure below.

t f 9 . f ---- Hybridization 2p- - -

41 I + 2s - 1.

Four sp3 orbitals

This hybridization is called sp3, one s and three p atomic orbitals. Each hydrogen atom 1s is atomic orbital overlaps with one sp3 hybridized atomic-orbital on the carbon atom forming a sigma-bond. The four C-H bonds are oriented in space as a tetrahedron to diminish electrostatic interaction between the hydrogen nuclei. See the following figure. For some molecules, a

- pair(s) of non-bonding electrons in the central atom’s valence shell occupies a hybridized orbital(s). Because this hybridized orbital(s) is the same as the hybridized orbitals which form sigma-bonds, it also contributes equally to the molecule’s geometry.

48

Therefore, the orientation of all hybridized valence shell orbitals (those forming the sigma-bonds and those containing pairs of non-bonding electrons) determines a molecule’s gdometric shape. A summary of hybridization models for valence shell orbitals and corresponding geometries are illustrated in Table I.

VALENCE SHELL ELECTRON PAIR REPULSION(VSEPR) THEORY

VSEPR Theory proposes that a molecule’s geometry is determined primarily by the repulsive interaction of electron pairs in the valence shell of its central atom. The orientation is such that their separation of the electron pairs is maximized and the electrostatic interactions minimized.

In methane, C H 4 , the carbon atom, the molecule’s central atom, has four bonding electron pairs in atom’s valence shell. Repulsive interactions between these four pairs decrease when they are arranged in a tetrahedral geometry. One can generalize that all molecules having four electron pairs in the valence shell of its central atom have a tetrahedral arrangement of electron pairs. This arrangement is also true for the oxygen atom in H2O.

The preferred arrangement of bonding electron pairs around the central atom gives rise to the corresponding geometric shapes of the molecules. In addition, the number of non-bonding pairs affects the orientation of these bonding electron pairs. The various molecular shapes are in Table I, using the following VSEPR notations:

A - refers to the central stom X - refers to number of bonding pairs of electrons on A E - refers to number of nonbonding pairs of electrons on A .

If a molecule has the formula AXmEn, it means there are m + n electron pairs in the central atom’s valence shell - m bonding and n non-bonding electron pairs. The general shapes ,possible are linear, trigonal planar, tetrahedral, trigonal bipyramidal, and octahedral indicating a total of 2, 3 , 4 , 5, and 6 electron pairs on the central atom respectively.

I

49

POLARITY

Once a molecule’s three-dimensional structure is determined, its degree of polarity can be established. If a bond forms between two atoms having different electronegativity values, the bond is polar. The magnitude of this difference in the bonded atoms measures the degree of the bond’s polarity. A vector, the dipole moment, represents this difference. For example, HF is more polar, that is, has a greater dipole moment, than HC1 because the fluorine atom is more electronegative that the chlorine atom.

If more than one polar bond exists in a molecule, the entire molecule may be polar or nonpolar, depending on its 3-dimensional shape. Consider the two molecules, BF3 and H 2 O . The BF3 molecule has the VSEPR formula A X 3 , where each bond is polar. Table I shows that the BF3 molecule has a trigonal planar geometric shape. The three fluorine atoms attract the bonding

0 .

electron pairs from the boron atom with equal magnitude. As a result, the geometric sum of the dipole moments of the molecule’s three bonds is zero, and the molecule is nonpolar, even though each B-F bond is polar. The HzO molecule has the VSEPR formula AX2E2, and each bond is polar. Table I shows that H2O has

H H a V-shaped geometry. Since the geometric sum of the bonds’ dipole moments in the molecule is not equal to zero, the molecule is polar.

The LikQncss o f Valence Bond ThQorY t o VSEPR ThQorY VALANCE BOND THEORY VSEPR THEORY

Non-bondin6 thbilats

I

2 ~

0

I

1

I

2 I

Hybridization

sp’d

rp’d’

& I

Valance Shell

Electron Pairs

3

4

5

6

VSEPR Formula

AX, E

AX4 E

--------

Electron-Pair Orientation

-

Geometric Shapes

linear ,-..--.-.-... trigonal planar

V-shaped

tet rahedral

t rigonrl pyramidal

.-.-.--..-.-a

V-~haped .------.....-* trigonal by pyramidal

irregular tetrahedral

T-shaped

linear --.-----...- octahedral

square pyramidal

squire planar

Example

NH,, PCI,, H,O+

‘SF4, TcCl.

ClFs

la;. XcF,

BrF,

ICI;. XeFI cn 0

I

51

In this study assignment, a series of molecules and ions will be given as formulas with the central atom underlined. Students will write Lewis structure and determine hybridization of the central atom’s valence shell orbitals, the VSEPR formula, and the molecule’s geometric shape and polarity.

ASSIGNMENT

With the help of Table I and molecular models, construct the following molecules or ions and write the Lewis formula. Determine the hybridization of the central atom’valence shell orbitals, VSEPR formula, geometric shape and polarity. The molecule, CF4 is down as an example. On a separate sheet of paper, construct a table similar to Table I for the following molecules and/or ions.

EXAMPLE

Molecule Lewis Formula or Ion ..

: F :

: F : C : F :

: F :

1. CF4 .. .. .. .. .. ..

.. 3-D Structure

F

Bonding Orbitals or pairs 4 t.-nbonding Orbitals or Pairs 0 ti;,. ridization: s P3 v’; L , L. Formula: AX4 Geometric Shape: tetrahedral Polar or Nonpolar: nonpolar

F

2 . cH3C1 3. H3Q+ 4 . BH3 5 . uH4+ 6. SF4

7 . &F2 8 . k F 4

10. &F3 11. u 0 2 -

9. EC12F3

1 2 . BF4- 13. EF5 1 4 . 2 0 4 - 3 15. s F 6 - 2 1 6 . SF6

Questions

1.

2 .

Regardless of which theory of covalent bonding is used, the predicted molecular shape of a given molecule or ion is the same. Comment on the the rationale for having several different bonding theories in lieu of one particular accepted theory. A pi-bond (VB theory) or a double-bond (VSEPR theory) does not affect a molecule’s geometric shape. a . What hybridization of valence shell orbitals on the S and

C atoms in SO2 and C O 3 - 2 are predicted? Draw the Lewis formula for each.

b. What geometry would you predict for SO2 and C O 3 - 2 ?

I

52

LXMITING REACTANT

Many times an analysis will be based on the fact that you have a limiting reactant involved in the chemical reaction. The limiting reactant determines the percent yield of product. Because chemicals react stoichiometrically, only a limited amount of product forms from given amounts of reactants. For example, the analysis of a sample containing soluble sulfate salt such as sodium sulfate can be performed by dissolving the sample in water and adding a solution containing barium chloride to form an insoluble barium sulfate salt

NazS04(aq) + BaClz(aq) - - > BaS04(s) + 2NaCl(aq)

and sodium chloride which is soluble. Since all of the salts, except BaS04, are soluble, the net ionic equation for the chemical reaction is b

Ba+2(aq) + SO4-2(aq) - - > BaS04(s).

One mole of barium ion, from BaC12- 2H2O (244.2g) in solution, reaets with 1 mole of sulfate ion, from 1 mole of Na2SO4 (142.lg) in solution, to produce 1 mole of barium sulfate precipitate (233.4g) if the reaction takes place completely. The reaction is referred to as being quantitative since it will allow the determination of the amount of sodium sulfate in the sample. Sodium sulfate is the limiting reactant and to ensure that a total, complete reaction takes place an excess of barium chloride solution is added.

A 1.000 g sample of unknown salt mixture containing sodium sulfate is weighed to the nearest 1 mg. The number of moles of sodium sulfate are calculated as if this were a pure sample of the

1.OOOg NazS04 X 1 mole NazS 0 4 = 7.030 x 10-3 142.1 g NazSO4 moles Na2SO4

sodium sulfate, thus representing maximum number of moles of sodium sulfate which may be present in the sample. Using the balanced equation

1.00 mole sO4-2 produces 1.00 mole of Bas04

or

7.030 x 10-3 mole Sod-2 produces 7.030 x 10-3 mole Bas01

and

7.030 x 10-3 mole Bas04 x 233.4 LI R as04 = 1 . 6 4 1 g 1 mole Bas01 Bas04

53

Thus, 1.641 g Bas04 represents the maximum amount of precipitate that can be formed if the sample is 100% pure. This is referred to as the theoretical yield of the product. Since the sample is not pure sodium sulfate, you will obtain less than the theoretical amount which we will call the experimental value. The percentage of sodium sulfate in the unknown can be calculated using the percent yield formula which is defined as:

percent yield = experimw ' tal Yield x 100 theoretical yield

Many industrial products such as food ingredients, deodorants] mineral waters, and cosmetics would use this procedure for analysis of sulfate ion.

In this experiment an unknown salt mixture containing NazSO4 is analyzed for the percentage of NazS04 by adding excess BaClz solution to ensure completeness of reaction. The other ingredient in the mixture is a nonreactive, soluble substance which will not interfere with the quantitative reaction. The precipitate, BaSO4, wi1.l be vacuum filtered, dried, and weighed to obtain experimental yield of BaSO4. The theoretical yield will be based on 100% activity of the salt mixture sample weighed.

Procedure

Weigh a clean, empty 400 ml beaker to the nearest 0.001 g. Add about 1.0 g of the unknown salt mixture to the beaker and reweigh the beaker to the nearest 0.001 g. Calculate the number of moles of sodium sulfate that theoretically are contained in this sample weight. Also calculate the theoretical yield of Bas04 as demonstrated in the introduction section of this experiment. Add 200 ml of distilled water to the beaker and stir with a stirring rod. In the hood, measure 1 ml of concentrated HC1 in a small graduated cylinder and with the aid of a stirring rod add to the solution in the beaker. (CAUTION: If spilled concentrated HC1, hydrochloric acid, is a severe skin irritant-flush affected area with large amounts of water. Check MSDS of hydrochloric acid and include other handling precautions and spill and disposal procedures in your notebook.) In graduated cylinder, measure 50 ml of 0 . 5 M BaClz and add to the solution in the beaker and stir for several minutes with a stirring rod. Allow the precipitate to settle. Verify that the amount of barium chloride added to the solution does, in fact, represents an excess of the reactant.

With the stirring rod in the beaker, cover the beaker with a watch glass and place the entire apparatus on a hot plate. Heat at low setting to maintain solution temperature between 80 - 9OoC for 1 hour. Check the temperature of the solution with your thermometer oocasionally. Avoid boiling the solution. Heat approximately 100 ml of distilled water in a beaker to be used as the washing solvent. Refer to filtration techniques in the Physical Separation of Mixtures experiment or pages 266 - 268 in Chemical Technician's Ready Reference Handbook, CTRRH.

54

While the solution is heating, assemble the vacuum filtration apparatus. Two types of filters will be used: a Buchner funnel with a paper filter and a fritted glass filtering crucible. Students should read pages 272 - 276 in the CTRRH. The advantage of vacuum filtration is the speed of isolation for a solid product from solution. One trial will be done recovering the precipitate on the paper and a second trial will recover the solid in the filtering crucible. It will require the recorded weight of both, paper and crucible. The order in which these two isolation methods are done is not a significant factor.

Before the filtration process is started, tip the beaker and allow the precipitate to settle under the lip of the beaker. While the solution is still hot, using beaker tongs, decant the supernatant liquid trying to minimize solid transfer to the filter. Wash the precipitate several times with warm distilled water, using a rubber policeman to remove any solid adhering to the sides of the beaker. Using the wash bottle, transfer the precipitate to the filtering medium and rinse several times with the warm distilled water. Allow air. to be drawn through the filtering medium for several minutes after the last rinse. Break the vacuum seal at the filtering flask and remove the filtering medium to a clean, dry, preweighed, marked watch glass. Place the watch glass and filtering medium in drying oven for several hours or overnight. When the precipitate is dry, allow it to cool and weigh watch glass and filtering medium. Calculate the weight of Bas04 isolated. Calculate the percent yield of the reaction. Report the average percentage yield on two trials.


Following the established guidelines for the report and in the conclusion section compare the two vacuum filtering mediums with respect to their advantages and disadvantages.

Quest ions

1. Using a suitable reference source, define the following words: a) ionic equation b) molecular equation c) net ionic equation d) spectator ion e) supernatant liquid

2 . The solublity of the precipitate, BaSO4, is 1.05 x 1 0 - 5 M in water at 250C. Based on a total volume of 250 ml, a) how many moles of Bas04 are not isolated in the precipitate? b) how many grams of Bas01 are not isolated in the precipitate? c) how significant is this weight in relationship to the

during the filtration process if coarse filter paper is used. How

4 . What is the purpose of washing the precipitate? 5 . If NaaPO4.is an unknown contaminant of the salt mixture, how would

this affect the yield of Bas04 in the experiment? Baa (PO4 12 is insoluble also. Write balanced equations as part of your explanation.

measurement you can make? 3 . Since Bas04 is a finely divided precipitate, some may be lost

. would this affect the reported percent yield? Explain.

55

PERCENT OF OXYGEN IN AIR

This is a simple experiment which may be used to introduce students to the need for monitor conditions when measuring gases. Air is a homogeneous mixture of gases such as nitrogen, oxygen, argon, and trace amounts of other elemental gases and carbon dioxide. The amounts of each gas can be measured both by weight and volume to determine the percent composition. In this experiment, the students will measure gas volumes using gas measuring burets. A 50 ml sample of air undergoes a chemical reaction in which only the oxygen is consumed. Consequently, the measurement of the original volume of air minus the remaining volume of gas after the reaction results in the volume of oxygen. To calculate the percent of oxygen, the volume of oxygen is divided by the air sample volume times 100.

Since gases are very sensitive to changes in temperature and pressure, the students should carefully note atmospheric pressure, labaratory and water temperatures. The instructor will demonstrate the-use of a mercury barometer. The students are required to read pages 479 - 4 8 2 in the Chemical Technician Ready Reference Handbook, CTHRH. The two reagents used in this experiment are a 1M NaOH solution and pyrogallic acid. Each should have the MSDS precautions for use listed in the laboratory notebook.

Procedure

To ensure that exactly 50 ml of air will be taken in this experiment, the volume of the sodium hydroxide solution used, will be equal to the volume between the 50 ml mark on the gas measuring tube and the bottom of the rubber stopper used to close the tube. To determine this volume, fill the tube with water so that the meniscus is at the 50 ml mark. Insert a rubber stopper and mark, with tape, the position of the stopper’s bottom. This will allow you to reproduce the exact volume each time. Invert the tube and read, using the graduations, the volume of air which was between the bottom of the stopper and the 50 ml mark. This is the exact volume of sodium hydroxide solution to use each time. Empty the water out of the tube. Pour sodium hydroxide solution into the buret and use a medicine dropper to get to the exact volume measurement required.

Measure 0.58 of pyrogallic acid on a weighing paper and add it to the tube. Quickly close the tube inserting the stopper to the marked position. Invert the tube several times to ensure thorough mixing of the solution and solid. Record your observations. Gently rock the solution back and forth in the tube for several minutes. Record the laboratory temperature.

56

Plug a lab sink with an overflow device and fill sink to the top of the overflow. Record water temperature in the sink. Adjust the temperature of water in the sink to the laboratory temperature using a mixture of warm and cold tap water. Submerge the stoppered end of the tube and remove stopper. Note how water is forced in to replace the oxygen. .Allow the tube to cool in the water for several minutes, then raise or lower the tube so that the water levels, inside and outside the tube, match. It is not necessary to correct the gas volumes to STP in this method because all operations are done within a few minutes. However, if there is a significant difference between laboratory temperature and the water temperature in the sink, use Charles’ L a w to correct the volume of the gas remaining in the tube. It is left for the student to calculate what represents a significant temperature difference. The atmospheric pressure should also be monitored periodically during the laboratory period to ensure that a s i g n i f i c a n t pressure change is not taking place due to weather changes. Student can calculate what would represent a significant change by using Boyle’s Law. Repeat the procedure a minimum of three times.

Calculate the average percentage of oxygen in the air. Check the CRC. Handbook or suitable reference for the actual percentage of oxygen in air. Calculate the percentage of relative error in your determination.

Requirement for Report

Use established format for reporting results of the experiment. Please note all observations and significant temperature and pressure calculations you are required to make.

Quest ions

1. Why are the water levels inside and outside the tube matched? 2 . Which correction factor, temperature or pressure, would have a

greater effect on your experimental data? Explain. 3 . If 50 ml of air is analyzed according to the above procedure and

the remaining volume of gas is 4 8 . 5 ml; a) what would be the percent of oxygen and the percent error? b) what factors could have contributed to this percent error? c) if laboratory temperature was 200C, what would the water

temperature have to be to totally account for the error? d) is the temperature calculated in c) a realistic value to

explain the error? Comment. 4 . What is another name for pyrogallic acid? What is its chemical

formula?

I

57

GRAM MOLECULAR WEIGHT OF CARBON DIOXIDE

Carbon dioxide occurs as a variable component in the atmosphere. It is formed by the decay, fermentation, and combustion of organic matter. In this experiment, carbon dioxide will be produced by reacting marble chips, predominantly CaC03, with hydrochloric acid. To obtain dry carbon dioxide, the gas is bubbled through a concentrated sulfuric acid which acts as a dessicant and collected in an Erlenmeyer flask.

The procedure involved in the determination of the molecular weight of carbon dioxide is called the gas density or vapor density method. In some texts it will be referred to as the Dumas method named for the Frenchman who is given credit for originating the method. It is based on the principle that equal volumes of gases contain the same number of molecules at the same temperature and pressure. This principle represents Avogardo’s Law which has been used to define a standard molar volume of any gas, 2 2 . 4 liters, at 1 atmosphere pressure ( 7 6 0 torr) and 273 K ( O o C ) . The ideal gas law, PV = nRT when solved for volume at STP defines this standard volume. If n, number of moles, in the ideal gas law is redefined as mass of gas/gram molecular weight (GMW)

the formula can be rearranged to solve for the gram molecular weight by using measured values obtain in the laboratory for any

PV

gas sample. The variables in the formula are mass (grams), R - gas law constant (0.0821 liter-atm/mole-K), T - temperature (degrees Kelvin), P - pressure (atmospheres), and V - volume (liters). Alternatively, the volume of gas measured at laboratory conditions could be corrected to the volume which it would occupy at STP. A simple proportion relationship would then be used to

mass of collected aaS = GMW volume occupy at STP 2 2 . 4 liters

solve for the GMW.

Procedure

Take two clean, dry, marked Erlenmeyer flasks and obtain

mark the bottom position of the stopper with a piece of tape. Weigh the flasks to the nearest 1 mg and record the weight. Record the laboratory temperature and pressure. Using CRC Handbookof Chemistry and Physics, record the density of air at laboratory conditions.

~ rubber stoppers that fit each tightly. Stopper the flasks and

58

In the fume hood, set up the apparatus as shown in the following diagram. This will require that a number of right angle glass bends must be produced by you. The student should read glass cutting (830 - 8321, glass polishing (835 -836) and glass bending (840 - 841) in the Chemical Technician's Ready Reference Handbook, CTRRH. Remember to lubricate glass with glycerin before inserting it into the tubing or stoppers. Place 40 g of marble chips in a generating flask and place approximately 30 ml of concentrated sulfuric acid, HzSO4 , in the bubbler bottle.

oc

f

Be sure to check all of the rubber and glass tubing for constrictions or blockage as the apparatus is assembled. Use a a paper punch and 3 by 5 index card to make the paper cover for the collection flask. Making sure that the screw clamp is closed between the funnel and generator, add approximately 50 ml of 6M HC1 to the funnel. The liquid level in the funnel should never exceed three-fourths of its total volume capacity nor should it be allowed to drain completely since air would enter the generator. Record in your laboratory notebook the handling precautions, and the spillage and disposal procedures for all chemicals used in the experiment from MSDS notebook.

Opening the screw clamp slightly, allow the HC1 solution to pass from the funnel onto the marble chips. A moderate rate of gas generation should be maintained and can be monitored by watching the bubbler bottle. The generator should be carefully shaken occasionally to avoid the'use of an unnecessary excess of HC1. Permit the gas to flow for 5 minutes to ensure the displacement of air in the apparatus. Touch the bottom of the generator flask and record the temperature effect. Remove the paper cover and delivery

. tube from the collection flask and quickly insert a rubber stopper to the marked position. Insert the delivery tube into a second flask to collect a sample while weighing the first flask to the nearest 0.001 g. Alternate sample collection in this fashion until a constant weight (50 .005 g) is obtained for both flasks. When all mass measurements have been completed, fill each of the

59

flasks with water to the marked position and measure the volume of water using a graduated cylinder. To disassemble, first loosen stopper on the generator flask to avoid siphoning concentrated sulfuric acid into the generator. Drain the HC1 from the funnel into a beaker and use proper disposal procedures for excess acid in the beaker, generator and bubbler bottle. R e m e m b e r that concentrated solutions are always poured into more diluted solutions.

To determine the weight of carbon dioxide it is necessary to calculate the weight of air in the flasks at the initial weighing. The density of air at laboratory conditions is multiplied times the volume of the flask to obtain mass of air. The mass of air is then subtracted from the initial flask weight to obtain weight of the stopper and flask. This is used to determine the mass of carbon dioxide collected. Using one of the two methods outlined in the introduction, calculate the gram molecular weight of carbon dioxide for each trial and the average value. Calculate the relative error for your results.

Exchange two clean, dry 250 ml flasks at the stockroom for two filled with samples of unknown gases. Do not disturb the 'stopper or warm the flask unnecessarily by handling. Mark the position of the stopper and weigh to nearest 0.001 g. In the hood remove the stopper and displace the unknown gas with laboratory air by use of the aspirator. Replace the stopper and weigh the flask of air accurately. Fill the flask with water to the mark and measure the volume using a graduated cylinder. Repeat the calculations that were necessary to determine the gram molecular weight of carbon dioxide.


Follow the guidelines fo,r experiment writeup. Be sure to included unknown sample identification numbers in the report.

Questions

1.

2 .

3 .

4 .

~ 5.

What evidence did you observe during the experiment which would enable you to decide whether the reaction generating carbon dioxide is exothermic or endothermic? Explain. If the carbon dioxide were not dried, would this tend to result in experimental values for the gram molecular weight to be too large or too small? Explain. State clearly as many practical uses for the gram molecular weight of a substance as you can. The most common drying'agent used in laboratories is anhydrous calcium chloride. Would this be a suitable substitute for the concentrated sulfuric acid used as a dessicant in this experiment? Explain. Since the laboratory air contains some water vapor, the density which you obtained from the reference is not the exact density of the laboratory air. Will this tend to cause your value for the gram molecular weight of carbon dioxide to be too high or too low? Outline your reasoning in coming to your conclusion.

60

EQUIVALENT WEIGHT OF AN UNKNOWN METAL

In the early development of the periodic table, the mass of elements were established by the Law of Combining Weights. It simply states that the number of grams of any element that will combine with, or displace 1.008 g of hydrogen or 8.00 g of oxygen. This mass is also known as the gram equivalent weight (GEW) of the element. Many of the metals on the periodic table, but not all, react with acid to form hydrogen gas which may be collected and measured. The chemical equation for zinc metal reacting with

Z n ( s ) + 2HCl(aq) - - - > ZnClz(aq) + Hz(g)

hydrochloric acid illustrates the type of reaction which takes place. Analyzing the equation, we observe that 1 mole of zinc produces 1 mole of hydrogen gas. In terms of mass, o n e gran; atomic weight (GAW) of zinc, 65.37 g, produces two GAW’s of hydrogen, 2.016 g. Therefore, the GEW of Zn is equal to one ha l f the GAW, by definition. We can also relate the volume of the gas

1 GAW of Zn = 2 GAW 8 65.37g 2.018g

1 GAW B - 1 GEW of Zn - 32.69g 1.0O8g

collected to the GEW of the metal since 1 mole of any gas at STF occupies 2 2 . 4 liters. Because hydrogen is a diatomic gas its standard molar volume actually represents 2 equivalents of hydrogen or 2GAW’s. Therefore, 1 equivalent of hydrogen will always be equal to 11.2 liters at STP. Since the number of equivalents of hydrogen will always be equal to the number of equivalents of metal, we can determine the GEW of an unknown metal. Note that the units on the GEW (g/equivalent) are similar

GEW of metal = mass of metal, number of equivalents

to units on units on GAW (g/mole). Students should a l s o observe the correlation between the oxidation number of the metal. In the case of Zn, and one mole of zinc always equals two equivalents of zinc. Zinc always has a +2 oxidation number, therefore the gram equivalent weight of a metal can be determined by dividing its gram atomic weight by the oxidation number.

6 1

To calculate the volume of hydrogen gas produced from the reaction, the student may use the Ideal Gas Law relationship, PV = nRT, or Combined Gas Law relationship, Pi Vi /Ti = PzVz/Tz . Regardless of which relationship is used, the laboratory pressure, Ft, must be corrected for the vapor pressure of water since the

Pt = P g a s + P w a t t r v a p o r

hydrogen is collected by displacing water from a collection flask. The vapor pressure of water at the laboratory temperature can be determined from a table in the CRC Handbook or other reference source. Students are encouraged to run through the procedure using zinc metal as a known until the relative error of their results are within 3%.

Procedure

Obtain a sample of unknown metal from the instructor. Weigh it to the nearest 1 mg. Tie a string around the metal sample.

Set up the apparatus shown in the illustration above. The apparatus consists of a 125 ml Erlenmeyer flask with a one hole rubber stopper attached by rubber and glass tubing to a 500 ml Florence flask, filled with water, which is supported in a water trough. Place 50 ml of 6M HC1 in the Erlenmeyer flask and suspend the metal sample on the string above the acid solution by using the rubber stopper to hold it in place. Clamp the flask in an upright position and attach rubber tubing to the glass tube extending from the rubber stopper. On the other end of the rubber tubing attach a short right angle glass tube to act as a delivery tube to the Florence flask. Submerge the delivery tube in the \

it into the flask. Record the laboratory pressure. . water trough to eliminate air in the end of the tube and insert

62

When ready to start the reaction between the metal and a c i d , tip the flask to about a 4 5 0 angle so that the sample contacts the acid solution. Allow the reaction to proceed to completion. Remove the delivery tube from the collection flask. Before removing the Florence flask to measure the volume of gas collected, allow the gas to stabilize to the water temperature by standing for 10 minutes. Record the temperature of the water trough. Remove clamp from collection flask and tip to the side, keeping the mouth of the flask submerged. Match the water level inside the flask to the level in the trough. Slip a glass plate over the end of the flask and remove it from the trough. Measure the volume of water remaining in the flask using a graduated cylinder. Completely fill the Florence flask with water and measure its total volume capacity.

Repeat the procedure for a second sample and report the average equivalent weight of the sample. Some educated guesses can be proposed for the identity of the unknown sample. When two trials have been completed, consult MSDS information on disposal of excess HC1 and record it in your laboratory notebook.


The information collected on the known zinc sample must be included in the report along with the percent error achieved on the results. Follow the established guidelines for the report.

Questions

1. If the volume of collected hydrogen is mistakenly recorded to be higher than actual, how does this affect the reported equivalent weight of the metal? Explain.

2. If all connections in the apparatus are not air-tight and some hydrogen gas escapes, is the reported equivalent weight larger or smaller than the actual value? Explain.

before the reaction for the sample is started, how would you correct the problem? If you did not observe these air bubbles until after the reaction was started, how would this affect the equivalent weight of the metal? Explain.

4 . A 0.375 g sample of a metallic element is completely dissolved a 100 ml of 0.500 N sulfuric acid. The liberated hydrogen gas occupied a volume of 350 ml when collected over water at 25oC, and the pressure was equalized to the atmospheric pressure of 750 torr. a) Calculate the equivalent weight of the metal. b) How many equivalents of acid were used to produce the

c) How many equivalents of acid remain in the solution?

3. If several air bubbles are trapped in the collection flask

hydrogen gas?

6 3

RELATIONSHIP OF TEMPERATURE AND PRESSURE TO CHANGE IN PHYSICAL STATE

A substances may exist in solid, liquid or gaseous state. The transition of matter from one of these states to another requires that a quantity of heat be absorbed ( or released) and is often difficult to measure with equipment available in a general chemistry laboratory. However, an indication of heat transfer may be obtained by monitoring the intensity factor, temperature.

The vapor pressure of a liquid is the pressure exerted by the vapor or gaseous form of the liquid when the vapor and liquid are in equililbrium with each other at a specified temperature. When the two states are in equilibrium, there is no change in pressure if the temperature remains constant; however, if the temperature is increased,the vapor pressure of the liquid increases.

The normal boiling point of a liquid is the temperature at which the vapor pressure of the liquid is exactly 1.00 atmosphere. Since the pressure acting on the surface of the liquid may differ Irom l . U O atm, the actual temperature at which boiling occurs may not be the same as the normal boiling point. When liquid boils in an open container, the pressure acting on the surface of the liquid is the prevailing barometric pressure. In a closed container from which air has been removed, the pressure acting on the liquid is that of the vapor itself. In the process of evaporation, a liquid vaporizes at the surface only; in boiling, vapor is formed in bubbles within the liquid as well. The temperature stays at the boiling point as long as any liquid remains. The absorbed heat is consumed in overcoming the attractions between molecules of the liquid and in forcing the vapr;r sut against the atmospheric pressure. Thiz quantity of absorbed heat, when expressed in calories per gram of liquid, is referred to as the specific heat of vaporization. Chemists also use the term molar heat of vaporization which is the quantity of heat, in calories, required to convert one mole of liquid to vapor at the normal boiling point. The quantity of heat energy required to overcome the cohesive forces of the molecules in the liquid is different from one liquid to another, but the quantity of heat used in overcoming atmospheric pressure depends only on the temperature and on the number of moles of liquid vaporized.

The freezing point of a pure liquid is the same as the melting point of the corresponding solid when the substance is pure. ThIs may be defined as the temperature at which pure liquid is in equilibrium with pure solid at some specified pressure, usually 1.00 atm. When the two states are in equilibrium, the solid and liquid coexist with no change in their

discussed in the preceding paragraph, the temperature remains constant at the freezing or melting point until the change in state is complete. The quantity of heat energy required to

. relative amounts with time. For reasons analogous to those

.convert a solid to a liquid, when expressed in calories per gram,

64

is referred to as the specific heat of fusion. Similarly, the molar heat of fusion is the number of calories required to convert one mole of the pure solid to liquid at a specified pressure.

Procedure

Section A

Wrap a small piece of Kimwipe or Kleenex around the bulb of a thermometer and fasten securely with a fine copper wire or rubber band. Record the temperature of the room by simply reading the thermometer. Then dip the covered bulb of the thermometer in methanol. Remove the thermometer from the liquid and suspend in the fume hood. Watch the mercury and record the lowest temperature. CAUTION: All liquids used in this section of the experiment are flammable and can constitute a combustion hazard. Check the MSDS notebook for the precautions, spillage and disposal procedures of all chemicals used in this experiment and record in your notebook.

Repeat these steps using the following liquids: propanol, isopropanol, acetone, diethyl ether, pentane, hexane, and water. Use a fresh piece of Kimwipe with each liquid. Using the CRC Handbook or suitable reference source, record the normal boiling poigts of all the liquids.

Section B (

make certain that the heat resistant glass test tubes are clean. Before assembling the apparatus shown in the figure below,

65

Remember to lubricate glass tubing and thermometers before inserting them into rubber stoppers. Do not attempt to change the position of the thermometer during the course of the experiment. Heat the water in both tubes to the boiling point with a moving flame. Continue the boiling until steam venting from the delivery tube gives a peculiar crackling sound and bubbles of gas no longer rise through the liquid in the beaker. Observe the change in the size of the bubbles as this condition is approached. At this point both thermometers should register nearly l O O o C and should agree within lo.

Close the screw clamp and immediately remove the flames from both tubes. Seal the rubber stopper with a ring of water where the stopper meets the test tube as a means of detecting any leakage. If leaks occur, the apparatus must be reconstructed and the experiment repeated. When the temperature has fallen to 9 O o C in the closed tube, note the temperature in the other tube, and record both values. Readings of the thermometer m ~ s t necessarily be omitted while the mercury in the thermometer is obscured by the stopper. Below the stopper, visibility may be clouded by condensed water droplets. To remedy this situation, rock the assembly stand until the water from the lower part of the tube washes away the droplets. Record similar observations at subsequent 100 intervals until no boiling is observed in either tube. The water will continue to boil in the closed tube down to at least 60oC and often a much lower temperature. If such results are not obtained, reconstruct the apparatus and repeat the experiment. During the course of the experiment, touch the upper part of both tubes and compare the temperature effects. Record the lowest point at which boiling occurred. At this point, cool the upper part of the closed tube with a stream of cold water from a wash bottle. Observe the rubber tubing connection on the delivery tube, Release the screw clamp with the delivery tube under water and record your observations in your notebook.

Section C

Place approximately 10 grams of crystalline sodium thiosulfate pentahydrate in a dry test tube and immerse the lower half of the tube in a water bath which has a temperature above 70oC. When all the solid has melted, remove the tube from the water bath. Stir the melted s a l t continuously with the thermometer. Record the time and the temperature readings at one minute intervals until the temperature has fallen to 50oC. At this time, add a crystal of the salt to prevent undercooling. Continue to stir and record time and temperature readings for about ten minutes after solidification begins.

Reheat the thiosulfate in the hot water bath until all solid -material has melted. All solid particles adhering to the walls of the tube must be melted or removed. Remove the the test tube from the bath and cool it under running water until the temperature is below 40oC. Observe and record the time and temperature.

66

Discontinue the cooling of the tube and drop a small crystal of the sodium thiosulfate into the melted material. Stir with the thermometer. Record the temperature and time every 30 seconds until the temperature of the mixture reaches a maximum. Then continue the observations at five minute intervals until the temperature has fallen to about 430C.

Reauirements for Report

For information in Section A , construct a table containing the names of the compounds, compound’s structure’ compound’s polarity, compound’s formula weight, compound’s normal boiling point, and compound’s lowest temperature observed. For the data collected in Section B, report the lowest temperature at which the water boiled in the open and closed tube. The data collected in Section C , should be presented in a graph using temperature as the ordinate and time as the abscissa. Both data groups may be plotted on the same graph.

Quest ions

1.

2.

3 .

4 .

5.

6.

7 .

a .

On the basis of observations recorded in Section A , make a statement regarding the relationship between the boiling points of liquids and their effectiveness in cooling by evaporation. What other factors included in the table of information for Section A appear to have a relationship to effective cooling. Considering the observations in Section B: a) Why did the rubber delivery tube collapse during the

experiment? b) Why did the water rush into the tube when the delivery tube

was opened under water? c) Why d i d the boiling continue in the closed tube at a much

lower temperature than in the open tube? d) Why did the upper part of the closed tube cool more slowly

than the upper part of the open tube? Explain why one feels cooler in a wet bathing suit than a dry one, especially when the wind is blowing. Why is a hot day less disagreeable when the atmosphere is relatively dry than a day with the same temperature when the relative humidity is high? Explain how cooling by evaporation is used in one of the methods for determining the relative humidity. Explain briefly how a pressure cooker utilizes a principle studied in this experiment? What is the meaning of supercooling? Was it observed in this experiment?

67

VISCOSITY

Viscosity is the resistance to flowing of a liquid, and like other properties of liquids, result from the intermolecular forces of attraction. As temperature is increased, the liquid’s viscosity decreases. The viscosity of a liquid can be determined by measuring the time required for a specified volume of liquid to empty through a capillary tube in a simple laboratory apparatus called a viscometer. Liquids with moderate to low viscosities may be readily measured in this manner. Using the simple viscometer, a set of simultaneous equations can be derived for determining the viscosity, n, of any liquid by measuring the flow time in seconds

for the liquid. In equation (l), A is a constant, dependent on the volume capacity of the viscometer, B is a constant, dependent upon the length and radius of the capillary, D is the density of the liquid, and t is the flow time of the liquid in seconds. So that

constants A and B may be conveniently determined for your viscometer, hexane and water must have their flow times measured. By obtaining their respective viscosities and densities from the CRC Handbook of Chemistry and Physics, the following two simultaneous equations may be solved for A and B by using determinants or substitution. In equation ( 2 ) , the viscosity

nz = ADztz - BDz/tz ( 3 )

of water, ni , and the density of water, Di , at laboratory temperature are combined with water’s flow time at the same i.=?m:.erature resulting in only A and B as unknowns in the e q i i s t i ~ n . Likewise, in equation ( 3 ) , the viscosity of hexane, n2, and the density of hexane, Dz, at laboratory temperature are combined with hexane’s flow time at the same temperature. Since viscosity is a temperature dependent property, the viscosity value for hexane may not be found in the CRC Handbook for the specific laboratory temperature at which the flow times were measured. If this is the case, use linear interpolation for two viscosity values closest to the laboratory temperature. Once the constants have been determined for the viscometer, the viscosity of any liquid may be determined by measuring the flow times in the viscometer and the density of the liquid.

For highly viscous materials, such as molasses and motor oils, a Brookfield viscometer is used to mechanically measure the viscosity by the drag of the liquid on a moving metal cylinder at - a fixed speed. Both methods will be used in this experiment.

68

Procedure

Section A

Fill your viscometer with hexane so that the liquid level at the bottom does not reach the capillary tube in the viscometer. Using a pipet bulb draw the hexane up to fill the calibrated volume area, ensuring that the liquid is drawn above the top volume line. Start the flow time measurement as the liquid passes the top line and stop timing when the liquid passes the bottom line of the volume area. All flow times should be measured at a constant temperature. If the viscometer is submerged in a beaker of room temperature water, a fairly constant temperature can be maintained. Measure the flow time of hexane several times. Pour the used hexane into the appropriate container in the fume hood. Use a pipet bulb to force air through the viscometer’s capillary to remove all traces of hexane. The viscometer can also be placed in the drying oven but must be allowed to cool before measuring the flow time for water. After the flow time for water at room temperature is measured several times, placed the viscometer in an ice-water bath for ten minutes. Record the temperature of the ice-water bath and then measure and record the flow time for water at this temperature. Calculate the constants A and B. Measure the flow times for the following compounds: methanol, propanol, isopropanol, acetone, reagent alcohol and heptane. Dispose of the liquids each time in the container in the fume hood. Be sure that the viscometer is free of all traces of residual liquid before starting the measurement of subsequent flow times. The density of reagent alcohol will have to be determined since it is a mixture of alcohols. See the appropriate sections in the Chemical Technician’s Ready Reference Handbook, CTRRH, for the pycnometer or hygrometer method.

Section B

The instructor will demonstrate proper use of the Brookfield viscometer. You will be supplied with a group of commercial products such as shampoo, creme rinse, olive oil, corn oil, and glycerin. The accuracy of the Brookfield can be checked by comparing values for the oils with those found in the CHC Handbook of Chemistry and Physics.

Rewiremen ts f o r Report

items: name of compound, compound’s formula weight, compound’s density, compound’s structure, compound’s polarity, compound’s flow time average, compound’s calculated viscosity, compound’s handbook viscosity, relative error for compound’s calculated viscosity for Section A . Show the calculations for the constants A and B. For Section B construct a table containing the following data: sample, spindle number, speed, dial reading, multiplying factor, viscosity, handbook viscosity if applicable.

Construct data and results tables containing the following

69

Questions

1. Analyze the data and results from Section A and answer the following questions: a) Why are the viscosities of propanol and isopropanol

different? b) Why are the viscosities of isopropanol and acetone

different? c) Does the viscosity of reagent alcohol occur in the sequence

of alcohols where you would expect to find i t s value? d) Reagent alcohol is a mixture of the alcohols you have

measured and ethanol. Using the percent composition of reagent alcohol from its label and the handbook values for the pure alcohcls, calculate an expected viscosity and compare this to the measured value.

2. H o w do the measured viscosities from the Brookfield viscometer compare to the handbook values?

3. Use the CTRRH or chemical supplier catalog to describe another device for measuring viscosity. Take care to describe the method used and how the device operates. Cite your reference source.

70

SURFACE TENSION

The surface tension of a liquid can be thought of as that property which draws liquid molecules together at the surface, forming a liquid-vapor interface, thereby distinguishing liquids from gases. The surface tension is a characteristic property of each liquid and differs greatly in magnitude from one liquid to another. Of the various methods available to measure surface tension, such as tensiometer, drop weight, bubble pressure, or capillary rise, capillary rise is considered the standard method. The capillary rise method is based on thc fact that most liquids, when brought in contact with a fine glass capillary tube, will rise in the tube to a level above the liquid outside the tube. This will occur only when the liquid "wets" glass, that is, adheres to it. If a liquid does not "wet" glass, as in the case of mercury, the level inside the tube will fall below the outside level. A normal liquid in contact with glass forms a concave surface which ic referred to as a meniscus. The mercury will exhibit a convex surface which appears as an inverted meniscus to an observer. To understand the theory of the capillary rise method, consider a fine capillary tube of uniform radius, r, immersed in a vessel containing a liquid that wets g l a s s (Figure 1). By wetting the inner wall of the capillary, the surface of the liquid is increased. To decrease its free surface, the liquid must rise within the capillary tube. As soon as this happens, however, the glass wets again, and the liquid draws itself upward. This process does not continue indefinitely, but stops when the force of the surface tension acting upward becomes equal to the force due to the column of liquid acting downward. If we assigned T, to be the surface tension in degrees per centimeter of the inner circumference and consider the force to be acting at an angle theta, called the contact angle with the vertical dimension, then the force due to the surface tenaici. I; given by cquation (1).

Fr = Z(pi)rTcos(theta) ( 1 )

This force is balanced by that due to the column of liquid of height, h, with a density, d, and where g is acceleration due to gravity,

980.67 cm/sec2 in equation ( 2 ) . Therefore, at equilibrium Fr = F2 which results in equation ( 3 ) and may be solved for the surface tension, T. The resulting equation ( 4 ) can be simplified

Z(pi)rTcos(theta) = (pi)r2hdg ( 3 )

71

by cancelling common terms in the numerator and denominator. For

T = (pi r2 hda =rhda Z(pi)rcos(theta) Zcos(theta)

( 4 )

most li'quids which wet glass , theta is essentially zero and thus cos(theta) = 1. Therefore, the surface tension of a liquid may be

T = rhda ( 5 ) 2

calculated if the radius of the capillary, normally 0.035 cm, the density of the liquid, and the height to which the liquid rises in the capillary are determined or known values.

Figure 1. Figure 2.

72

Procedure

First the radius of the capillary tube must be verified. The surface tension of hexane can be obtained from the CRC Handbook of Chemistry and Physics in conjunction with its density. The capillary rise of hexane is measured by filling the device in Figure 2 so that the liquid level is near the zero point on the attached scale. Record the scale measurement for the liquid level. Using a pipet bulb, draw the liquid up in the capillary above the point to which it has risen. Remove the bulb and allow the liquid to drop to its equilibrium level. Read the height of the liquid column. Remember to subtract the liquid level value if it is not exactly zero to obtain the true value for the capillary rise. This information along with the density and surface tension for hexane will allow you to solve equation (5) for the radius of the bore.

Dispose of the hexane and all other organic liquids u s e d in this experiment in the appropriately marked container in the fume hood. Using a pipet bulb, force air through the capillary to remave traces of the hexane before measuring the next liquid. Repeat this step before measuring the capillary rise of each liquid. Measure the capillary rise f o r the following liquids: methanol, propanol, isopropanol, acetone, and reagent alcohol. Record the true value of the rise. Check MSljS notebook for necessary precautions in handling these chemicals and record in your notebook. If you have recorded this information for a previous experiment, indicate on what pages in your notebook the precautions may be found.

Requirements for HeDort

Calculate the surface tension for each liquid using equation ( 5 ) . Construct a data table including the following information for each liquid; compound’s name, compound’s formula weight, compound’s density, compound’s capillary rise, compound’s calculated surface tension, compound’s actual surface tension, and relative error in the calculated surface tension. Include in the report the data table and a sample calculation for the surface tension. Discuss the any relationships that you observe from the calculated surface tension in relationship to formula weight, structure, and polarity. Discuss possible errors which may have affected your results.

Quest ions

1. Which measurement, surface tension or viscosity, appears to represent a measure of the intermolecular forces? Explain.

2. Using the CTRRH or chemical supplier catalog, cite another method of measuring surface tension and describe the operation of the device. Indicate your reference source.

73

STRENGTHS OF ELECTROLYTES

Compounds are frequently divided into two categories: electrolytes and nonelectrolytes. Electrolytes are compounds that ionize or dissociate in aqueous solution. As a result of this behavior, their aqueous solutions conduct an electric current. Nonelectrolytes are compounds that do not ionize or dissociate in aqueous solution, and therefore their aqueous solutions do not conduct an electric current.

Electrolytes are divided into two types. Strona: electrolytes are compounds that ionize or dissociate completely, or nearly completely, in dilute aqueous solution. Such solutions conduct an electric current well. Strong acids, such as HC1, H N O a , and H2S04, and strong bases, such as NaOH and KOH, are examples of s t r o n g electrolytes. In addition, most soluble salts, such as N a C 1 , KzCrz07, and MgSO4 are included in the group of aqueous solutions which conduct an electric current very well.

electrolytes are compounds that ionize or dissociate o n l y slightly in dilute aqueous solutions. Their solutions conduct an electric current very poorly. Weak acids, such as acetic acid, citric acid, and ascorbic acid, and weak bases, such as ammonia, are examples of weak electrolytes. Certain salts which are only slightly soluble in water are also very poor conductors. Examples of these salts are AgCl and MgC03.

1 7 - - 1-

Whether a salt is a strong or weak electrolyte can be predicted from the water solubility rules in your textbook. Whc? inorganic salts dissolve in water they produce ions in the solution. When aqueous solutions of salts are mixed together, they can produce a chemical reaction because of the ions. I'd ;kAis point in the course, we have represented reactants and products in molecular form for chemical equations. When silver

AgN03(aq) + NaCl(aq) ---- > AgCl(s) + NaNOs(aq)

nitrate is mixed with sodium chloride in aqueous solution, silver chloride precipitates and sodium nitrate remains in the solution. An alternate way of representing the chemical equation is to write all of the soluble salts as ions in solution and all of the

> Ag+ (aq) + Nos- (as) + Na+ (as) + C1- (aq) ---- AgCl(s) + Na+(aq) + N03-(aq)

insoluble salts in molecular form. This is the i o n i c equation for a chemical reaction. Note that AgNOs, NaC1, and NaNG3 are

form. -indicated as ions in solution and AgCl is written in the molecular

Only those ionic and molecular species which represent the

74

observed chemical change are used to write the net ionic equation. The non-participating ions, called spectator ions, are omitted from the equation. Since Na+ and N O 3 - appear on both reactant and product side of the ionic equation and they remain in the same physical state without an observed change, they are considered spectator ions. Thus, the net ionic equation for the chemical reaction is as shown below.

Ag+(aq) + Cl-(as) - - - - > AgCl(s)

This type of chemical equation is classified as a metathesis or double displacement reaction because an exchange of ions takes place between the reactants. A requirement of this type of reactions is that a chemical change must be observed by the formation of a precipitate] a change in color, evolution of a gas, a change in acidity or basicity] or the release or absorption of heat energy. If none of these changes are observed, then the double displacement of ions to produce a molecular species does not take place. You simply have all the ions reamaining in solution unchanged.

The procedure requires you to mix pairs of salt solutions together. First, write a predicted metathesis reaction in molecular form. Use the solubility rules to help you. Second, you,will write the ionic equation designating the physical state of each reactant and product as they exist in solution. Finally, you will write the net ionic equation based on your experimental observations.

Procedure

Clean and dry 12 test tubes and label them with numbers 1 - 1 2 . Prepare 250 ml of the following solution from the solutes indicated and store them in appropriately labelled bottles. Cht-k your calculations for the amount of solute to be weighed with the instructor. Transfer 2 to 3 ml of each solution to the correspondingly labelled NHdNOa and place in test CaC03 in test tube $12.

Test tube Number 1 2 3 4 5 6 7 8 9 10 11 1 2

test tube. Obtain a few crystals of tube + I l l . Obtain a few crystals of

Solu tion or Reagent 0.1M FeCls from FeC13- 6HzO 0.1M Na2COs from NazCO3 0.1M BaClz from BaClz- 2HzO 0.1M NasPO4 from NasPO4. 1 2 H z O 0.1M NiClz from NiClz. 6HzO 0.1M CuSO4 from CuSO4- 5HzO 0.1M CoCl2 from CoClz- 6H20 0.1M AgN03 from AgN03 1.OM HC1 from concentrated HC1 1.OM NaOH from NaOH Crystals of NH4 NO3 Crystals of CaC03

75

For each pair of solutions that will be mixed:

a) write a predicted metathesis equation in molecular form

b ) observe and note any properties of the reactants

c) mix he recommended amounts of the two reactants and note any chemical reactions that occur (refer to the introduction section for chemical change information)

d) write the net ionic equation that describes the chemical react ion.

CAUTION: I f a gas is evolved, do not attempt to detect odor unless spezifically told to do so and then only in the appropriate manner.

Combine the following solutions and/or reagents in the amounts specified:

a) Combine 1/2 of #l with 1/3 of # l o b ) Combine 1/3 of #2 with 1/3 of #6 c) Combine 1/2 of # 5 with 1/3 of #2 d) Combine 1/2 of #3 with 1/3 of #6 e) Add 1/2 of # 4 to the remainder of 86 f) Add 1/2 of #7 to the remainder of #l g ) Add 1/3 of #lo to #ll; heat over cool flame and cautiously

h) Add 1/2 of #9 to #12 i) Add the remainder of #9 to the remainder of # l o ; insert

detect odor

a thermometer in #lo before addition and note any temperature change.

j) Add the remainder of # 4 to the remainder of #t5 k) Combine 1/2 of #8 to the remainder of tt7 1) Add the remainder of #8 to the remainder of #2

Reauiremnents for Report

As procedure requires, show all equations and evidence of chemical reactions from experimental observations.

Quest i ong

1. List five ways of identifying a chemical change. 2. What is the difference between an ionic and a net ionic

3 .

4 .

equation? In general, what can you conclude about most salts containing C1- and S O I - ? In general, what can you conclude about most salts containing C03-2?

76

5. If equal volumes of 1.OM N H 4 O H and 0.1M CuSOs are mixed an intensely blue colored solution results. This blue color is significantly different than the original CuSO4 solution color. a) Predict a metathesis equation. b) Check you textbook or reference source to verify or

c) Write the net ionic equation for the chemical reaction. d) How is this reaction different than the typical metathesis

disprove your prediction.

reaction.

77

STANDARDIZATION OF A BASIC SOLUTION

When a solution of strong acid is mixed with a solution of a strong base, a chemical reaction occurs that can be represented by the following net ionic equation:

> H 2 O ---- H+ + OH-

This is called a neutralization reaction and chemists use it extensively to change the acid-base properties of solutions. The equilibrium constant for the reaction is 1 0 1 4 at room temperature, so that the reaction can be considered to proceed completely to the right, using up whichever of the ions is present in the lesser amount and leaving the solution acidic or basic, depending on whether H+ or OH- ion was in excess.

Since the reaction is essentially quantitative, it can be used to determine the concentrations of acidic and/or basic Solutions. A frequently used procedure involves the titration of an acid with a base. In the titration, a basic solution is added from a buret to a measured amount of acid until the moles of OH- ion added is just equal to the number of moles of H+ ion present in the acid. At that point, the volume of base that has been added is measured.

The equivalence point or end point in the titration is determined by using a chemical, called an indicator, that changes color at the proper point. The indicators used in acid-base titrations are weak organic acids or bases that change color when neutralized. One of the most common indicators is phenolphthalein, which is colorless in acid solutions but becomes red or pink when the pH of the solution becomes 9 or higher. The point at which the phenolphthalein changes color is the endpoint of the titration. Indicators are selected so that the reaction’s equivalence point and the indicator’s endpoint occur at essentially the same pH in the titration.

In Section A of the procedure, dry potassium hydrogen phthalate, KHCeH404, is used as the primary standard acid for determining the normality of the sodium hydroxide solution. This standard is used because of its high purity and because it is not hygroscopic. The accurately weighed acid sample is dissolved in distilled or deionized water. The equivalents of KHCsH404, KHP, used for the titration is calculated from the acid’s measured mass and its equivalent weight:

mass (E) = eq KHP equivalent weight (g/eq)

The equivalent weight of K H P is 2 0 4 . 2 g/eq. Recall that an equivalent is that amount of pure substance necessary to generate one mole of H+ ion or one mole of OH- ion. The NaOH solution is then added from the buret to the KHP solution until the equivalence point is reached, signaled by the phenolphthalein indicator changing from colorless to pink. At this point the volume of NaOH solution added to the K H P solution is determined and recorded. Because one equivalent of KHP reacts with one equivalent of NaOH, the normality of the NaOH solution is calculated using the following equations:

eq KHP = eq NaOH ( 2 )

Normality (N) NaOH = ~ C I NaOH L solution ( 3 )

In Section E of the procedure the equivalent weight of an unknown acid is determined. The standardized NaOH solution f rom Section A is used to titrate an accurately measured amount of unknown acid to the equivalence point. By knowing the volume and the normality of the NaOH, the equivalents of acid in the unknown are determined using the following equations:

Volume ( L ) X Normality (N) = eq NaOH ( 4 )

eq acid = eq NaOH ( 5 )

From the measured mass of the unknown acid titrated, the equivalent weight of the acid is calculated using:

Gram Equivalent Weight (acid) = mass of acid ( E ) eq acid ( 6 )

Procedure

Section A

Obtain a clean 1500 mL flask from the stockroom, add boiling chips and more than one liter of distilled water. Heat the water to a rolling boil for 10 minutes. Cool the water in an ice bath. Weigh approximately 4 to 4 . 5 g of NaOH pellets. Obtain a one liter plastic bottle from the stockroom and clean it well with soap and water being careful to rinse it thoroughly with distilled water. Transfer some of the cooled water to the plastic bottle and add the previously weighed sodium hydroxide pellets. Shake and swirl the mixture until the pellets have dissolved. Add the remaining water to fill the container completely. Cap the container and shake it to ensure a homogeneous solution. Label the container as approximately 0.1N NaOH with your initials and ‘the date on which the solution was prepared. Be sure to shake the solution well before you use it.

79

Prepare three clean Erlenmeyer flasks and weigh three samples of KHP while at the balance. Weigh the weighing bottle containing the KHP to the nearest 0.001 g . Into each clean Erlenmeyer flask, weigh 0.4 to 0.5 g of KHP to the nearest 0.001 g . Pour the KHP samples from the weighing bottles and reweigh the weighing bottle to obtain sample weight. Be sure to mark the flasks with numbers for identification purposes. Add 50 ml of distilled water to each flask and 2 drops of phenolphthalein indicator.

Wash a 50 ml buret thoroughly with soap and water using a long buret brush. Also thoroughly clean a funnel. Flush the buret and funnel with tap water and rinse several times with small quantities of distilled water. Fill the buret with distilled water and allow it to drain. A clean buret shows no water droplets clinging to its inner walls. If the buret is glass, care must be taken to properly grease the stopcock to avoid leaks during the titration. Rinse the buret with three 5 ml portions of the NaOH solution, making certain that the NaOH wets the entire inner surface of the buret. Drain the NaOH through the buret tip.

Using the funnel, fill the buret with the NaOH solution and kecord its initial volume to the nearest 0.01 ml. Do this by holding a white card with a black line in back of the buret and reading the bottom of the meniscus. Be certain that all air bubbles are removed from the buret and buret tip. Place white paper beneath the Erlenmeyer flask. Slowly add drops of NaOH solution to the acid, swirling the flask after each drop. Note that the pink color of phenolphthalein appears and disappears as the drops hit the acid solution and are mixed with it. When the the pink color begins to persist, slow down the rate of NaOH addition. Rinse the inner wall of the flask with a wash bottle periodically during the titration. In the final stages of trie titration add NaOH drop by drop until the entire solutiori just t r i r r l q a pale pink color that persists for 30 seconds. Many times a full drop of NaOH solution added to the acid solution may be too much near the endpoint. If this is the case, then a partial drop is formed on the buret tip and washed into the flask with a wash bottle. Again the wall of the flask should be thoroughly rinsed and the solution swirled. Continue the titration until an endpoint is reached. buret and record.

Read the volume of NaOH solution in the

Refill your buret and repeat the titration at least two more times with varying but accurately known amounts of KHP. Calculate the normality of the NaOH solution based on each sample of KHP. The values should agree within one percent. If there are any discrepancies, additional samples of KHP should be weighed and titrated. Calculate the average normality for the NaOH solution clith its standard deviation. Record the normality on the label of

~ the bottle of NaOH solution. When not in use, the NaOH solution should be tightly sealed as it absorbs CO2 from the atmosphere to form an acid solution, thereby weakening the concentration of the solution. When you have completed the titration process for

8C

the day, drain the buret into the sink and rinse several times with distilled water. Fill the buret with distilled water and place in the buret stand with a small beaker below the tip. This is the proper storage procedure for a glass buret and will allow you to know when it will be necessary to regrease the stopcock due to leakage. If teflon stopcocks are used, the empty buret should be inverted with tip up and stopcock loosened slightly.

Section B

Two unknown acid samples are to be titrated. Three trials of each unknown must be performed. Weigh each unknown sample by difference as described for weighing the KHP. The sample size may vary greatly due to the wide range of gram equivalent weights found in the unknown samples. It is suggested that the student start with 0.3 g (to 0.001 g ) sample size. If after the first titration has been completed, this sample size required more than 50 ml of NaOH, reduce the sample size accordingly.

Add 50 ml of distilled water and two drops of phenolphthalein to each unknown acid sample. Fill the buret or refill the buret as the case may be and record the initial volume of standardized NaOH solution. Titrate to the phenolphthalein endpoint. Head the final level of NaOH solution in the buret and record. Repeat the titration at least twice using varying amounts of unknown acid. For each acid, calculate the average gram equivalent weight and standard deviation. Again, discard any NaOH solution remaining in the buret and clean and store the buret as previously described. Save the remaining NaOH solution in the plastic bottle for future experiments.

Requirements f o r Report

Construct a data table for each section. Section A’s table should contain the following information: weight of KHP sample, equivalents of KHP, equivalents of NaOH, total volume of NaOH solution used, normality of NaOH for each sample, the average normality of NaOH solution with the standard deviation. For Section B, each unknown acid sample must have the following information included in the data table: unknown number, sample weight, total volume of NaOH used, equivalents of NaOH, equivalents of acid, gram equivalent weight for each trial, and average gram equivalent weight of each unknown with standard deviation.

stions

1. Explain the difference between an equivalence point and an endpoint.

. 2. Is it quantitatively acceptable to titrate all KHP samples with the NaOH solution to the same dark red endpoint? Explain.

81

3 . If the endpoint in the titration is mistakenly exceeded ( t o o pink), what effect does this have on the calculated normality of the NaOH solution? Explain.

4 . ”If 2 drops are good, then 20 drops are better.” Explain why this reasoning is not acceptable when adding phenolphthalein indicator for endpoint determination.

dissolved in a 100 ml of diqtilled water. If 2 1 . 6 4 ml of the sodium hydroxide solution are required to reach the equivalence point, what is the normality of the NaOH solution?

6. A 0.394 g sample of an unknown acid requires 23.75 ml of 0.123 N NaOH to reach a phenolphthalein endpoint. What is the unknown acid’s equivalent weight?

7. If KHP were hygroscopic, how would this property tend to affect the normality values of NaOH? Explain the reasons for predicting values too high or too low for normality.

5. A 0.325 g sample of potassium hydrogen phthalate, KHP, is

F 82

VINEGAR ANALYSIS

Many commercial products contain or are low percent acid solutions. Vinegar is a water solution that is 4 to 5 percent by weight acetic acid, CHaCOOH. The minimum federal standard for all vinegars is 4%. Many manufacturers add flavorings and color to make the product sell better. The purpose of this experiment is to compare the acetic acid concentrations in various commercial vinegars and determine a best buy for the purchase price.

The percent by weight of acetic acid in vinegar can be determined by titrating a measured weight of vinegar to a phenolphthalein endpoint with a measured volume of standardized base. The number of equivalents of acetic acid can be calculated using:

Volume (L) X N ( equivalents ) = eq base L solution

A t the equivalence point, the number of equivalents of standardized base is equal to the number of equivalents of acetic acid.

The number of grams of acetic acid titrated can be determined by multiplying the number of equivalents of acetic acid by its equivalent weight, which is the same as its molecular weight since there is only one active hydrogen in CH3COOH.

eq acetic acid X ( 60.05 gram acetic acid.) = grams acetic one eq acetic acid acid (2)

The percent acetic acid in the vinegar cnn be ~ : ~ l c u l z t e ! d as

grams acetic acid X 100 = % acetic acid by weight grams of vinegar (3)

Procedure

Select a brand of vinegar from the samples available and record the brand name and purchase price. Pipet 5 ml of vinegar into a clean, dry Erlenmeyer flask, previously weighed to 0.001 g. Reweigh the flask and contents. Add 2 drops of phenolphthalein to the vinegar and dilute with 50 ml of distilled water.

Rinse a clean buret with the previously prepared standardized NaOH solution, making certain no droplets of liquid cling to the inside walls of the buret. Fill the biiret ~ I t h tkAe stadardized

. NaOH solution. Eliminate all air bubbles from the buret and tip and record the level of NaOH solution to the nearest 0.01 ml.

83

Titrate the vinegar to a phenolphthalein endpoint. Record the level of the buret at the endpoint to the nearest 0.01 ml. Repeat the procedure on a duplicate sample.

Select another brand of vinegar and perform duplicate analyses to determine the acetic acid content.

Reauirements for Report

Construct a data table containing the following information: brand name, purchase price, vinegar sample weight, total volume of standardized base, normality of standardized base, equivalents of acetic acid, weight of acetic acid, percent acetic acid in vinegar, average percent acetic acid in vinegar. From the this information determine which vinegar is the "best buy".

Questions

1.

2 .

3 .

4 .

5 .

A total of 3 4 . 2 ml of 0.105 N NaOH was required to reach a phenolphthalein endpoint in titrating 6.15 g of vinegar. a) Calculate the number of equivalents of acetic acid in the

b ) Calculate the number of grams of acetic acid present in the

c) Calculate the percent by weight of acetic acid in the

A drop of standardized NaOH solution adheres to the side of the Erlenmeyer flask and is not washed into the vinegar with the wash bottle; how does this affect the reported percent acetic acid in vinegar? Explain. In determining the percent acetic acid in vinegar, the samples are weighed rather than measured volumetrically. Explain. If the vinegar were measured volumetrically, what additional piece of data would you need to complete the calculations in this experiment? Phenolphthalein changes color in the pH range 8 . 2 to 10.0; methyl red changes from 4 . 2 to 6.0. Comment on the advisability of using methyl red indicator in this experiment and the consequences if indsed the substitution is made.

vi ne gar .

sample.

vinegar.

84

ACIDITY OF FRUIT JUICES

Some citrus fruits taste more sour, and therefore are more acidic,' than others. The purpose of this experiment is to determine the total acidity of apple, orange, lemon, grapefruit and/or lime juice. Citric acid is the acid responsible for the sour taste in citrus fruit. Citric acid is a polyprotic acid in which three ionizable hydrogens may be neutralized by a strong

H O I ti

H - C - C -OH

H O I I I

H - C - C -0-Na+

H - C - C -OH + 3 N a O H ---- > H - C - C -0-Na+ + 3 H z U

H H

base. It is not the only water soluble acid found in citrus fruits, since ascorbic acid, Vitamin C, is also contained in fruit juices. Ascorbic acid is a monoprotic acid with only one titratable

HCsH706 + NaOH ---- > NatCsH706- + H2O

hydrogen per molecule. Consequently, only total acidity can be determined on the basis of equivalents of acid per.volume of juice solution. This will allow the pH of the fruit juice to calculated

( 1 ) total eq of acid = moles of hydrogen ion in juice

by taking the moles of H+ and dividing by the volume in liters of pH = - log ( mole of H+) = - log (H+ molarity) (2)

liters of juice

the juice sample titrated with the standardized NaOH solution. The pH values for the individual fruit juices should be found in the CRC Handbook of Chemistry and,Physics (p. D151) prior to starting the experiment.

85

E rocedure

Section A

Apple cider will be used as the representative sample of apple juice. Cut the citrus fruit in half and use the juicer apparatus to obtain juice from the fruit.. Filter the juice using a Buchner funnel to remove the pulp. Be sure to use a trap bottle between the filtering flask and aspirator to avoid contaminating the juice with tap water. Into a 250 ml Erlenmeyer flask, pipet 5.00 ml of juice. If the juice is highly colored, a lesser quantity, 1.00 or 2.00 ml, may be used instead. Dilute the juice to approximately 100 ml total volume with distilled water and add two drops of phenolphthalein indicator.

Rinse and fill your buret with standardized NaOH solution from a previous experiment. Record the initial volume level of NaOH solution. Titrate the juice solution to the phenolphthalein endpoint and record the final volume level in the buret. Repeat the titration on a duplicate sample to check for reproducibility. C-lciilate the pH of the fruit juice using equations ( 1 ) and ( 2 j . kompare your results with the CRC Handbook value for the corresponding fruit juice. Choose another citrus fruit or apple juice and repeat the above procedure on duplicate samples.

Section B

Using a pH meter which has been calibrated, measure the pH of each fruit juice. Take care to rinse the electrode with distilled water between individual measurements. Record the pH values in your notebook.

Keauirements for Report

Construct a data and results table containing the following items: name of the juice, volume of the juice, normality of the NaOH, total volume of the base used in the titration, equivalents of the acid, moles of H+, molarity of H+, the calculated pH of the juice sample, measured pH value and pH range from the CRC Handbook. Comment on the consistency of the sour taste to the actual pH range found in the CRC Handbook and whether your experimental results parallel the actual values. Discuss the two methods of determining pH used in the experiment and state reasons why one method may be preferable. (Hint: Check your textbook for the buffering capacity of salts of weak acids.) Show a set of sample calculations for the experiment.

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86

Q u e s t i on s

1. Is it acceptable to report the total acidity of a juice sample as equivalents of citric acid? Explain your reasoning.

2 . Why are pH ranges for various fruit juices reported rather than exact values?

3. If a 5.00 ml sample of lemon juice require 37.78 ml of 0.125 N NaOH solution to reach a phenolphthalein endpoint; a) calculate the equivalents of acid. b ) calculate the moles of H+ ion. c) calculate the molarity of H+ in the juice sample. d ) calculate the pH of the juice sample.

a7

STANDARDIZATION OF AN ACID SOLUTION

Neutralization titrations are widely used for determining the concentration of substances that react directly or indirectly with a solvent to form hydrogen or hydroxide ions. For most applications, water is chosen as the solvent because of its availability] low cost and low toxicity. In most neutralization reactions between strong acids and bases, a pronounced pH change takes place at the equivalence point. It is for this reason that standard solutions for such titrations are always prepared from strong acids or strong bases. Hydrochloric acid is the most commonly used standard acid solution for the titration of bases. Dildte solutions of the reagent are stable indefinitely. A standard acid solution is ordinarily prepared by diluting an approximate volume of the concentrated reagent and subsequently standardizing the diluted solution against a primary standard b a s e . In this experiment] you will determine the concentration of a 0.1 M HC1 by standardizing it directly against the weak base called ?'RAM or TRIS, which means tris(hydroxymethy1)aminomethane. THAM is a primary standard base whereas KHP is a primary standard acid. The results for the HC1 molarity] from the direct standardization against THAM, will be compared to the results obtained indirectly by titration against the previously standardized NaOH solution.

The equation for the reaction of THAM with HC1 is similar to H H t

H H HN: + HC1 ---- > HN:H + c1-

that for the reaction of ammonia, NH3, with HC1. Sometimes this reaction is written in the chemically equivalent way where NH4OH

NH4OH + H+C1- ---- > NH4tCl- + HOH

is 1 ammonia molecule associated with 1 un-ionized water molecule. Note the formation of a covalent bond between the hydrogen ion and the ammonia molecule, using the lone pair of electrons on the nitrogen atom.

THAM is a derivative of NHa and a covalent bond between the hydrogen ion and a lone pair of electrons also forms during its neutralization by HC1. It is clear that there is a one to one

H + H

H H (HOCH2)aC-N: + H+C1- -e-- > (HOCH2)sC-N:H + c1-

88

mole relationship in the reaction of THAM and HC1. Knowing this, we can write the equation for the calculation of HC1 molarity f rom

Molarity (HC1) = ( Qr ams THAM) ( 1000 1 (GMW of THAM) (ml of HC1)

the titration results as indicated above. The gram molecualr weight of THAM is 121.4 g/mole.

To obtain results for the comparison value against the standardized NaOH solution, a fixed volume of acid will be pipetted into a clean Erlenmeyer flask with phenolphthalein indicator added and titrated to the normbl endpoint. Since the reaction of HC1 with NaOH is a one to one mole relationship, the

HC1 + NaOH ---- > HOH + NaCl

equation for the calculation of the molarity of the HC1 can be

Molarity (HC1) = (ml of NaOH) (molarity of NaOHl (ml of HC1) ( 2 )

written from the titration results

Procedure

Section A

Weigh four samples of THAM contained in the labelled weighing bottles, by difference, into 250 ml Erlenmeyer flasks. The THAM has been dried in an oven at 1000C for one hour to remove adsorbed moisture. The samples should weigh between 0.5 and 0.6 g to the nearest 1 mg. Dissolve each samyic in 75 1111 of distilled water. Add water to the flask from a wash bcttle to ensure that any powder adhering to the neck and walls of the flask is washed into the solution. Add 3 drops of Methyl Red Indicator ( 0.2% methyl red in ethanol) to each flask. The acid color of this indicator is light red or pink while the basic color is yellow. At the endpoint the solution will be orange.

Rinse a clean buret with three 5 nl portions of the HC1 solution to be standardized. Fill the buret with the HC1 solution taking care to remove air bubbles from t h e buret and tip. Record the initial volume of HC1 solution. Titrate the THAM solution with the acid solution to an orange colored endpoint. Using a hot plate, heat the resulting solution gently for 3 minutes to remove COz dissolved in the solution. If the solution is not boiled, you may obtain a premature endpoint due to the buffering effect of the dissolved C O z . T h i s will be pxriderlt because the 'indicator color will revert to the yellow color when the solution is heated. Heat until small bubbles appear on the bottom of the flask. Do NOT boil too long or too hard, otherwise the HC1 will be lost and the indicator will be destroyed. After cooling the flask under the tap or in an ice bath, add an extra drop or two of . - :r t,c, the sc.Jllution. If t)i? e n j i l t ; -1m 1 - , 2 = t r r r n G J tf- ;i

' yellow color, continue the titration until the orange endpoint is reached.

I

89

Finish each of the four titrations completely befclre proceeding to the next titration; this avoids errors in the buret readings for the volume of HC1 solution. Read and record the final volume of HC1 solution. If the endpoint is exceeded, a back titration with your standardized NaOH solution is possible. You must remember to adjust the equivalents of acid accordingly before calculating your acid normality.

Section B

Drain the remaining HC1 solution from the buret into a beaker and save this solution for later use. Clean the buret and fill with standardized base from a previous experiment. Add more H C 1 solution to the beaker and pipet three 25 ml aliquots into separate 250 ml flasks. Add two drops of phenolphthalein indicator to each flask. Record initial level of NaOH solution in the buret and titrate to the normal light pink endpoint. Record t h e final level of NaOH solution in the buret. Refill b u r e t w i t h NaOH solution and titrate the remaining samples completing each sample before starting the next. Record all volumes and the qormality of the NaOH solution.


Construct a data table for information obtained in the experiment and also for results from the necessary calculations in the experiment. Show one set of sample calculations for each section of the procedure. Report the average molarity of the HC1 solution with the standard deviation for both sections. Compare the molarity of the acid as determined by THAM titration to the molarity of the acid as determined by NaOH titration. Express the difference in molarities as percent difference.

Glues t i ons

1. Why is it proper to use phenolphthalein indicator for the KHP titrations with NaOH while a methyl red indicator is reqilired in the THAM titrations with HCl? Explain.

2 . If 2 3 . 5 6 ml of the hydrochl3ric acid solution is required in the titration of 0.536 g of THAM by the procedure in the experiment, what is the molarity of the acid?

a) calculate the percent error in the molarity of the acid solution which would result from an error of 0.02 ml in the titration volume.

b) calculate the percent error in the molarity of the acid solution which would result from an error of 5 mg in the weight of THAM.

3. Using the results from problem 2:

90

4 . Compare the percent errors in a) and b) of problem 3 . Using this information, state which error would have a more significant effect on your results. What measurements should you take great care when making?

5 . Why might the values of the molarity of the HC1 solution, determined by THAM, be slightly different from the values determined by NaOH? Which should be more accurate?

NaOH, is quantitative; between HC1 and a weak base, THAM, is quantitative; between NaOH and a weak base, KHP, is quantitative. Why is the reaction between THAM and KHP not quantitative?

place of THAM. In this case, the reaction is

6. The reaction between a strong acid, HC1, and a strong base,

7. Borax, Na2E407. lOHzO, may be used as a standard in

NazB.107 + 2 HC1 t 5 H 2 O ---- > 2 NaCl + 4 H3BO3

Compute the equivalent weight of borax and compsre it to THAM. Does this suggest an advantage that borax may have over THAM?

91

ANTACID ANALY S I S

Various commercial products claim to give the "best relief" for acid indigestion. Pharmaceutical companies issue quantitative results from their laboratories that verify their advertisements. Rolaids, for example, claims to "consume 47 times its own weight in excess stomach acid. This information was published in an article in Chemistry, Volume 4 4 , page 28 in 1971. The normal pH of the stomach ranges from 1.0 to 2.0. Acid indigestion and heartburn normally occur at a lower pH. The purpose of an antacid, regardless of its claims of effectiveness, is to neutralize or to buffer the excess hydrogen ion in the stomach and, therefore, relieve acid indigestion. Milk of magnesia, an aqueous suspension of magnesium hydroxide, Mg(OH)2, is a simple antacid that

Mg(0H)z + 2H+ ---- > Mg+z + 2H2O ( 1 )

neutralizes H+. The "fast relief" antacids that buffer the excess acid in the stomach are those that contain calcium rqrhonate, CaC03, or sodium bicarbonate, NaHC03. A HC03-, C O 3 - 2 buffer system is established in the stomach with these antacids as shown in the following equations. Rolaids, which

contains dihydroxylaluminum sodium carbonate, is a combination antacid, that is, it both neutralizes and buffers. It reacts with a c i d according to the following equation:

NaAl(OH)2COa + 3H+ ---- > Na+ + Al+3 + 2HzO + HC03- ( 4 )

This experiment determines the total effectiveness of several antacids by means of a strong acid-strong base titration. TG avoid the possibility of a buffer system being established during the titration, an excess of hydrochloric acid will be added to the dissolved antacid forcing the equilibrium to the right (Refer to equations 3 and 4 above). The solution will be heated to drive off the carbon dioxide as a gas. The excess hydrochloric acid will be titrated with standard NaOH solution.

The number of equivalents of antacid in the commercial sample plus the number of equivalents of NaOH used in the titration

equivalents of base = equivalents of acid (5)

eqantacid + eqNaOH = eqHCl ( 6 )

92

equals the number of equivalents of H C 1 added. Therefore, the equivalents of antacid can be obtained by solving equation ( 6 ) . The number of equivalents of antacid per gram allows the comparison of effectiveness of the various brands.

Procedure

Weigh one tablet of antacid and record the weight and brand name. Grind the tablet using a mortar and pestle. Record the cost and number of tablets per bottle for each brand which you analyze. Weigh two clean, dry 250 ml flasks and mark each for identification purposes. Transfer one half of the pulverized antacid tablet, approximately 0.7 g, to each flask and reweigh the flasks to the nearest 1 mg. Pipet 60.0 ml of 0.1N HC1, previously standardized, into each flask and s w i r l to dissolve. The solution may appear cloudy due to the insoluble binder, an inert ingredient, used to prepare the tablet. Record the normality of the HC1. Heat the solution to a gentle boil for 3 minute to remove any dissolved C02. Add 4 - 8 drops of bromophenol blue indicator. Bromophenol blue indicator is yellow in acidic solution and blue in basic solution. If the solution is blue, pipet an additicnal 10.0 ml of the 0.1N HC1 solution into the sample flask and boil again. Cover the flask with a watch glass or Parafilm and place in an ice bath to cool to room temperature. Rinse a clean buret several times with small quantities of the standardized NaOH solution before filling. Record the normality of the NaOH solution and the initial volume of the NaOH to the nearest 0.01 ml. Be sure there are no air bubbles in the buret tip. Titrate the acidic solution with the standardized NaOH solution to a blue endpoint. Record the final volume of NaOH in the buret to the nearest 0.01 ml. Repeat the titration with s second sample of the same brand.

Choose a second brand of tablet and repeat the procedure described above. Compare the strengths of antacid tablets on the equivalents per gram basis. Calculate the cost of antacid tablets per gram so that the "best buy" can be determined.

Requirements for Report,

Construct a data table containing the following pieces of information: brand name, weight of sample, normality of HC1 solution Used, volume of H C 1 added, equivalents of acid added, normality of NaOH solution used, total volume of NaOH, equivalents of NaOH added, equivalents of acid in excess, equivalents of antacid in the tablet, equivalents of antacid per gram, cost of antacid per gram, and equivalents per cent. Show one complete set of calculations for a'sample. On the basis of information collected, compare the brands, and rate the best buy for effectiveness.

93

Quest i on s

1 .

2.

3 .

4 .

5 .

Write the balanced equation for the reaction of 1 mole of the active ingredient in Rolaids with excess H+ ion. If the pH of the stomach is 1.0 and the volume of the stomach is one liter, how many grams of milk of magnesia will be required to raise the pH of the stomach to 3.0? Salts of weak acids can also produce a neutralizing effect in the stomach by reacting with water, this is called hydrolysis. Write a balanced equation for sodium citrate, Na3CsH507, whose parent acid is citric acid, H 3 C s H s 0 7 . If the C 0 2 is not removed from the solution by boiling after the 0.1N H C 1 is added, how will this affect the amount of NaOH required to reach the bromophenol blue endpoint? If results from trials on the same sample differ by a substantial amount, more than 5%, what should you do before presenting your results?

94

ANALYSIS OF ACTIVE INGREDIENT IN COMMERCIAL TABLETS

Many of the commercial pharmaceutical products have weak. acids as active ingredients which are titrable using a standardized base solution and an appropriate indicator such as phenolphthalein. In this experiment, two products, Vitamin C and aspirin, will be investigated for activity. Each tablet will contain a specified amount of active ingredient as indicated by the label on the bottle. After titration with the standardized base solution, the experimental value can be calculated and compared

% activity = E a ctive ingredient (exp eriment) X 100 g active ingredient (reported) (1)

to the reported value on the bottle to verify the manufacturers’ claims. Both acids are monoprotic so calculations can be done using molarity or normality.

Aspirin, acetylsalicylic acid, is both an organic ester and an organic acid. It is used extensively in medicine as a pain killer and fever reducing drug. When ingested, the acetylsalicylic acid remains intact in the acidic stomach, but in the basic medium of the upper intestinal tract, it hydrolyzes forming the salicylate and the acetate ions. Aspirin’s analgesic action is probably due to the salicylate ion, however, its additional and physiological effects are still not totally understood. It is known that when salicylic acid is ingested, the same therapeutic effects are observed, however, it causes an upset stomach whereas aspirin is less l i k e l y tc do so. The quantitative reaction of acetylsalicylic acid with a standardized

ii 0 II

+ OH- - - - -> allrcH’ + H z O O-C-CH3

C-OH It 0

II 0

sodium hydroxide solution occurs according to the following reaction. The volume of base times the normality yields equivalents of base which are equal t-o equivalents of acid. The grams of acetylsalicylic acid may be calculated using the following relationship.

eq acetylsalicylic acid X 1 8 0 . 2 a = g acetylsalicylic ( 2 ) acid eq

95

Vitamins are a group of organic substances required in the diet of man and animals. Vitamin C , ascorbic acid, can be obtained from citrus fruits, tomatoes, potatoes, and fresh vegetables. All animals, except man, are capable of synthesizing their own ascorbic acid; man must obtain it from food sources. The long established use of citrus fruits, particularly limes, to prevent the occurrence of scurvy was readily justified when it was established that the condition was a result of a Vitamin C deficiency. A well-balanced diet provides adequate amounts of the vitamin as measured by the Recommended Daily Allowance (RDA) of 75 mg per day.

Ascorbic acid is a cheap, water soluble, organic compound which has a molecular formula, C s H 8 O 6 . It is a weak monoprotic acid and may be titrated with standard NaOH solution to a

phenolphthalein endpoint. Ascorbic acid oxidizes readily in air, particularly at high temperatures, and therefore the vitamin is easily destroyed in cooking and long storage. Similar to the calculations for aspirin, when the equivalents of ascorbic acid are known, the grams of ascorbic acid may be calculated from the following relationship.

eq ascorbic acid X 1 7 6 . 2 ~ = grams ascorbic acid eq (3)

Procedure

Obtain three tablets of each type from the instructor and grind each tablet using a mortar and pestle. Transfer this powder to a 250 ml Erlenmeyer flask and be sure to wash the residue of each tablet into the flask with a wash bottle. Check the bottle of both the aspirin and vitamin C and record the potency of each tablet as rated by the manufacturer. Add 50 ml of distilled water to each sample and swirl to dissolve the tablet. The solution may have a cloudy appearance which is the result of the binder, an inert ingredient, used to produce the tablet. Add two drops of phenolphthalein to each flask. Rinse a clean buret with several small quantities of the standardized NaOH solution. Fill the buret with the NaOH solution and record the inital volume to the nearest 0.01 ml. Titrate to a light pink endpoint and record the final volume of the NaOH solution.. Repeat the last two steps on each sample. Record the normality of the NaOH solution used in the titration process.

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96


Construct a data table for information collected on each sample and the results of each tablet should include equivalents of acid, grams of acid, and percent activity. Show a sample set of calculations necessary to reach the final result.

Questions

1. If 6 fluid ounces of a well known.vegetable juice contains 35% of the RDA of vitamin C, how many milliliters are necessary to provide the recommended daily allowance of 7 5 mg?

than fresh vegetables. 2. Explain why cooked vegetables have lower vitamin C content

3. What reaction occurs when aspirin enters the intestinal tract? 4 . What is the active ingredient in aspirin and why it is not

5. A 325 mg tablet of aspirin was titrated to a phenolphthalein ingested directly?

endpoint with 16.3 ml of 0.107N NaOH. What is its percent activity?

97

IRON - COPPER SULFATE REACTION

Most chemical reactions are classified as being either acid- base or oxidation-reduction, redox. Acid-base reactions produce water as a reaction product. In redox reactions, electrons are transferred from one substance to another, resulting in a change in oxidation number. The substance achieving a higher oxidation number by losing electrons is oxidized. The substance achieving a lower oxidation number by gaining electrons is reduced. Electrons are never considered as reactants or products. The total number of electrons gained must equal the total number lost. For example, magnesium metal reacts with oxygen gas to form magnesium

Mg(s) + 0 2 ( g ) ---- > MgO( s ) (unbalanced)

oxide. Magnesium increases its oxidation number from 0 to +2 by l o s i n g 2 electrons per mole of Mg, that is, it is oxidized.

Mg ---- > Mg+z + 2e- (an oxidation half-reaction)

Ehch oxygen atom in the 0 2 molecule decreases its oxidation number from 0 to -2. Each mole of 0 2 must gain 4 moles of electrons to

0 2 + 4e- ---- > 20-2 (a reduction half-reaction)

to be reduced to two moles of 0-2. The half-reactions show what is taking place with each element in the chemical reaction. To yield the balanced equation, the number of electrons lost in the oxidation half reaction must equal the number of electrons gained in the reduction half-reaction. Balancing f o r electrons gained and

2 (Mg ---- > Mg+2 + 2e-) 1 (02 + 4e- ---- > 3.0-2 1

4e- + 2Mg + 02 ---- > 2Mg+2 + 20-2 + 4e- 2Mg + 02 ---- > 2Mg0

lost and adding the two half-reaction together produces the balanced redox reaction. In this reaction, Mg, in causing 0 2 to accept its electrons, is the reducing agent and 02, causing Mg to release its electrons, is the oxidizing agent. A reducing agent causes reduction or is oxidized itself and an oxidizing agent causes oxidation or is reduced itself.

In this experiment, you will perform a simple redox reaction by allowing iron filings to react with a copper sulfate solution. An excess of copper sulfate will be used to ensure that all of the iron will react. By weighing the amount of iron used and the amount of new substance formed, you will be able to determine the quantitative relationship between reactants and products for this reaction. It is a redox reaction, therefore, you can verify your results by half-reactions for the reactants to produce a balanced

. redox reaction.

98

Procedure

On a weighing paper or boat, measure approximately 1 gram of iron filings to the nearest .001 g . In a clean 100 ml beaker, weigh approximately 5 grams of copper sulfate to the nearest .1 g. Add 25 ml of distilled water to the salt in the beaker. Heat the solution until it just starts to boil. DO NOT let the solution boil over. Remove the heat source when the solution is hot. While stirring the hot copper sulfate solution with a glass rod, add iron filings in small amounts until all has been added to the solution. Reweigh paper or boat after addition to obtain an accurate mass of iron added. Record the weight and any observations made during the reaction.

Weigh a filter paper and record its weight to the nearest 0.001 g. Filter the solution by gravity, carefully decanting the solution and transfer the solid with the aid of a wash bottle and rubber policeman. Wash solid several times with 10 ml portions of distilled water. Allow the solid and paper to dry until the next laboratory period and weigh it to obtain the weight of the product.

Reauirements for Rep ort

Construct a data table for the information collected during the experiment. The table should include the following items: moles of iron used, moles of copper produced based on product weight, mole ratio of Fe/Cu using the appropriate number of significant digits.

Questions

1. Write a balanced redox equation using half-reactions for the

2. What evidence did you observe to indicate that some of the

3 . What evidence other than product formation would indicate

4 . Starting with 0.500 g of iron fillings, how many moles of

5 . If 1.545 grams of copper are produced from the reaction, how

chemical reaction which took place.

copper solution remained unchanged?

that a chemical reaction is taking place?

copper can be formed? H o w many grams of copper are formed?

many grams of iron fillings were needed to start with?

1 99

REDOX TITFWTION WITH POTASSIUM PERMANGANATE

Reactions in which substances undergo changes in oxidation number are referred to as oxidation-reduction reactions or redox reactions. Oxidation is defined as an algebriac increase in oxidation number, or a process in which electrons are lost. Reduction is defined as an algebraic decrease in oxidation number] or a process in which electrons are gained. Oxidation-reduction processes must occur simultaneously.

The species that gains electrons is called the oxidizing agent, therefore] it is reduced. The species that loses electrons is called the reducing agent, therefore, it is oxidized. One gram equivalent weight, GEW, of oxidizing agent is the weight that gains 6.02 x 1023 electrons and one gram equivalent weight of the reducing agent is the weight that loses 6.02 X 1023 electrons. Note that the term gram equivalent weight is defined so that one GEW of oxidizing agent reacts with one GEW of reducing agent.

c-

GEWox = GEWred

Consider the reaction of potassium permanganate with oxalic acid in the presence of excess sulfuric acid. The balanced molecular and net ionic equation are:

2KMnO4 + 5HzC204 + 3H2SO4 ---- > lOC02 + 2MnS04 + K2SO4 + 8H2O

2Mn04- + 5HzC204 + 6H+ ---- > 10C02 + 2Mn+2 + 8 H z O

In WnO4 (or in the Mn04-) the oxidation state of Mn is +7, while in MnSO4 (or in the Mn+2 ion) it is + 2 . Therefore, each Mn undergoes a change in oxidation number of five. Since each formula unit of WnO4 contains one Mn, and each Mn gains five electrons, one mole of WnO4 is five gram equivalent weights in this reaction. Consequently, WnO4 produces 5 moles of electrons per mole of WnO4 or has five equivalents per mole of KMnO4.

GEW of WnO4 = 158.0 R X 1 mole = 31.60 R mole 5 eq eq

Therefore] the gram equivalent weight of WnO4 in this reaction is 31.60 grams.

100

In H2C204 the oxidation state of carbon is + 3 while in COz it is + 4 . Each carbon undergoes a change in oxidation number of one. Since each formula unit of H2C204 contains two carbons, and each carbon loses one electron, one mole of oxalic acid is two gram equivalent weights in this reaction. Consequently, H2C204 produces 2 moles of electrons per mole of oxalic acid or has 2

GEW of HzC204 = 90.0 g X 1 mole = 4 5 . 0 4 1 mole 2 eq eq

equivalents per mole of oxalic acid. The gram equivalent weight A 2 C a C 4 is 45.0 grams. The simple rule for for obtaining the

grax equivalent weight of an oxidizing or reducing agent is divide the weight of one mole of the substance by the change in oxidation number per mole of substance.

In this experiment, you will prepare an approximately 0.1N KMn04 solution and standardize this solution by titrating against a standard solution of H2C204 solution. You will then use the

unknown oxalic acid solution. Just as in acid-base titrations where the number of equivalents of acid must be equal to the number of equivalents of base, in redox titrations the number of

C. - standardized KMnO4 solution to determine the concentration of an

eq of oxidizing agent = eq of reducing agent

equivalents of oxidizing agent must be equal to the number of equivalents of reducing agent. For this reaction of KMnO4 with

eq of Wn04 = eq H2C204

HzC204. Alternatively, we can express this relationship as

litersoxid X Noxid = litersted X N r , d

At the end of a titration, three of the four variables will be known and the unknown variable can be calculated. In the standardization process for potassium permanganate, the normality will be the unknown variable. After the standardization process has been completed, the oxalic acid normality will be the unknown variable.

Procedure

The preparation of a stable solution of KMnO4 can be time consuming process. Traces of MnO2 and many other substances, as well as exposure to light, catalyze the decomposition of KMnO4 solutions. It is possible to make a WnO4 solution by diluting a concentrated WnO4 solution which has been p r t p u a d well in advance. The diluted solution should be standardized and used for analysis within a re1 ativelv s hort per iod of time in order to obtain good results. Alternatively, calculate the weight of

10 1

KMn04 required to prepare 250 ml of a 0.1N WnO4 solution. Check your calculations with your instructor. Weigh the required amount KMn04 and transfer it to a dark brown bottle. Using your graduated cylinder add distilled water to give an approximate total volume of 250 ml. Mix the solution thoroughly by vigorous swirling and shaking. This solution must be used the day it is preuared.

Assemble a clean buret and rinse it three times with 5 ml portions of your KMn04 solution, allowing the rinse solution to drain through the tip of the buret each time. Discard the rinse solutions. Fill the buret with KMnO4 solution and allow it to drain through the buret tip until no air bubbles remain in the tip. Record the buret reading before beginning the standardization. Because the KMn04 solution is opaque, the buret is read at the top of the meniscus.

Standardize your Wn04 solution by titrating it against the standard k c 2 0 4 of known concentration indicated on the bottle label. Pipet separate 10.00 ml samples of the standard oxalic acid

25 ml of distilled water to each flask. In the fume hood, cautiously add 10 ml of 1 2 M H2SO4 solution to each sample using a stirring rod. CAUTION: &SO4 is very corrosive and should be handled with great care. Check MSDS notebook for the precautions and spill cleanup method and record in your notebook. Using a hot plats, heat the samples to 80 - 9OoC and titrate the hot solutions with the WnO4 solution. The WnO4 solution must be added very slowly initially. If it is added rapidly, brown MnOa often precipitates and the titration must be discarded. The reaction of WnO4 and HzC204 is apparently catalyzed by the presence of Mn+2 ions. After a few milliliters of WnO4 solution has been added, the rate of addition can be increased. As the endpoint is approached, the Wn04 solution should again be added slowly as in any titration. Since KMn04 is so intensely colored, it serves as its own indicator. The endpoint is the point at which a faint pink coloration persists throughout the stirred solution for at least thirty seconds. When you think you have reached the endpoint, read the buret and record the volume. Then add one more drop of m 0 4 . If an intense pink color is obtained, this indicates that the endpoint was reached before the last drop was added. Calculate the normality of the KMnO4 solution and repeat the titration until you obtain at least three values which agree within 2 2 in the third significant digit. Determine the average normality with the standard deviation. Use the average normality for the remainder of calculations in the experiment.

c- - solution into three 250 ml Erlenmeyer flasks. Add approximately

Pipet 10.00 ml samples of the unknown H2C204 solution into clean 250 ml flasks and treat these samples in the same manner as the oxalic acid solutions used in the standardization process. Follow the directions precisely. When you have finished the titrations, calculate the normality of the unknown oxalic acid sample, determine the average and the standard deviation.

102

When you have finished, discard the Wn04 solution in the lab sink with large amounts of tap water. Add 10 ml of 3% HzOz arid 2 ml of 12M HzSO4 to a beaker that contains 50 ml of distilled water. Use this solution to rinse all of your equipment that the WnO4 solution has touched because it leaves a thin film of MnO2 on glass. The acidic hydrogen peroxide solution will remove this residue. Then clean all glassware with soap and water using distilled water for the final rinse.


Construct a data table for the information collected during the experiment. A results table should include the following items: normality of standardized KMn04 for each sample, its average and standard deviation, normality of unknown oxalic acid solution for each sample, its average and standard deviation, and a sample calculation for each result.

Quest i ons

C. - 1. What weight of KMn04 would be required to prepare 300 ml of 0.150N KMnOs to be used in this experiment? What would be the molarity of this solution?

HZC204. 2 H z O would be required to prepare 1.50 liters of the standard oxalic acid solution?

standard for Wn04 solutions. However, since As406 is insoluble in H20, it is usually dissolved in NaOH solution,

2 . A standard oxalic acid solution is 0.300N. What weight of

3. Tetraarsenic hexoxide, AS406, is frequently used as a primary

4NaOH + AS406 ---- > 4NaAsOz + 2HzO

and the resulting solution is acidified with hydrochloric acid

NaAsOz + HC1 ---- > HAsOz + NaCl

and then titrated with WnO4 solution. The above reactions do not involve oxidation-reduction and serve only to prepare the solution for titration with KMnO4 solution.

a. Balance the following equation.

KMnO4 + HAsOz + HzO + HC1 ---- > KC1 + MnC12 + H3As04

b. A 0.04852 g sample of As406 was dissolved in NaOH solution, acidified with HC1 solution, and required 2 4 . 7 5 ml of KMnO4 solution for complete oxidation. Calculate the normality of the WnO4 solution.

4 . Great care must be exercised in the standardization process to avoid formation of MnO2 in the early stages of the titration. Explain why the sample must be discarded if this occurs. Use balanced equations to aid your explanation.

103

COLLOIDS

A colloid is a very fine dispersion of one substance in another, in which the substance is not soluble. For example, a starch-water mixture forms a colloidal dispersion in which solid starch is dispersed in water. A colloidal dispersion is like a true solution since the dispersed particles can not be removed by filtering, however, unlike a true solution all the particles carry the same electric charge. The principal distinction between true solutions, colloids and suspensions is particle size. The particle size in true solutions is less than 10 nanometers (nm) while colloid particles range from 10 to 1000 nm. Neither solutes or colloid particles are visible with an ordinary microscope, as are the larger particles found in suspensions. Colloidal solutions do scatter light at right angles to the light beam and, in addition, the size, shape and general behavior of these particles determine the intensity and color of the scattered light. The rapid, irregular movement of colloidal particles through the dispersion is called Brownian motion.

There are many kinds of colloidal systems, two common types being liquid dispersed in liquid called emulsions and solid dispersed in liquid called a s o l . The sols are divided into two large classes; lyophilic sols in which there is an attraction between the colloidal particles and the solvent and lyophobic sols in which there is little or no attraction for the solvent. Lyophilic s o l s tend to form gels or semirigid liquids, while lyophobic sols are usually quite fluid.

Colloidal systems may be prepared by condensation which _. “’-ol LA,* ,-=es building up molecules or ions to colloidil size o r b y dispersion which involves breaking down larger particles into clolloidal size. Colloids may be separated from ions or molecules in true solution by a process known as dialysis.

In a given colloid the particles usually have a charge, either positive or negative, due to ions adsorbed on the surface of the particles. Since the particles all adsorb the same charge, the colloidal dispersion is more stable because like charges repel. If the charge on the particles is removed or neutralized, they may form larger aggregrates and separate from the solution. The colloidal particles can be precipitated from solution, by adding an oppositely charged ion that has an affinity for the adsorbed ion, by adding an ion of different charge, or by mixing colloids of opposite charge.

In this experiment, the student will form a variety of colloidal systems and investigate their individual properties

1 104

Procedure

Colloidal Fe(OH13

add slowly, drop by drop, a solution of 0.5M FeCla until a highly colored liquid is obtained. This will require approximately 5 ml of the iron solution. Note the color of the liquid. Assuming complete hydrolysis of the FeC13, write the equation. Set this preparation aside f o r later use.

Heat 50 ml of distilled water to boiling and while boiling

Carbon Black S o l Put the amount of carbon black that can be held on the tip of

a spatula into a mortar, add 2 ml of water and some tannin. Grind until the substances are finely mixed. Add 50 ml of distilled water and set aside for later use. What is the purpose of the tannin?

Gal

and ad3 50 ml of reagent alcohol, mixing the solution thoroughly.

beaker. Cut a piece of this material, place it in an evaporating dish. In the fume hood try igniting this material with a match and record the results.

Foam Dissolve about 5 g of Alz(S04)3. 18Hz0 in 40 ml of hot water.

Add about 0.5 g of detergent. In another beaker dissolve about 5 g of NaHC03 in 50 ml of water. Mix 5 ml of each solution in a 50 ml graduated cylinder and record the results. Explain the results observed. In a laboratory sink, mix the remaining portions of the solutions in a large beaker.

Place 10 ml of saturated calcium acetate solution in a beaker

c- - Describe what happens. After 5 minutes elapses, invert the

Emulsion

To one of the test tubes add 4ml of soap solution. Shake both test tubes vigorously and record the separation times. E x p l a i n results.

Fill two test tubes with 2 ml of kerosene and 10 ml of water.

Fi 1 t rat i on

the same f o r the carbon black sol and the emulsion. Record your results.

Take some of the colloidal Fe(OH)3 and try to filter it. Dc

Dialysis

second test place a corresponding amount of 0.1M CuC12. Cover the top of each test tube with a piece of plastic wrap or other semipermeable membrane which should be held in place by a strong rubber band. Invert the test tubes, placing each in a separate beaker of water and allow to stand for 30 minutes. Have either of the two substances diffused out? Explain the difference.

Fill a test tube halfway with the colloidal Fe(OH)3 and in a

10 5

Tyndall Effect

narrow beam of strong light. Do likewise for some of the CuClz solution. Explain the results. Place the Fe(OH)3 solution against a black background and view the reflected light perpendicular to the light beam.

Take some of the Fe(OH)3 solution and hold it in a sharp

Coagulation

in test tubes or small beakers and label them 1 to 5. Into the first tube, from a dropper, add 0.1M KC1 counting the drops and stirring until the colloid precipitates. Into the second tube add in the same manner 0.1M KzSOs counting drops. To the third tube add in the same manner as before drops of 0.1M K 3 P 0 4 . To the fourth tube add in the same manner drops of 0.1M B a C 1 2 . To the fifth tube add some 0.1M sugar solution.

Divide the remaining Fe(OH)3 solution into five equal parts

Colloids, such as Fe(OH)a, are stabilized by electrical charges. Higher charged ions of opposite sign are more effective in neutralizing the charge and causing the colloid to coagulate. Tn view of these facts, explain why different amounts of the various solutions were needed to.coagulate the colloid. On the basis of your results, what is the charge on the Fe(OH)3 colloidal solution? Why was the BaClz solution not as effective as the K z S 0 4 solution since each contain bivalent ions? Why was the sugar solution not effective?

c- -

Requirements f o r Report

Write a brief summary of the observations recorded in the Be sure to include explanations where required and procedure.

answer all questions asked in the procedure.

Chi e 3 t ‘Lon s

1. How is a colloidal dispersion like a true solution? 2 . How is a colloidal dispersion different than a true solution? 3 . HOW does a colloidal dispersion acquire an electrical charge? 4 . How does a suspension differ from a colloidal dispersion?

106

COLLIGATIVE PROPERTIES

The addition of a nonvolatile solute to a solvent produces characteristic changes in the physical properties of the solvent. For example, when salt is used to freeze ice cream, the salt-ice- water mixture's temperature drops below the temperature observed for an ice-water mixture itself. Also we know that when antifreeze, ethylene glycol, added to an automobile's cooling system prevents freezing in the winter and boiling during the summer because the antifreeze solution has a lower freezing point and a higher boiling point than pure water.

The figure below is a plot of the vapor pressure of pure water, solid line, and the vapor pressure of a water solution containing a dissolved nonvolatile solute, dashed line. Pure

water boils at 1OUoC when its vapor pressure is 760 torr. When a nonvolatile solute is added forming a solution, solute molecules or ions occupy part of its surface area. This inhibits the water molecule's movement into the vapor state causing a vapor pressure lowering of water, a, under the 760 torr value. Since the solution's vapor pressure is less than 760 torr, it does not boil at 1 O O O C . For the aqueous solution to boil, its vapor pressure must be raised to 769 torr by increasing the solution's temperature above 1 0 0 c C . "hi: is t h e boiling point

C. "

c

0' t T m r t u r e I-Cl 1W t

elevation, b, of the water due to the solute. The freezing point depression, c, results from the solute molecules inhibiting the formation of the crystal lattice necessary for water to solidify. Water's normal freezing point is OoC, however, the freezing poir?t of an aqueous solution containing a nonvolatile solute is lowered.

Vapor pressure lowering, boiling point elevation, and freezing point depression are colligative properties of solutions. These properties are determined by the number, rather than the type, of solute particles dissolved in the solvent. For example, an equal number of moles of glucose or urea lower the freezing point of a fixed amount of water to the same temperature. One mole of NaCl and 2 moles of glucose in the same amount of water produce the same effect on the colligative properties of the solution because one mole of NaCl prodaces Lwt ; m ? e s of solute particles, Na+ and C1-, whereas the 2 moles of glucose remain unionized.

107

The change in the colligative proerties of an aqueous solution is directly proportional to the amount of solute dissolved in the solvent. For water, one mole of a nonvolatile, nonelectrolyte dissolved in 1 kilogram of water, a 1 molal solution, produces a solution with a boiling point elevation of 0.51oC and a freezing point depression of 1,860C. The freezing point depression, A T f , and the boiling point elevation, A T b , are proportional to the molality, m, of the solute in the solution.

ATi = Kfm

ATb = Kbm

m = moles of solute = jtz/ GFW) solute kg solvent kg solvent

Kf and Kb are molal freezing point and boiling point constants for the solvent.

c- In this experiment you will determine the gram formula weight

of a nonvolatile solute in tertiary butanol by measuring the freezing point for pure t-butanol and for the solution containing a mass of solute, The mass of the solvent and solute must be determined by weighing. The molal freezing point constant] Kf , is S.loC-kg/molal and pure t-butanol's freezing point is 2 5 . 5 0 C . Using the above equations you will calculate the moles of solute and also the solute's gram formula weight.

The freezing point of t-butanol and that for the solution will be obtained from a cooling curve, a graph of decreasing temperature as a function of time. The figure below illustrates that

1 108

t-butanol’s cooling curve achieves a plateau at its freezing point. Extrapolation of the plateau to the temperature axis determines its freezing point. The student should be aware that supercooling may occur and consequently, the temperature should be monitored for at least ten minutes after it appears that a plateau has been reached. The solution’s cooling curve does not achieve a plateau, but continues to decrease slowly as the t-butanol freezes. Its freezing point is determined at the intersection of two lines drawn tangent to the curves above and below the freezing point.

Procedure

Assemble the apparatus shown in the figure at the left. Weigh a clean, dry 200 mm test tube in a 250 ml beaker to the nearest 0.01 g. Add approximately 10 - 1 5 g of t-butanol to the test tube and reweigh the tube and contents to the nearest 0.01 g. Record both measurements. Prepare about 300 ml of an ice-water mixture in a 400 ml beaker. A 600 ml insulating beaker which may be filled with cold water, near OoC, to minimize heat absorption for the surroundings by the ice-water mixture in the 400 ml beaker.

hole rubber stopper using glycerin as a lubricant. Carefully remove all

“ O Y Insert a 1100C thermometer in a two

traces of glycerin from the thermometer. Insert the wire stirrer in the other hole before stoppering the test tube.

Heat the test tube and contents to 450C by inserting it in a beaker of warm water. Place the test tube and contents in the ice-water bath and while stirring with the wire, record temperatures to the nearest 0.loC at 1 minute intervals. The temperature remains constant at the freezing point until solidification is almost complete. Continue recording until the temperature begins to drop again. The freezing point determination may be repeated by simply immersing the test tube and its contents in the warm water beaker again and repeating the temperature-time measurements. On graph paper, plot the temperature, vertical axis, versus time, horizontal axis, to obtain t-butanol’s cooling curve.

Melt the t-butanol in the warm water bath to remove any of the solvent from the thermometer and wire stirrer. Remove the stopper from the test tube and dry the outside of the test tube and reweigh the test tube and contents in the 250 ml beaker to the nearest 0.01 g. Obtain an unknown from the instructor in a

/

weighing boat. Quantitati rely transfer the 1 nknc tube and reweigh the tube to the nearest 0.01 g.

109

Jn to the test Record the

weight measurements. If the unknown adheres to the sides of the test tube roll the solvent in the test tube until the solute dissolves. The wire stirrer may be used to aid in forming an uniform solution. Replace the stopper in the test tube and warm the solution to 4 5 o C hy placing it in a beaker of warm water.

Place the test tube in the ice-water bath and while stirring with the wire, record temperatures to the nearest 0.loC at one minute intervals. A break will occur when the solution nears its freezing point and the time interval should be shortened to 30 seconds. This break may not be as pronounced as the plateau was for the pure solvent. Plot the temperature versus time on the same graph as the cooling curve for pure t-butanol. Draw the t.angent lines to the curve above and below the freezing point. The intersection of the tangent lines is the solution’s freezing p o i n t .

Warm the test tube and its contents in the warm water beaker. Remove the stopper and wire stirrer. Dry the outside of the test tube and reweigh in the 250 ml beaker. Obtain from your instructor an additional amount of the same unknown sample and quantitatively transfer to the test tube. Reweigh the tube and contents. Record these measurements. Repeat the freezing point determination and plot the data on the same graph. After all experimental data has been completed, dispose of the contents in the test tube by adding a warm soap water solution to the tube and discard in the laboratory sink with large amounts of tap water.


The graph of experimental data must be clearly l z k c l l c d LZCI freezing points of the solutions and t-butanol indicated on the temperature axis. All weight measurements and results should be tabulated. A sample calculation for the gram formula weight of the solute should be demonstrated.

110

Questions

1 .

2 .

3 .

4 .

5 .

A 0.578 g sample of a nonvolatile solute dissolves in 15.0 g of t-butanol. The solution freezes at 2 3 . 0 o C . What is the gram formula weight of the solute? If ten grams of t-butanol and the same mass of solute were used in the above problem, would you expect the freezing point depression to be larger or smaller than that observed for the solution in the above problem. Explain. If the solution’s freezing point is mistakenly read 0 . 2 0 C lower than it should be, wiil the gram formula weight be too high or too low? Explain. If the t-butanol is initially contaminated with a nonvolatile, nonreactive nonelectrolyte, how does this affect the reported gram formula weight? If the thermometer is miscalibrated 0 . 5 0 C higher than the actual temperature over its entire scale, how will it affect t h e reported gram formula weight of the solute? E x p l a i n .

111

CALORIMETRY

In nearly all chemical and physical changes, a transfer of heat energy is observed. This heat flow for a chemical reaction, measured in a calorimeter, is quantitatively expressed as the heat of reactions or enthalpy of reaction, AH, at constant pressure and temperature. Reactions evolving heat are exothermic reactions and their AH values are negative. Reactions absorbing heat are endothermic reactions and their AH values are positive.

In this experiment three calorimetric determinations will be conducted. First, a calorimeter constant will be determined for the calorimeter used in subsequent measurements of enthalpy of reaztions. Second, a heat of neutralization, A h , for a strong acid/strong base reaction will be determined. Third, the heat of solution, A H s , will be measured for a salt dissolving in water.

Calorimeter Constant

c- When a chemical process is carried out in a calorimeter, some of the heat transferred during the reaction is absorbed by the environment. For example, hot coffee in a Styrofoam cup loses heat to the cup and surrounding air. In other cases, some heat is released by the calorimeter to the chemical system. An example of this would be ice melting in a styrofoain cup since it would absorb heat from the cup and immediate environment. Heat energy always flows from a hot system to a cold system. The calorimeter constant will be measured by adding warm water to cold water contained in the calorimeter. Since the heat energy which is lost by the hot system must equal the heat energy which is gained by the

Qhot = Q c o l d + Q c a l

cold system the calorimeter and surroundings must be considered as part of the cold system. If the heat transferred to the calorimeter and surroundings is divided by the experimental temperature change, the calories transferred to the calorimeter and surroundings per degree of temperature change is referred to as the calorimeter constant. The number of calories per degree will be essentially the constant for any similar reaction run in that particular calorimeter if the temperature change and the volume of the final solution is approximately the same as that used in the determination of the calorimeter constant.

The calorimeter constant can be calculated using the data from the mixing of hot and cold water in the calorimeter.

where Qcal is the quantity of heat transferred to the calorimeter and surroundings, Qhot is the quantity of heat transferred from the hot water, and Q c o l d is the quantity of heat transferred to

I 112

the cold water. Equation 3 allows the calculation of heat

Q = m ( g ) X Cp (cal/g-oC) X AT ( 0 C ) (3)

transferred by the water where m is the mass of water, Cp is the heat capacity of water, and AT is the temperature change of water. By substitution into Equation 2 , the following Equation 4

results. Rearrangement of Equation 4 allows the calculation of

the calorimeter constant where A T c a l is the temperature change of t h e calorimeter.

Heat of Neutralization

The reaction of a strong acid with a strong base produces C.

H+ + OH- - - - > Hz0 + heat

water and heat. The heat evolved during a reaction taking place

Q r x n = Q i q + Qcal (6)

in a calorimeter must be the sum of heat absorbed by the liquid reaction mixture and the heat absorbed by the calorimeter. The density and heat capacity of combined acid-base solution is assumed to be equal to those of water. The temperature change is measured for the reaction mixture and the same temperazure change Is assumed for the calorimeter. When th;= :31c:r'-es p r ~ d u c e d for

Q r x n = (mCpAT)liq + ( C c a l A T c a l ) ( 7 )

this reaction have been calculated from Equation 7, it can be converted to the enthalpy of neutralization by dividing by the number of moles of water produced. The acid will be the limiting

AHneut = Q r x n moles of HzO

reactant for the neutralization reaction. The enthalpy of reaction can then be compared to the expected AH for the reaction of HC1 with NaOH.

Heat approximately 200 ml of distilled water to 4 0 0 C on a hot plate. Support the calorimeter, the two cup assembly, in a 250 ml beaker. Place 50 ml of room temperature distilled water in the calorimeter. Suspend a thermometer, using a clamp and ring stand, so that the bulb of the thermometer is immersed in the water in the calorimeter. The thermometer should not touch the sides or the bottom of the calorimeter and should be capable of being read easily. Weigh a 100 ml beaker to the nearest 0.01 g. Place approximately 50 ml of the warm distilled water in a 100 ml beaker and reweigh. Record the weight of warm water. Suspend a thermometer in the warm water in similar fashion as described above. F o r a 5 minute period, measure and record the temperatures at 30 second intervals.

While stirring the sample of water in the calorimeter, rapidly p o u r the hot water in the beaker into the cold water. To e n s u r e total transfer of the hot water reweigh the 100 ml beaker st the erla of trial. Record time of mixing and continue to measure and record temperature-time data for a 15 minute period at one minute intervals. Construct a graph where the temperature is the vertical axis and time is the horizontal axis. Plot the temperature-time data on the graph. Draw a vertical line on the graph to represent the time of mixing. The best straight line curve is drawn through the plotted points prior to the time of mixing. These curves are extrapolated so that they intersect the line representing the time of mixing. The best straight lina curve is drawn through the plotted points after the time of mixing. This line is extrapolated so that it intersects the line of mixing. The point of intersection represents the maximum temperature achieved by the mixture in the calorimeter. The difference between the initial temperature of the cold water and

c-

38

37

36 u

28

2 8

26

1 114

Heat of Solution

The lattice energy of a s a l t , A H ~ ~ , and the hydration energy, AHh, of its composite ions govern the quantity of heat evolved or absorbed when it dissolves in water. Lattice energy, an endothermic quantity, is the energy required to vaporize 1 mole of the salt into gaseous ions. Hydration energy, an exothermic quantity, is the energy released when gaseous ions are attracted and surrounded by water molecules as a solution forms. The heat of solution, H s , is the difference between these two energy

AHLE = A h + AHs

values and may be either an exothermic or endothermic process depending on the magnitude of the lattice energy and hydration energy for a particular salt.

In this experiment the heat transferred by the dissolution of N H 4 C 1 can be calculated by using equation 9 . The heat transfer

for the calorimeter, Q c a l , is dependent on the temperature change observed in the solution. The heat transfer of solution, C a s o l , is determined by total mass of solution, combined salt and water,

Q s o l = (mass of solution) (heat capacity of solution) (AT) times the heat capacity of the solution times the temperature change observed. To find the enthalpy of solution, A H s , the

AHe = QS

moles of N H 4 C 1

heat transferred, Q s , is divided by the number of moles of N H 4 C 1 dissolved in the water. This value can be compared with a literature value.

Procedure

Calorimeter Constant

Obtain two 6 ounce Styrofoam cups and lid from instructor. Place one cup inside the other. Before starting measurements of the calorimeter constant, 150 ml of hot distilled water should be placed in the inner cup and allowed to stand for several minutes. This process should be repeated several times to remove any residual chemicals remaining after the manufacturing process for the cup. If this is not done, difficulty may be encountered in obtaining a consistent calorimeter constant.

115

the maximum temperature of the mixture is the temperature change for the cold water. In similar fashion, the temperature change for the hot water can be found. Calculate the calorimeter constant as described in the introduction.

Repeat the entire procedure and determine the second calorimeter constant. If values agree within 5%, proceed to the next section and if they do not, repeat the procedure until agreement is obtained.

Heat of Neutralization

Measure 50 ml of 1.00 M HC1 at room temperature into a clean, dry calorimeter. Suspend a thermometer in the solution in the same manner as described in the first section of the procedure. Weigh a clean, dry 109 ml beaker. Measure 50 ml of 1 . 0 5 M NaOH at room temperature into the 100 ml beaker and reweigh. Record b o t h weight measuremeiits. Suspend a thermometer in the N a U i i solution. Adjust the temperature of the NaOH solution by heating or cooling until the temperature is the same a s the HC1 solution.

f o r the t w o solutions at 1 minute intervals. --. F a r a 5 minute period, measure and record temperature-time data

While stirring the HC1 solution in the calorimeter with a g l a s s stirring rod, pour rapidly the NaOH solution into the HC1 solution. Record the time of mixing and continue to measure and record temperature-time data for 15 minutes at one minute intervals. Construct a graph similar to that described in the previous section of the procedure. Plot the temperature-time data and draw the best straight line for piotted points before mixing and the best straight line for the plotted points after m l x i r i g . Extrapolate both lines to intersect the vertical line representing the time of mixing. From the graph deLerrriiiie t i l e -axi-miirri temperature at mixing and calculate the temperature chs.nge for the solution. Reweigh the 100 ml beaker which contained the NaOH solution. Determine the heat transferred and the enthalpy of neutralization using the equations found in the introduction section. Repeat the entire procedure for a second determination. Compare average enthalpy value to literature value.

Heat of Solution

Measure 2 . 5 to 2 . 8 g of ammonium chloride in a weighing boat. Record mass of boat and ammonium chloride. Place 100 ml of room temperature distilled water in a clean, dry calorimeter. Suspend a thermometer in the calorimeter as previously described. For a 5 minute period, record temperature of the distilled water at 30 a e ~ o n d intervals. While stirring with a glass stirring rod, rapidly add the ammonium chloride and save boat to reweigh later. Note the time of mixing and continue to measure and record temperature-time data for a 15 minute period at 30 second intervals. Construct a graph of the data with temperature on the

I 116

ordinate and time on the abscissa. The same procedure should be followed to obtain the maximum temperature change by extrapolation. The heat capacity for the ammonium chloride solution is 0.966 cal/(g-oC). Remember the mass of the solution is the mass of water plus the mass of salt used. Reweigh the weighing boat to determine the mass of ammonium chloride added to the water. Refer to the equations in the introductory section to calculate heat transferred, and the enthalpy of solution. Repeat for a second determination and compare the average value to the literature value found for N H 4 C 1 .


Report must include all graphs and calculations for heat transfer, experimental enthalpy values, literature enthalpy values, and relative error percent on the comparison.

c-

3 117

RATES OF REACTION

A. Factors that affect reaction rates.

The study of chemical reaction rates, mechanisms ~7f reaction and control is called chemical kinetics.

The rate of a chemical reaction is defined as the change in reactants (or products) concentration divided by the time required for this change. The unit of time may be seconds, minutes, hours, days or years. Rates are also, affected by five factors which are listed as follows: (1) Nature of the reactants, i.e., once metal may react vigorously with acid while another does not react. (2) The particle size of the reactants, i.e., a lump of coal b u r n s slowly but powdered coal may explode. ( 3 ) Temperature increases in general increase the rate of reaction, i.e., a 1 O o C rise in -__ .temperature doubles the reaction rate. ( 4 ) Catalysts affect the rate by using or allowing a different pathway for the reaction to follow. There still is a great deal of things to learn in this topic. (5) Concentration affects the rate of reaction, i.e., if the concentration of one of the reactants is doubled and is an integral part of the reaction the rate increases appropriately.

chemical reaction they must be isolated one from another. To study the factors that affect the rate of a

Procedure

[l] Nature of the reactants: (a) Into 4 separate 150 mm testtubes place 1 ml each of 3M s u l f u l i ,

acid, 6M hydrochloric acid, 6M nitric and 6M phosporic acid. [Review the handling procedures for acids.]

testtube and record your observations.

acid. To one tube add a piece of zinc metal, to the next add a piece of lead and to the third add a piece of copper. Record your observations.

Next place a small piece of polished magnesium strip into each

(b) Into 3 separate 150 mm testtubes place 2 ml of 6M hydrochloric

[2] Size of the reactants: (a) Prepare a 250 ml side arm flask with a thistle tube and a one

hole stopper. Connect to the side arm a rubber tube with a piece of bent glass tube in the other end. Arrange a beaker of water with an inverted 250 mm test tube filled with water so as to collect evolved gas.(See page 131 in CTRRH) Next place about 3 g of marble chips into the flask and 50 ml of water. Insert the stopper with the thistle tube into the neck of the flask and make sure the end of the thistle tube extends below the water level in the flask. Now slowly add 10 ml of 6M HC1 through the thistle tube. Record your observations.

replace the marble chips with 3 g of ground up chips or with powdered calcium carhnnPte Repeat the p r c , c + d l l r + a q hef nre s n d compare the results with the first trial.

(b) Clean the generator flask and set it up again as before, but

1 118

[3]Temperature: (a) Some reactions lend themselves to rate determinations. The

reaction of HCl and thiosulfate is such a reaction:

HC1 + NazS203 = S + SO2 + NaCl + H2O

the formation of the sulfur is time related. In three testtubes place 5.0 ml of 0 . 1 M thiosulfate and in a second set of three test tubes place 5.0 ml of 0.1M H C 1 . Now taking one set of the tubes each; place them in (1) an ice water bath, (2) a hot water bath at 7 0 O C , ( 3 ) a room temperature setting. Starting with the room temperature pair, pour the acid into the thiosulfate solution and stopper and mix thoroughly far a few seconds. Record the mixing time and measure the time until the first cloudiness appears.

Next repeat the process with the tubes in the icebath. Again record the time elapsed from mixing to the appearance of the sulfur.

Last mix the heated solutions and replace them into the warm water and record the time. When all the data has been collected plot the

C. - temperature on the y-axis and the time on the x-axis. Discuss the graph. (b)Another reaction theat exhibits the affects of temperature is

the reaction of oxalic acid with permanganate.

5 H z C z 0 4 + 2KMnO4 + 3HzSO4 = 1OCOz + 2MnSO4 + 8 H z O + KzSOs

Place 10 ml fo 0.3M oxalic acid in a 10 ml Graduated cylinder and 2.0 m l of 0.01M permanganate solution in a 250" testtube with 8.0 ml of 3M sulfuric acid. When one student mixes the solutions the other should record the time and ther temperature. When the purple color disappears the time is recorded. Repeat the experiment at 40oC allowing some time for the solutions to reach the temperature of the water. Again pour the oxalic acid into the permanganate and record the temperature and time. Do the same thing at 800C and record the data. Plot the temperature versus time as before and discuss the graph.

decomposition exhibits the property of this factor.

provided for the reaction, In a second tube place another 5 ml of the peroxide solution and a pinch of manganese dioxide. Observe and compare the two solutions. Try warming the tubes in a water bath.

[5] The affects of concentration on reaction rate: A reaction called the time clock or iodine clock reaction gives a good answer to the affect of concentration on rate.

[ 4 ] Presence of a catalyst: The affect of a catalyst on peroxide

Procedure: In a 250" testtube place 5 ml of peroxide solution

2HI03 + 5HzSO3 = I2 + 5H2SO4 + H2O

I2 + starch = 12-starch (complex) blue

Procedure

Into a 150 ml beaker labeled ( A ) place the following, 3.0 ml of 0.1M iodate, 5 ml starch and 92.0 ml of distilled water. T - i t C a n o t h e r 150 ml beaker labeled (E] place 10.0 ml of 0.1M ,bisulfite ion plus 90.0 ml of distilled water.

3 119

Next place a sheet of white paper on the desktop and on top of that p u t a clean dry 250 ml beaker. Now while one student holds the flasks labeled A and B, another will mark the time of mixing and the running time of the reaction. The first student on the signal of the student doing the timing will pour the two beakers contents together into the 250 ml beaker. When the b l u e complex just shows the time is complete. Repeat the experiment, but this time adjust the concentration of the (A) beaker to 6.0 ml of iodate, 5 ml of starch and 89.0 ml of water. The solution in (B) is held constant.

With a clean dry 250 ml beaker repeat the above procedure. Thirdly, adjust the solution ( A ) to 10.0 ml of 0.1M iodate, 5 ml

of starch and 58.0 ml of water. Again repeat the prior procedure. When the three trials are complete plot the concentration of the

iodate versus time in seconds. Remember the solutions were diluted.

Requirements for report

report. Follow the established guidelines for the writng of the

c- - 0 1 e s t,i ons

1. What factors effect the rate most markedly? Why?

2. In the time clock experiment would the acid concentration be imp0 rt ant?

[Reference:Fundementals of Chemistry by Brady and Holum, published by John Wiley and Sons]

1 120

B. DETERMINATION OF A RATE CONSTANT

To study the properties of hydrolysis of an organic molecule and use these properties to gain knowledge of the rate of a chemical reaction.

on the factors that affect the rate, i.e., nature of the reactants, temperature, concentration of the reactants, catalysts and physical state and size. For reactions in solvents the nature of the solvent interaction must also be included as a factor that may affect the rate of hydrolysis of the organic molecule in its conversion to the product that will be studied.

be reacting this material with a mixed solvent to moderate the -- - speed of the reaction. The following is a representative example of the reaction equation.

The rate of a homogeneous reaction is dependant

We will be utilizing a molecule called an organic halide and will

RX + H 2 O = ROH + H+ + X-

‘The term RX represents the organic molecule and some examples for this are t-amyl chloride, t-butyl chloride or t-butyl bromide. ROH the reaction product formed after the hydrolysis of the halide. The important part of the products side of the reaction is the generation of the hydrogen ion which will allow monitoring of the reaction progress.

The rate of the reaction is dependant on the first step of the reaction; HC1 == R+ + C1- which is slow in comparison to the rest of the reaction steps. This step is designated as the rate determining step because the sequential steps occur quickly once the carbon cation(carb0nium ion) is formed. The resulting reactions may take one of the following paths;

(1) R+ + H2O = ROH + H+

(2) R+ + HOR = ROR + H+

( 3 ) R+ = R’C=CH2 + H+ In each case a mole of hydrogen ion is produced per mole of

reactant so our monitoring of the reaction is independant of the pathway after the cation formation.

and in equation(1) with water the solvent properties must be considered in all portions of the reaction or we lose overall knowledge of which factors are affecting our rate.

order reaction, i.e., the rate is proportional to the concentration of the organic halide to the first power.

Since in equation(2) a reaction with the alcohol solvent is noted

This reaction, the hydrolysis of the organic halide is a first

121

This basically is true unless there is competition with the solvent in the first step of the reaction. The hydrolysis is therefore adjusted by maintaining a constant composition of a mix of water and an alcohol such as isopropyl alcohol.

order rate its concentration decreased at an exponential rate thus if No is the initial concentration then at some later time (t) Nt is the concentration. It can be shown then that

where k is the specific rate constant and equals the k in the rate equation. The above equation can be rearranged to the following;

When a substance is being consumed by means of a reaction of first

Nt = No e-kt

2.303log(No/Nt ) = kt

If we know the initial concentration of the reactant No and measure the reactant Nt at measured time intervals we can graph the log(No/Nt) versus time and for a first order reaction the slope of the line will equal k/2.303. Since it can be shown that the units of concentration cancel in the ratio No/Nt it is not

reaLLdrit(vO1umes) at time zero and at time t or Ao and At.

with a solvent mixture of 60% isopropyl alcohol and 40% water by volume. A solution of 50:50 IPA and water is about a factor of 4 times faster than the 60:40 mix. The method is fairly easy fot the determination. A measured volume of the halide is added to a volume of the specified solvent. Then 10 to 15 drops of phenolphthalein is added and finally a starting volume sodium hydroxide made up in the same solvent mixture is added. The amount of NaOH may be from 1 to 5 ml. The flask is stoppered and shaken a i d dhen the indicator color fades the time is taken. Another Fgrtion of base is added, shaken and the time to fading is noted.

is set aside until the following lab period. At that time the final titration is done to determine the value of A o .

volume of titrant) in the numerator and Ao minus the volume at time (t) in the denominator. The logio of the ratio is taken and plotted against the time for that particular titrations fade time.

c_ necessary to use concentration but only the total amounts of

A measureable rate of hydrolysis for t-butyl chloride is obtained

This process continues throughout the lab period and then the mix

The calculations for the plot then are simply the value Ao(tota1

Procedure

Prepare enough of the solvent mixture(60:40,50:50) to allow for t w o trials of 100 ml each and 100 ml for the preparation of the titrating base.

solvent mixture and set it aside for filling your buret when you begin the experiment.

the solvent and add 10 drops of the indicator. Place the flask into a prepared water bath at the temperature you will use for the experiment. Most will start with a room temperature trial. Allow ten minutes for equilibration and then the flask to the HOOD

A b - qrganic halides are to be dispensed. Pipet 1 ml of the ~ ~ i \ i ~ u e ir11;o the flask and mix well and mark this as the starting time.

Next weigh about 0.40 g of NaOH and dissolve it in 100 ml of the

Now using a clean, dry Erlenmeyer 250 ml flask pour in 100 ml of

,-, r ~-

122

Now using the base solution that has been placed in the buret add on to two milliliters to the flask and measure the time until the indicator turns colorless. Mark this time for that base addition. Add another portion of base and continue the process until the time of the lab period has expired. Then place the stoppered flask in your locker until the next lab period where the final titration will be performed.

The experiment will also be done at a different temperature setting. The directions are the same except that the flask will be kept in a temperature bath during all the tiration activities. DATA: When all the data has been collected for all trials you will begin the calculations to generate your plotting points for the graph.

The value called Ao will be defined as the total volume used f o r the titrations through the final titration done the following l a b period. The time at the final titration is called infinite.

the difference between A0 and the total titration volume at time (t). For example if AO is 6 9 . 7 0 ml and the volume after 125 sec is 3 . 6 7 ml then At = 6 9 . 7 0 - 3 . 6 7 or 66 .03 ml. This is then divided into Ao and the loglo is taken and plotted against the value for time in seconds or minutes.

The values for the concentration A t are determined by taking

Reauirements for report Follow the guidelines established to write your report and in

the conclusion discuss the effects of temperature on your results

Questions

1. What other rate affecting properties could be studied with this react ion? 2. Does the volume of water versus alcohol raally effect the rate? 3. What other solvent system could be s t u d i e d ?

[Reference:Laboratory Studies and Problems in General Chemistry by Stone and McCullough published by McGraw-Hill]

1 123

C. DETERMINATION OF THE ORDER OF A RATE EQUATION

The utilization of our knowledge of the rate law for a chemical

In a reaction the rate equation;

is a function of the concentrations of the participating species raised to a power relating the order to which that reactant affects the rate. For example, if the concentration of A is doubled with everything else held constant and the rate is increased by a factor of four, the value for a is two( a=2). Likewise, if the concentration of B is changed and the other factors are held constant and there is a one to one change t h e n b is one (b = 1) and is first order. The specific rate constant k o n l y varies with temperature and is unaffected by concentration changes.

In this experiment we will determine from the reaction kinetics of [S2O8-2] and [KI] the rate law and the order with respect to each reactant.

reaction can be used to determine the order of reaction.

rate = k[A]a [B]b

rate = k[S208-2]x [KI]Y

c- " We will accomplish this investigation by a variation of concentration of each of the two reactants versus time. The results are then plotted as a function of the l o g rate versus log [ 3 . This method allows us to

and holding

where log k becomes

take the rate equation: rate = k[A]aC(constant)

one reactant constant and taking the log, log(rate) = log k + a log[A] + log C

and log C are constants for this reaction and the equation

where m is the lope of the plotted line and the value for th order of the reaction. Treating both reactants this way we can then substitute the values into the rate equation and solve for k.

formation of I2 with time, so we will use a constant volume of sodium thiosulfate to give a time interval before the formation of the free Iz. The reaction equations are as follows:

The experiment requires an intermediate to help measure the rate of

S208-2 + 21- = 2S01-2 + I2

I2 + starch = 12- starch( complex)

Procedure

To summarize the experiment a table will give the composition of the five test solutions we will use to calculate the values for the order of the reaction for the two reactants.

124

! 10 ml ! 2 ml 1 10 ml I 1 30 ml d 30 ml

1 ! 148 ml 2 ml ! 10 ml

, I 30 ml 2 ! 133 ml ! 20 ml ! 3 ! 128 m l ! 30 ml ! 2 ml ! 10 ml

5 1 88 ml ! 30 ml ! 2 ml ! 10 ml 1 , 70 ml 4 1

I 1

I I

1 1

1 1

I t 50 m 108 ml ! 30. 1 ml ! ml !

-Prepare solution A for trial 1 in a clean 250 ml beaker. Mix the solution and record the temperature.

-Next measure 30 ml of 0.1 M persulfate into a clean 100 ml beaker and record the temperature.

-Now add solution B to A and start timing the reaction with a stopwatch. The appearance of the blue end point is sudden so place the 250 ml beaker on a white background for easier observance of the end point.

bcth time and temperature.

to give us the needed information for our rate law determination.

amount of iodine(1z) that is formed in each trial. The moles of iodine is exactly one half of the moles of thiosulfate. This value is then divided by the time interval for the specific trial and the l o g of this value is taken. This is the first half of the information needed to plot the graph of l o g rate versus l o g [ 3 .

volume of the mixture of the two solutions is 200 ml. If for trial one we use the concentration of the 10 ml of 0.3 M KI diluted to 20U ml, we then get the value of 0.015M KI. Now we take the l o g of this value and plot it with the first value.

c- - -Repeat the above procedure for the next four test solutions recording

Calculations The data from the above experiment must now be adjusted

-First the concentration of the thiosulfate is used to determine the

The concentration of the iodide ion is next determined. The total

This is continued for the first three trials. Another graph of the trials 3 through 5 is done the same except the

When the order of the reaction is determined for each of the concentration of the persulfate is substituted for the iodide.

participants, then the value for k is found by simply substituting the concentrations of iodide and persulfate raised to their respective powers and the value for the rate at that trial and just calculate.

7 0 0 Follow the established guidelines outlined for report writing.

1. What is the effect on the experiment when the starch solution is old? 2. Would more starch increase the rate of the reaction?

[Reference:Fundementals of Chemistry by Brady and Holum published by John Wiley and Sons]

125

D. DETERMINATION BY SPECTROPHOTOMETRIC STUDY THE RATE OF BROMINATION OF ACETONE

We will utilize our knowledge of the rate law and the factors of the visible spectrum to study the kinetics of the bromination of acetone in the presence of an acid media.

A and B are concentration dependant factors. The values of x and y are dependant on their affect on the rate of the reaction and likewise the value for k is dependant on the rate law and temperature. With careful planning each of the nonconcentration dependant variables can be determined by varying the concentration of the reactants while holding other parameters in the reaction constant. This reaction is a case in point. We will be doing a study of the affect of three reacting species by varying the concentration of one while we hold the other two constant at constant temperature. Our reaction is

In the rate equation; rate = k[Alx[B]~ the values of

H3CCOCH3 + Brz + H+ = H3CCOCH2Br + HBr c- .

With three variables, the logistics become rather formidable but we are aware of the fact that the order of the reaction with respect to bromine is zero; and, therefore, its concentration has no affect on the rate of the reaction. We will test this fact but needless to say it reduces the work, needed to complete the study of this reaction. Another factor we can study in this reaction is the value called the activation energy (Ea). This is found by the following equation:

log k = -Ea/(2.303RT) + const.

if we take the value for log k at two different temperatures an;! plot L L V C ~ ~ U S the value 1/T, the slope of the line is equal tc - E a / / 2 . 2 2 2 E .

One other factor we need to discuss is the method of following the FrDgress of the reaction. Since neither acetone nor hydrogen ion is colored, they don’t absorb light, but bromine absorbs light until it is converted to the bromide ion. We can monitor the progress by the slow disappearance of the bromine color. To do this, we must use an instrument called a spectrophotometer which has the ability to distinguish between small concentrations in the material being analyzed. For a given analyte, the instrument at a specified wavelength will respond according to the following equation:

Absorbance(A) = a [Brz]

At 390 nm of wavelength bromine absorbs strongly and therefore we will use this as our selected wavelength. We will run a series of different concentration of the bromine to determine the value of (a) a constant of absorbance for bromine. Once this value is known, we can compute the concenterations of the bromine from the absorbance data and use this in our plots of conc changes versus time to determine the rate of the experiment. By varying one concentration, we can determine the order due to that species. When this is determined the value for k can be determined; and, thus, the value for the activation energy of the reaction is found.

126

Procedure

Prepare the following solutions in volumes of lOOml for the experimental procedures: 4M acetone solution is made by diluting 2 9 . 5 ml of acetone with water to a volume of 100 ml total volume, 1M HC1 solution, and from the instructor get 100 ml of 0.02M bromine-water solution.(This must be handled very carefully and protective gloves and eyewear is mandatory.)

Determination of the absorbance constant is done by testing various concentration of bromine solutions made by pipeting 10 ml of bromine stock into a fifty ml volumetric flask and then adding 10 ml of HC1 stock and diluting to the mark. Next pipet 5 ml of bromine and follow the above directions, repeat with 2 ml of bromine stock. The instructions for the use of the spectrophotometer will be given separately by the instructor.

values from the following table:

T r i a l N u m b e r Volume of Volume of Volume of Volume of

Determination of the rate and rate law constants. We will use the

water ( 4 M ) ( 1 M ) ( 0 . 0 2 M ) (ml>

1 1 0 . 0 10.0 1 0 . 0 2 0 . 0 2 1 0 . 0 10 .0 1 0 . 0 2 0 . 0 3 5 . 0 10 .0 1 0 . 0 2 5 . 0 4 1 0 . 0 5 . 0 10 .0 2 5 . 0

1 0 . 0 5 - 10.0 10.0 2 0 . 0

acetone HC1 bromine

c- - (ml) (ml) (ml)

Step 1: Pipet 10.0 ml of the acetone into a clean 125 Erlenmeyer flask; and, to this, add 10.0 ml of acid and 20.0 ml of distilled water. In a similar flask, pipet 10.0 ml of the bromine solution. Measure the temperatures of the two solutions being sure that they are within 0.5 degrees of one another.

Make sure your machine has been setup for the measurements and then carefully mix the two flasks by pouring the solutions back and forth a few times. This is the starting time for the reaction. Carefully plac=c some of the solution into the cuvette for the instrument and begin monitoring the experiment by taking readings periodically on one minute intervals. If the solution is being tested at a temperature other than room temperature, it will be necessary to remove the cuvette from the machine and replace it in the temperature bath being used between measurements. Make between five and ten readings and then start Trial 2 and repeat for the rest of the trials. You will be required to run the above experiment at another temperature to obtain necessary information for the activation energy determination. When all data has been collected you may proceed with the calculations in the next section.

Calculations: -The spectrophotometer constant (a) is determined by using the concentration [Brz] for each of the trials in the first step of the experiment. This value is then divided into the absorbance reading for that concentration and the three values are then averaged.

concentration of the bromine[Brz] as a function of time. Rate measurement and rate law calculations. In each trial determine Lh6

[Brz 3 = Absorbance/a

127

k41~lw g#il:lt t h e cun lzun t r a t i an versus time. tiring the S i o p U ; $ram t h i s piiZlt. and equating it to -rate you can evaluate the order of the reaction with respect to each of the participating reactants. The following sequence will allow for calcualtions of each of the species.

rate for trial 1 {I acetonel}x rate for trial 3 = { acetone]}^

rate for trial 5 r B r z y rate for trial 1 = [Br2]Y

rate for trial 1 = rH+lz rate for trial 4 [ H + l z

Once the exponents for each participant is determined the value for k at the temperature of the reaction is found by the following equation;

k = - slope/[Alx[H+l~

Use the concentrations for trial 1 for each temperature determination. When the values for k at the two temperatures are found, then the log

k is plotted versus 1/T and the slope of the line is equal to -Ea/2.303R. l A A G uLcs:ptable value to use for R is 1.99cal/mole-oK. -- -- -


the l a b discussion. Follow the established guidelines for report writing as presented in

Que s t i ons 1. What is the acceptable value in the literature for Ea? 2. Write out the chemical equations and suggest the mechanisn? for the reaction.(How does it work?)

rR+ference:Modern Laboratory Programs i n Chemistry by W . Roy Mason published by Willard Grant Press]

b 128

EQUILIBRIUM-LeCHATELIER'S PRINCIPLE

To study equilibrium and the effects of changes on that equilibrium. We will look at liquid and solid systems of equilibrium and the things that influence these systems.

Our studies, to this time, assume the completion of a reaction when we place the ingrediants in contact with one another. This assumption is true only in a relatively small number of reactions. Therefore, the fact that most reactions do not go to completion implies that we have a system where reactants are mixed with products. When a reaction system reaches the point where no further reaction is apparent, we have a state of "dynamic equilibrium". This is a property of a reversible reaction and at a given temperature can be equated to a constant value called the equilibrium constant(Keq).

of the reactants or products, we should find some evidence of a change in the system. i.e. 2NO + 0 2 = 2H02 and if we add NO to the system there should be a shift to the right in the equation indicated by an intensifying of the brown color of the N O z . This effect was first proposed by Henri LeChatelier in 1888 and is now called the LeChatelier

-Principle . The principle states that if a system in dynamic equilibrium is effected by an external stress, the system will shift the equilibrium in the direction that minimizes the effect of that stress.

If our system is in equilibrium and we change the concentration of one

Procedure

In this experiment we will look at a series of equilibrium situations and the effect of a stress(temp, conc., etc.) applied to the equilibrium. From our observations we will gain understanding crf the reaction equation and learn to predict the outcome of stresses applied to other equilibria.

a 1.U1-l NH3 solution. Add NH3 until a defiriite coior change occurs in f 5 e test tube. Note the results. Write the equ?+,icn of the reaction.

write the equation.

sulfuric acid, shake the tube and observe the results. Write an equation to explain what you see. Repeat the procedure with the nickel- ammonia solution but use 1.OM HC1. Again write the equations to explain the results.

A . To 5 ml of a 0.1M aqueous solution of Cu+2, we will add dropwise

Repeat the experiment with 0.1M Ni+z solution. Note these results and

N u w to the tube containing the copper-ammonia add dropwise 1.0 M

B. To a 5 ml quantity of 0.01M silver nitrate in a test tube add 5 ml of 0.1M sodium carbonate. Note the reaction. Write an equation to explain what you see.

to settle and carefully pour off all but about 4 to 5 m l of the clear liquid. To the test tube, add concentrated NH3(use the EOOD) with an eyedropper until the solution just disscllvqs. Ncw rarefully add 6M HN03. What happens? Write an equation. Add an excess of NH3 dropwise, what happens?

Now add 5 ml of 0.1M HC1 and describe what happens. Allow the solution

129

Now add 5 ml of 0.1M KI to the test tube- what happens? Write a sequence of equations to explain the results.

C. To 5 ml of 0.1M acetic acid add a few drops of universal indicator. Note the color and then add 0.1 g of solid sodium acetate. When it

Separate the solution into two equal quantities in test tubes and add a dissolves compare the color of the indicator.

few ml of water to each. Add several drops more of the indicator and add the same amount of indicator to two other tubes with distilled water.

NaOH. Note the colors for the specific tubes. To the other t w o tubes, add 1.0 ml of 0.1M HC1 and again note the colors for the specific t u b e s . Explain the differences. Look up the definition of the word buffer.

To one of the first tubes and one of the water tubes add 1.0 ml of 0.1M

Requirements for report Follow extablished guidelines for report writing as presented in

lab discussion and include all pertinant equations to exnlain what is observed in the laboratory experiments.

&a s t i ons

1. Define hydrolysis on the basis of your experience in this laboratory. 2. Find and write out the Henderson-Hasselbach equation.

[Reference:Laboratory Manual for General Chemistry by Bcran and Brady published by John Wiley and Sons]

130

SPECTROPHOTOMETRIC DETERMINATION OF AN EQUILIBRIUM CONSTANT

To use the knowledge of the spectrophotometer to study and

Many metal ions interact with nonmetallic ions or to determine the equilibrium constant of a complex-ion reaction.

molecules to form a relatively stable complex as we found in our study of Cu+2 + 4NH3 = Cu(NHs)4+2 complex in the previous experiment.

This property is also found with ferric ion (Fe+3) and thiocyanate ion (SCN-). Both the ferric ion and the thiocyanate ion are colorless, but the complex forms an intense red-brown color which will allow us to use the spectrophotometer. The reaction follows:

Fe+3 + SCN- = [Fe(SCN)+21 (colorless) (colorless) (red-brown)

The spectrophotometer will allow us to determine the concentration of the complex for the given reaction mixtures. With the concentration v a l u e of the complex, the equation Keq=[Fe(SCN)+21/[Fe+3 ICSCN- 1 and the beginning concentration of the ferric ions and thiocyanzte ions, it is possible to determine each concentration and then

- -calculate the L q value for a variety of the concentration conditions. For example if we define the starting quantities:

[Fe+3]init = known initial conc Fe+3 [SCN-]init = known initial conc SCN-

[Fe(SCN)+z]final= measured final conc of complex then

[Fe+3]final = [Fe+a]init - [Fe(SCN)+2]final [SCN-]final = [SCN-]init - [Fe(SCN)+z]final

and our calculations reduce to

We will make our measurements for a series of five different concentration mixtures and evaluate each and average the results to obtain a statistical value representing our work. Since we are presuming that Keq is a constant, our values should fall within a set of parameters consistent with experimental error.

To effectively analyze each mixture for the complex ion concentration, we will determine the absorbance of a set of standards prepared from our stock solution and use this to prepare a calibration curve. This can then be used to determine the final concentration of the complex at any point in our series of measurements.

Procedur e

The wavelength we will use for the spectrophotometer readings is 447 nm. This wavelength is sensitive to the complex ion color and will give good response to small changes in concentration.

iicid solvent to keep the ferric ion from any hydrolysis reactions with water that will interfere with our desired reaction. The concentration -' - i i r stock ferric ion will be 0.2M and the thiocyanate will be u . d02M.

The stock solutions for the ferric ion must be made in a 0.05M nitric

I 131

MaLerials Needed: 2 5 ml flasks .,urct pipets 0.u5M nitric acid (HNO3) 0.2M ferric ion stock (Fe(N03 )3 ) 0.002M thiocyanate ion stock (KSCN) spe.2trophotometer cuvette

I-.

Using the following chart, prepare the standard solutions f o r the calibration curve which will be used for the rest of the experiment. Proceed to make the standards by measuring the quantites using a buret and a pipet. The total volume of solution is 25 ml so use an apprjpriate flask.

Solution 0.00200M KSCN 0.200M Fe(NCh 13 1 5 . 0 ml 5 . 0 ml 2 4.0 ml 5.0 ml 3 3.0 ml 5.0 ml 4 2.0 ml 5.0 ml 5 1.0 ml 5.0 ml

-- -

Now set the wavelength on the spectrophotometer as suggested in the beginning of this section and prepare to measure the absorbance of each of the solutions. Use the 0.050M nitric acid solution as the blank or reference and starting with solution # 5 , rinse the cuvette three times with the solution and then wipe the surface carefully after filling with the solution and take the reading. Repeat with solution # 4 and the rest up the line until you finish with #l. Now plot the absorbanc2 for each solution versus concentration of the complex [Fe( SCN)+Z] which in essence is equal to the concentration of the thiocyanate ion since we ndve greatly over matched the ferric ion concentration to t i le t L : r r ~ : . - ~ ~ . i t e . Remember these solutions were diluted to 25 ml for t,kAE final -,%' :' 1 urn5 .

Next we will prepare the equilibrium solutions from our stock KSCN and a diluted Fe(N0313 according to the following chart:

Sol ut ion 0.00200M Fe(N0a 13 0.00200M KSCN 0.05M HNO3 1 5.0 ml 1.0 ml 4.0 2 5.0 2 . 0 3.0 3 5 . 0 3.0 2 . 0 4 5.0 4.0 1.0 5 5.0 5 . 0 0 . 0

Measure the absorbance of each of these solutions as we did with the standards and record the values of the concentrations from the working c3libretion curve. From the discussion section, use the formulas to determine the concentrations of the Fe+3 i o n and the SCN- ion in each solution. Then substitute these values in the equilibrium equation and caliilate the Keq for each solution and average the results.

I

132

Requirements for remrt

a l l graphs and tables as they apply with the proper labels and t i t l e s Follow the guidelines established for report writing and include

Questions

1. If the blank used to set the spectrophotometer parameters was distilled water would this effect y o u r answer? 2. Why is nitric acid used in the ferric ion solution?

[Reference:Laboratory Manual for General Chemistry by Beran and Frady p u b l i s h e d by John Wiley and Sons.]

C.

133

DETERMINATION OF TH6 Ksp OF A SLIGHTLY SOLUBLE SALT

We will use the knowledge of equilibrium and solutions to determine the approximate solubility product of a slightly soluble salt.

a precipitate of two unknown ions versus the concentration of these ions after succesive dilutions. The equation for the equilibrium will be given, but the material being tested will not be known. We, therefore, must rely on our knowledge of the equilibrium process for slightly soluble species in solution. We know, for instance, that when the product of the concentrations of the ions in solution becomes greater than a given value, the excess material will precipitate out from the solution as a solid. If the concentrations of the ions is such that their product is less than the given value, we will have no precipitate.

In this experiment we will be comparing the formation of

Eor example: Ba+ 2 + SO4-2 = BaS04( s ) Ksp = [Ba+z][SO4-2] = 1 X 10-10

C. -

It tne concentration of the barium ion is one millionth of a mole per liter and the sulfate ion is the same their product is 1 X 10-12 which is below the ion product for the Ksp and there will be no precipitate formed. If the concentrations of the ions is lO-4M then their product is 10-8 and this is greater than the Ksp and the excess will form a solid.

the order of magnitude of the equilibrium constant for the unknown species. By using careful dilution and then combining the solutions and viewing any evidence of a precipitate, we should be able to estimate raLrry accurately to the power of ten for the equilibrium constant. A it.:Jle dettl:;d st i idy woKld allow for the deterininaticn of the mcltislier associated with the power of Ksp value, but for our purposes this zpgroach will suffice.

In this experiment we will use a method of serial dilutions to estimate

Procedure

This experiment will attempt to calculate an approximate value for the Ksp of an unknown salt generically called MA. You will, therefore, be given 'a solution that contains 0.200M M(NO3)z and 0.20OM K z A . Also, another trial of a different ratio of anion to cation will be determined and it is called XY2 where starting solutions are 0.200M X(NO3 ) 2 plus 0.200M KYz.

the procurement of 10 to 20 ml of M(N03)z and 10 to 20 ml of KzA.

and 10 drops of the A solution, stir vigorously and wait. If a precipitate forms within 20 minutes proceed to the next step.

a. The procedure for the determination of the Ksp for MA begins with

In a clean, dry test tube careful-ly place 10 drops of the M solution

134

b. Place 30 drops of the solution M and 30 drops of water into a different clean dry test tube. In another clean dry test tube place 30 drops of A solution and 30 drops of water and calculate the concentrations of these solutions. Now place 10 drops of the new M solution in a clean, dry test tube and add 10 drops of the new A solution. Mix thoroughly - does a precipitate form? Continue the experiment by placing 30 drops of the new M solution in a new clean, dry test tube and add 30 drops of water and repeat this with the A solution. Again calculate the concentrations of M and A. Again add 10 drops of the newest M solution to 10 drops of the newest A solution. If a precipitate forms continue the precedure until no more precipitate is noted. Calculate the Ksp by multiplying the concentration calculated for the M solution times the concentration for the A solution.

c. The procedure for the XYZ unknown is the same, but the X ( N 0 3 ) 2 and the KY solutions are used. The calculations proceed by multiplying the concentration of the X solution times the square of the ccncentration of the Y solution.

Requirements for report Follow the established guidelines for report writing. -- "

w u~ 5 t i on s

1. From your data and the knowledge that the generic formula is MA can you determine a set of possible compounds that might fit your data? 2. What effect would a high concentration of one ion of the pairs have on the precipitate formed?

[Reference: Experiments in General Chemistry by Whitten and Gailey published by Saunders College Publishing.]

135

tJETEROGENEOUS EQUILIBRIUM DETERMINATION OF TJB Ksp OF A SLIGHTLY SOLUBLE SILVER SALT

The principle of equilibrium is used to determine the Ksp a slightly soluble silver salt.

The basis for gravimetric quantities analysis rests in the equilibrium between solid ionic materials and solutions of their ions.

In general, this is a simple concept where the solution is saturated with the counter ion to force precipitation of the desired ion and then the product is washed, filtered and dried for weighing. Our experiment could follow this method and we probably would obtain satisfactory results, but by careful selection of our ions we can use alternate methods.

experimently determine determine the Ksp by mixing varying concentrations of each solution, allowing equilibrium to be established, then titrating the excess amount of one of the reactants. Since the starting concentrations are known it easy to relate the final concentrations of the ions and using the equation for the salt qa.lculate the Ksp. i.e.

In this experiment we will use silver ion and oxalate ion and

A g z C 2 0 4 = 2 A g + + c 2 0 4 - 2

KSP = [ A g + ] 2 [ C z 0 4 - 2 ] = # We will determine the oxalate ion concentration at equilibrium by

titrating it against potassium permanganate in acid solution. The equa t i on ;

2KMn04 + 5 K z C z O 4 + 8 H z S O 4 = 2MnS04 + l O C 0 2 8 H z O + 6K2SO4

titration and, also, the concentration of the silver ion, since it will be a stoichimetric relationship with the oxalate ion in solution.

The presence of the oxalate ion is determined by the permanganate

Procedure

Prepare 200 ml of 0.250 M A g N 0 3 and 200 ml of 0.250M NazCzO4 solutions by accurately weighing reagent grade chemicals and dissolving them volumetrically using a 200 ml volumetric flask.

Next obtain three 125 ml Erlenmeyer flasks that have been cleaned and dried. Number them in order and to flask number one, add by pipet 50 ml of 0.250M A g N 0 3 and 25 ml of 0.250M NaaC204 to the flask and stopper it. To flask number 2 add 7 5 ml of the silver nitrate solution and 25 ml of the oxalate solution, stopper and set aside. To the last flask add 50 ml of the silver nitrate and 50 ml of the sodium oxalate and stopper and set aside.

them to prevent supersaturation from occuring.

primary standard grade sodium oxalate.

Allow the flasks to sit for 30 minutes, but periodically shake

Next prepare and standaridize 250 ml of 0.1000 M mnO4 with

136

When the flasks have set for thirty minutes, then filter the flask contents through a dry retentive fliter paper and a dry funnel into a dry beaker. Then transfer by pipet 50 ml of the filtrate(fi1tered liquid) into a clean dry 250 ml Erlenmeyer flask. To the flask add 3 ml of concentrated H2SO4 and heat the flask gently to 800C. Titrate carefully with the permanganate until a persistent pink color remains for thirty seconds. If the titration takes to long a time you may have to reheat the solution.

titration volumes. Repeat the procedure for each of the flasks and record all the

Now using the equation: 2Mn04- + 5 C Z O 4 - 2 + 16H+ = Mn+2 + lOCO2 + 8H2O

we can calculate the amount of oxalate not precipitated as the salt. From this calculation and the fact that silver ion must react in two to one ratio calcualte the amount silver ion left in solution. N o w using the equation for Ksp and the concentrations determined calculate the Ksp for each trial.


your data for the standardization of the permanganate solution. -- Follow the established guidelines for report writing and include

Questions

1. Would the reaction be effected with a cold solution? Write the equation. 2. Could you suggest a different method to analyze this experiment to get the needed results?

[Reference: Experiments in General Chemistry by Drago and Brown published by Allyn and Bacon, Inc. pg 153-581

137

HYDROLYSIS

, We will use the knowledge of equilibrium and an understanding of pH to determine the interactions of ionic substances and water.

result is a definite acidic (basic) solution. The total ionization of this acid (or base) increases the hydrogen ion concentration (or decreaes) well above the normal or neutral level of water. This phenomenon is expected when dealing with strongly ionized acids and bases.

is no apparent effect on the concentration of the hydrogen ion concentration. On the other hand when a weak acid (or base) is dissolved in water tI lere is an increase in the concentration of the hydrogen ion or(0H-) but rlot to the extent of the strong acid or base. When a salt of the weak acid (or base) is placed in water the effect is to balance the equilibrium equation, i.e.;

When a strong acid (or base) is placed in water the

When the salts of strong acids and bases are dissolved in water there

HA + HzO = H30+ + A- A- + HzO = HA + OH-

C. ~

where the anion reacts with water to generate the undissociated acid (or base).

ionic: salts on the hydrogen or hydroxide ion concentrations with respect to their interaction with water.

The experiment will acquaint the student with the effect of variuos

Procedure

In this experiment the use of the pH meter will be

ine instructor will provide detailed precedures in tila use arid ~ ‘ a ~ e ‘11

Frepare 0.1 M concentrations of the following salts:

rlecessary to evaluate the hydrogen ion concentration of the solutions.

t L I - F‘! Toter and the associated electrodes. -

NaCl NaCz H3 Oz NaHCO3 NaHSOs NaHz PO4 NH4 C1 ( NH4 ) 2 SO4 NH4 Cz H3 0 2

When sufficient solution is prepared for each of the students to obtain about 100 ml of each one then proceed to test each solution according to the directions outlined by your instructor. Record the pH reading of each Yoltltion and what ions it contained. For each determination write an equation that explains the observation that was made.

1 138


the equation for each of the hydrolysis reactions. Follow the established guidelines for report writing and include

Quest ions 1. What would you expect when a salt such as NazHPO4 is placed in water? Write the equation. 2. How is t h e concept of hydrolysis used in chemistry?

[Reference: published by ;oh Wiley & Sons]

Laboratory Manual for General Chemistry by Beran and Brady

139

DISSOCIATION CONSTANTS

We will use the knowledge of equilibrium and pH to determine the dissociation constants of a weakly dissociating acidic or basic campound.

Usually an equilibrium reaction takes a little time to establish the equilibrium when the participants are joined together. In the case of the weak acid or base with water the equilibrium is in place almost immediately. Therefore their study is used in this experiment.

generic sense by the following; HA = H+ + A-, where the corresponding equilibrium equation becomes;

and the equilibrium constant Ka (called dissociation constant) is the constant for that acid. If we used acetic acid as our representative acid the value would be about 1.8 X 10-5 when the concentrations are in moles per liter and the temperature is 25oC. -A The expression can be written for a weak base much the same way and the CoI lc ius ions amount to much the same except we are working with a base instead of an acid.

If we are aware of the K value for our particular reaction we can determine the concentration of the various species with relative ease, but if the acid is unknown it becomes a little more difficult to determine the concentrations af the various species unless we take advantage of the fact that the equilibrium equation is totally reversible and we can then evaluate the information obtained in our titration data to determine the K value.

our uniLrLown material. The first method is based on the fact that when . - 1?1szc thc acid or base in solution the concentratian of the hydrzgen j o n or the hydroxide ion and the corresponding counter ion will have th? same value, If we know the total concentration of the acid or base then at any given pQint in the equilibrium [HA] = F - [H+], where F is the formal concentration of the acid, and [H+] = [A-I

Using the above information we can substitute into the equilibrium equation as follows:

The equation for the reversible reaction can be represented in a

Ka = [H+][A-l/[HAl

There are two methods that allow for the calculation of the K value fclr . .-

Ka = [H+]2/ F - [H+] From this arrangement we can calculate Ka. The error of this technique

is potentially high due to the fact that the value used to determine [H+] is then squared and divided by a difference using the value of [H+]. The second method reduces the potentially large error because the measurements are not magnified so extremely. In this technique the reaction of the acid with the base produces the following;

and we see there is a quantity of anion produced in relation to the undissociated acid. When one fourth of the original HA has been titrated then one fourth of the anion has been generated. Then using the equation;

Ka = CH+l{CA-I/[HAl) the values for Ka can be calculated at various ratios of anion to undissociated acid and the values can be averaged.

HA + OH- = A- + HzO

I 14 0

Requirements f o r report

calculation set up for each result. Follow the established guidelines for report writing and show a

Questions

1. Calculate the result using both methods decribed in the discussion. Is there any difference? 2. What would you do to improve the results?

[Reference:Laboratory Studies and Problems in General Chemistry by Stone and McCullough published by McGraw-Hill, Inc.]

14 1

DETERMINATION OF pKa OF AN ORGANIC ACID

We will utilize the knowledge of the pH plot and the equation of the equilibrium constant to determine the pKa of an unknown weak organic acid.

Most organic acids are less soluble in water solution as a result of the intramolecular hydrogen bonding of the carboxyl group on the molecule. This leads to a lowering of the ionization of the molecule and thus a lower solubility. The equilibrium constant (Ka) is a function of the strength of the acid. If the organic acid has an attached group that reduces the ability of the hydrogen ion to separate from the molecule the acid is said to be weak. If the group attached to the carboxyl group is an electro-negative group such as a nitrogen containing group then we may see a strengthening of the acid and it may become closer to the ionizing capability of a mineral acid such as phosphoric acid.

con-ihination we can more easily dissolve an organic acid such as benzoic acid and then study the dissociation constant of these acid types.

carefully plot the pH versus volume relationship it is possible tc, estimate the pKa of the unknown acid. At the point on the graph where one half of the acid has been neutralized the ratio of undissociated acid to anion concentration is 1:l. At this point the pH then is equal to the pKa. (See the Henderson-Hasselbach equation).

By using a slightly less polar solvent such as alcohol and water

If we titrate the acid with a known concentration of a weak base and

Procedure

Weigh out an approximate 0 . 0 2 g quantity of the assigned acid and dissolve them with 25 ml of denatured (reagent) alcohol in a 100 to 150 ml beaker and then add 25 ml of distilled water. Prepare to titrate with a 0.10 N NaOH solution using a pH meter and a stirring apparatus. Be sure to insulate the beaker from the heat of the ztirring motor or use a A T C ( A u t o T e m p C o m p e m s a t o r ) attachment with the meter.

pH readings. Reduce the increment size when you approach the endpoint. continue the titration well beyond the endpoint to get a good plot. Plot your pH versus volume data and determine graphically the equivalence point and the pKa at the point where the volume is equal to 1 / 2 oh the neutralization volume.

Add your titrant slowly inb 1.0 ml increments recording both volume and

Requirements for rep ort

your graph of the titration properly titled and labeled. Follow the establishied guidelines for report writing and include

Quest i ons 1. Would it be more advantageous to use higher weights and concentrations to do this lab? Why? 2 . Would it be necessary to Know the molarity of the base?

[Reference: Experimental Methods in Organic Chemistry by Moore, Palrymple, and Rodig published by Saunders College Publishing pg 1 8 9 - 911

142

BUFFERS

The use of the knowledga of hydrolysis and acid-base

When a weak acid or base is placed in water there is dissociation of the acid or base to a small degree and the properties used to define an acid or base are present in the solution.

A l s o , when the soluble salt of one of those weak acids or bases is placed in water the dissociation is basically complete and the conjugate species interacts with water to exhibit the opposite properties of the weak acid or base and as such could be analyzed as the opposite species.

exists as if the solution were in equilibrium at a set state of conditions and neither additional acid or base added can greatly effect this stable state. This situaltion is described as a buffer solution.

There ios e n o u g h of the undissociated acid molecules present to neutralize a finite quantity of base when added. Also, there is enough of the dissociated conjugate base(anion) present to react with any acid that maybe added. The overall equilibrium remains status quo.

buffers that can utilized.

of possible buffer solutions and the determination of the ratio of the buffering species in solution.

equilibria can be used to study the properties of buffer solutions.

When a weak acid and its salt are present in the solution a situation

Depending on the particular species there are a large variety of pH

Henderson-Hasselbach equation allowsfor the calcualtion of a variety

pH = pKa + log [A-] / [HA] for a weak acid

pH = pKb + log [B-]/[BOH] for a weak base

Bilffers are generallly obtained commercially but in some instances it is advantageous to know how they might be prepared. Our experiment is to generate a few buffers.

Procedure

( 1 ) The first step is to prepare a buffer that is easily done arid test the solutions.

Prepare a 100 ml solution of 0.10 M N a C z H 3 0 2 and 100 ml of 0.10 M HCz H3 0 2 .

Place 20 ml of each in a 100 ml beaker and test the pH of the solution with a pH meter. It should read 4 . 7 - 8 according to the Henderson- Hasselbach equation. If it does not check the calibration of the pH meter.

Next add to the solution 5 ml of 0.05 M HC1 and recheck the pH meter. Does it remain the same? N o w add 0.5 M HC1 carefully recording the volume added until there is

a one unit change in the pH meter reading( i.e. 4 . 7 -3.7.) How many ml’s does it take?

and repeat the above procedure but using 0.05M NaOH and 0 . 5 M NaOH. reiterate the questions and determine the answers.

Again place 20 ml of each of the starting solutions in a 100 ml beaker

14 3

(2) Now using the Chemical Rubber Company Handbook prepare and test a buffer that will give you a pH of 7 and not change with the addition of acid or base over a range of 5 ml of 0.02M concentration.

Requirements f o r report

prepared buffer to the instructor for test. Follow the established guidelines f o r report writing and submit your

Questions 1. What is meant by the capacity of a buffer? 2. Choose another buffer system then the one you prepared and calculate the ratio of the salt to acid or base needed.

144

ELECTROCmMISTRY VOLTAIC CELL - "ST EQUATION

We will utilize the knowledge of a spontaneoous oxidtion- reduction reaction and the electromotive series to study the voltaic electrochemical cell.

reducing agent there is a reaction and the electrons are transferred from the reducing agent to the oxidizing agent. If the reaction species are separated by a semipermeable membrane then the electrons must pass from the reducing agent through an outside conductor to the oxidizing agent for the reaction to proceed. This is accomplished by connecting a wire(conductor) from the electrode in contact with the reducing agent to the electrode in contact with the oxidizer. The electrolytes in the solution will act as elctron carriers within the cell to complete the ccircuit and allow the reaction to take place.

as the anode for a spontaneous reaction and is the negative terminal of the voltaic cell. The oxidizing agent,(may be ions in solution with an inert electrode), is the cathode and is positive. The two electrodes 'are is contact with a solution of their ions, or other electrolytes, but are separated form one another by means of a porous cup, salt bridge, or gravity dddue to density of the two solutions.

be a current generated and a voltage. The voltage may be determined by cornparin-g the location of the electrode pair onb the Electromotive Force Table, or a table of Standard Reduction Potentials. For example, zinc is located below hydrogen on the Standard Reduction Potential(SRP) table and has a value of -0.763 Volts as an oxidizing agent, or its ability to be reduced. While copper has a SRP value of + 0 . 3 3 7 Volts. If we place Cu(I1) i o n in the presence of zinc metal the copper will be displaced in solution by the zinc. The potential voltage of the cougie would be +0.337V - (-0.763V) = 1.1OOV. This reaction woald taka p1a.z.s a5 a spontaneous reaction as shown in the equation GO = -nFEo, since EO is positive, GO would be negative and by definition this denctes a spontaneous reaction.

When an oxidizing agent is placed in the presence of a

The reducing agent, generally in the form of a metal strip, is defined

When an electrode pair is set up with an external conductor there will

The reaction equation is as follows:

where the zinc is oxidized and the copper is reduced. Now with G = - RTln Kc, then EO = RT/nF(lnkc) and if l o g l o is

used then EO = 0.059l/n(log kc), or the potential of the cell is a function of the concentration of the ionic species in solution and the E c e l l is equal to EO - 0.0591/n (log Q), where Q equates to the mass-action quotient. If our cell has a change in concentration from the values expressed in the SRP tables this equation allows for calculation of the change in potential.

I 14 5

Procedure

In this experiment we will compare a series metal-metal ion redox couples and determine the cell potentials and compare them to the literature values from various sources.

The metal-metal ion couples we will consider are; Zn-ZnSO4(1M), Mg- MgSO4(1M), Cu-CuS04(1M), Ni-NiS04(1M) and Sn-SnClz(1M).

The cells will be arranged by placing in a 250 ml beaker, the first of the metal-metal ion and then in a porous cup place the second of the metal-metal ion pairs. Before starting any of the reactions be sure the surface of the electrodes(meta1 strip) is polished clean.

resistance and to the second terminal of the resistance connect another conductor which will be connected to the second electrode. Across the two terminals of the resistance connect a voltmeter (watch your terminals). Now connect the second electrode with the alligator clip and quickly read the voltage and record this value.

distilled water and place the next metal-metal ion pair to be tested in C.”. - -.r? connect as before but check to see if the terminals of the voltmeter are in the right places. Each cell is different.

all the other pairs and you have values for all.

in the experiment. Nernst equation: Now take the ZnSO4 solution and reduce it to a U . O l M concentration by dilution of 1 ml to 100 ml. Set up the Zn-ZnS04(0.01M) cell along with the Cu-CuSO4(1M) half cell and determine the voltage.

Use the equation from the discussion to calculate the anticipated v a i i i e - are they the same? If not, why not?

To the first metal strip(e1ectrode) attach an alligator clip and a high

Remove the porous cup and disconnect the leads. Wash the cup with

A 1 - - .- Repeat the procedure until each of the ion pairs has been tested with

Now arrange the metal-metal ion pairs in order of potentials determined


drawing of your cell arrangement. Fallcx the established guidelines for report writing and include a

Quest ions 1. If the zinc electrode was used as the reference value on the SRI? table, what would your voltage determinations be? 2. Why is it necessary to isolate the two solutions?

[Reference: Fundementals of Chemistry by Brady and Holum, John Wiley publisher]

14 6

ELECTROLYTIC CELL DETERMINATION OF AVOGADRO’S NUMBER

We will use the knowledge of an electolytic cell and the Laws of Faraday to verify Avogadro’s Number.

The electroyltic cell is nonspontaneous and the redox reactions performed in this cell must be forced by an external source (i.e. a battery or power supply.) This type of a cell has is uses such as electroplating or refining of metals like aluminum, ets. They, also, are useful in studying the various phenomnon of oxidation-reduction reactioris or analyzing theory that has been studied in various scientific presentations.

Number. Avogadro suggested the value 6.022045 X 1 0 2 3 particles in his study of the properties of gases although the actual number was determined a w h i l e later.

value called Avogadro’s Number by depositing a measureable quantity of -- an element such as silver or copper onto a preweighed electrode and using the electrochemical law (Faraday’s)that the amount of mass of an element deposited by an electrical source is directly proportional to the amount of electricity used. The electrical quantity Q in coulombs is equal to the amperes (I) times time (t) in seconds.

as measured by a scientist named Millikan. Therefore;

and by knowing the number of electrons needed to deposit the mass of the element we worked with, we can determine the number that would equate to the number of atoms in one gram-atomic weight(mo1e) of that e l2=a:en t . This should equal Avogadro’s Number.

We will apply the properties of this ccell to a study of Avogadro’s

By using the electrolytic cell we will experimentally determine the

The electrical charge carried by an electron is 1.6 X 10-19 coulomb

Q X 1 electron/l.6 x 10-19 coulomb = # electrons

. Procedure __ - __ -

Prepare 200 ml of 0.5 M CuSO4 in a 250 ml beaker. To the heaker suspend 2 polished and preweighed copper strips to be used as electrodes. To the copper strips attach two leads with alligator clips.

the leads and complete the circuit. Adjust the amperage to about 50 milliamps by varying the voltage or the resistance and allow the current to flow about one hour. Monitor the flow rate of electrons periodically(every 10 minutes) and record the value.

anode(+) and the cathode(-) from the cell. Rinse them carefully withe distilled water and then with a little acetone to dry. Weigh each of them separately and record the weight gain and loss for each.

Determine the average mass transferred and use this value to calculate the value for Avogadro’s Number.

Compute the value for Q (multiply I X t(sec)) then diviae Q by the charge per electron (1.6 X 10-19) to get the number of electrons.

Now arrange a proportion to determine the number of electrons transferred per gram-atomic weight. Then determine the number of electrons transferred on the basis of an equivalent weight.

Using a D.C. power supply, a milliammeter, a 90 ohm rheostat connect

At the end of the time disconnect the leads and carefully remove the

I 14 7

Requirements f o r report Fol low the established guidelines for report writing.

Questions 1’. What is the precision of this experiment? Why? 2. State Avogadro’s theory. 3. How does your experimental value compare to Avogadro’s? 4 . What is the % error? 5. What caused the difference?

[Reference: Laboratory Experiments for Foundations of Chemistry by Toon and Ellis, published by Holt, Rinehart and Winston, pg 1611

1 4 8

ELECTROLYTIC CELL COULOMETRIC PRODUCTION OF CUPRIC ION

We will use the knowledge of Faraday’s laws to study the electrolytic cell and to compare techniques in measurement of various quantitatively produced elements.

Faraday stated two truths as follows: ( a ) the weight of an element generated at an electrode is directly

(b) that the weights of different elements that are produced by proportional to the quantity of electrons transferred through the cell,

equivalent amounts of electron are proportional to their respective equivalent weights.

to quantitative measurements as do mass quantities or volume or gases.

equals coulombs, which when divided by the value 96,500 equals equivalents we can study the various parts of an electrolytic cell(nonspontaneous).

cathode(-) in an electrolyte solution such as dilute H 2 S O 4 . With the ~ t e of an inverted buret or a eudiometer tube we can collect the hydrogen gas evolved at the cathode. During the reaction there will be copper (11) ions evolved at the anode and a measure of the time and amperage will give us a record of the coulombs of electricity utilized in the cell to compare with the amount of each of the other two species produced.

In studying the above statements it becomes clear that electrons relate

With understanding and the fact tthat amperes times time in seconds

In this experiment we will arrange a copper anode(+) and a Chromel C.

The reactions for each of the electrodes is as follows:

Cu(e) = Cu+2(aq) + 2e-

Using the simple loss of weight at the anode as a measure of the 2H+(aq) + 2e- = H2c tz)

copper(I1) put into solution and the time in seconds times the average amperage we have two measurements of the equivalents utilized. When we finally compare these to the equivalents of hydrogen at the cathode there should be good agreement in the results.

in the sulfuric acid solution may be carried out as a fourth measurement technique.

In addition, if desired, an iodometric determination of the copper ion

Prepare a solution of 0.25M H2SO4 sufficient to fill a 50 ml buret and have about 75 ml in a 150 ml beaker. Next clean a piece of #12 copper wire, of which the last 6 inches of a 12 inch piece have been twisted around a pencil and bent at right angles to the straight stem, using 50 ml of warmed 0.5 M nitric acid solution. Rinse the wire with distilled water and then a little acetone to dry it. Weigh the wire to thc nearest 0.1 mg and attach it as the anode. Next a piece of 12 inch Chromel is prepared as the copper wire but the bend is in a U-shape and the wire not inside the inverted buret must be sealed form the solution using

..., either glass tubing and an epoxy sealant or shrink A* -44. insulation used for electrical junctions. cr-

-

149

With the electrodes prepared place them in the beaker with enough of the H2S04 solution to fill the buret. Invert the buret and insert the Chromel electrode into the open end of the buret and with a suction b u l b draw the solution into the buret above the last calibration mark.

none lower the meniscus to the calibration mark(50 ml or 25 ml for what ever size buret used.) Add the rest of the liquid and prepare thm circuit for the reaction.

sufficient resistance in the circuit can be used. A milliammeter is also necessary in the circuit. Connect the positive terminal of the power supply to the copper wire and the negative lead to the Chromel wire. Begin timing the reaction as soon as the leads are connected and a current is flowing. A current of 30 to 50 milliamperes is suggested to give enough data to evaluate the consistency of the current. Take readings of the amperage and the hydrogen gas production on a periodic basis and record these values. At the end of the experiment average the the ampere readings and the hydrogen gas production and this will verify the consistency. Continue the experiment until 25 to 40 ml of the hydrogen has been collected and then disconnect the leads.

n A r l A c\ iAttle acetone to dry it and weigh it.Do not place it i n a drying oven. Why? Record the volume of hydrogen and measure the liquid volune in ml in the buret that is above the liquid level in the beaker.

Seal the buret with the stopcock and watch for any leakage. If there is

A 6-12 volt D.C. source is needed or a higher voltage source with a

Remove the copper wire, rinse with a little distilled water and then C.

iiecord this value. Repeat the experiment once more and collect the data in the same

Now compute the number of equivalents of copper generated by dividing manner .

the weight loss on the electrode by the weght per equivalent according t o the equation inthe discussion section. Next compute the number cf equivalents of electrons by multiplying the average amperage by the . . - : xk - r 2f sFconds the reaction was run and dividing by the factcr Q P ' 1 1 , I -niJl@mhs/equivalent. Finally, compute the number of equvalents of hydrogen generated using the ideal gas law and the laboratory temperature and pressure and the change in the pressure due to the suspended column of liquid remaining in the buret. The volume in ml measured times the density of the liquid (about 1.014g/ml) and then divided by the density of mercury (13.6g/ml) will give the amount of pressure drop in the buret. This is then subtracted from the pressure reading plus the vapor pressure due to water at the lab temperature to obtain the proper values to substitute into the gas law equation.

Compare the values obtained for each determination and compute the percent relative error in the methods.

&wirenents for report

tables for your data. Follow the established guidelines for report writing including

Quest i ong 1. Which one deviates the most? Can you'explain why? 2 . Nhat would improve the method?

[Reference: Laboratory Studies and Problem in General Chemistry by StorLe and McCullough published by McGraw-Hill]

150

ISOLATION OF CAFFEINE FROM A SOFT DRINK

The purpose of this experiment is to extract or isolate caffeine from Pop. Although pop is mainly a solution of flavored sugar in water there are other interferring materials. Some soft drinks contain ingrediants such as caffeine to enhance the flavor of the drink. Some of the impurities are the flavoring agents which may include some complicated organic structures and preservatives such as potassium benzoate.

The other water soluble products are eliminated when the caffeine is extracted with the organic solvent such as dichloromethane.

Procedure

In a 250ml Erlenmeyer flask place 100.0ml of pop then attach a hose to an aspirator and evacuate until the solution ceases to evolve carbon dioxide bubbles. At this time add 2.ug of -- calcium carbonate and warm gently. Filter using a qualitative filter paper, filter the solutionby gravity into a 125ml Erlenmeyer flask. Rinse the residue left in the flask with two 10 ml washings of hcjt water and add to the filtered material. Discard the residue after extracting the remaining liquid by carefully pressing the residue and filter paper.

Cool to room temperature and extract with 4.0ml of dichloromethane. Be careful to avoid an emulsion by to vigorous agitation of the

Allow the layers to separate and withdraw the bottom layer solution.

with a Pasteur pipet and transfer to a small filtering flask by means of a funnel with a small layer of anhydrous sodium sulfate.(about 2.0g) Repeat the extraction procedure with four additional 2.0ml portions of dichloromethane and add these to the filtering flask.

The organic solvent is now evaporated to dryness using a steam bath or sand bath in the HOOD.

The product is now ready for purification and characterization which requires sublimation of the crude caffeine.

Using an aspirator, a filtering flask and a cold finger the sublimation canbe accomplished. First clamp the flask on a ring stand and connect to the aspirator. Next seal the cold finger in the flask about 1/2cm above the crude caffeine. Then with a micro- burner begin the heating process. Be sure not cause the crude caffeine to begin melting for it will then decompose. Continue this process until the sublimation is complete, no more caffeine forms on the cold finger. Remove the heat and allow the system to cool to room temperature under reduced pressure.

the cold water connections. Then remove the cold finger from the flask and scrape the crystals of caffeine onto a weighing paper or other preweighed collecting container. A melting point and an IR spectrum should be included with this experiment.

When the cooling is completed carefully remove the vacuum and

151

Reqirements for r e p o r t

include the satructural formula of caffeine and its properties. Follow the established guidelines for report writing and

Questions 1. What.is the term alkaloid defined as? 2 . What is the LDora1 for a rat?

[Reference: Microscale Organic Laboratory by Mayo, Pike and Butcher published by John Wiley & Sons]

PREPARATION OF ASPIRIN

152

Aspirin has been used for years as a fever reducer, anti- inflamation drug, and pain reliever. Its effect on the body is still not fully understood, but its chemistry is not to difficult.

The molecule is actually an ester of salicylic acid and the acetate ion. Most esters are the condensation of an organic acid and an alcohol, but aspirin is actually the condensation of anorganic acid and an organic acid.

an acid molecule with a hydroxy group also attached. This is actually not the case, since salicylic acid is composed of

CsH5(0H)(COOH) + H3C202H = CsHs(H3C20z)(COOH) + H2O

The reaction equation decribed above is an abreviated form of the reaction. The complete reaction must also include a catalyst to enhance the reaction kinetics. If hydrogen ion is used it will help to facilitate t h e speed of the reaction and push it towards completion.

Erocedure

In a 18 X 150 mm test tube place 1 g of pure salcylic acid and a b o u t two ml of acetic anhydride. To this add two drops of concentrated sulfuric acid and heat the tube gently in a hot water bath to about 85- 900C. Heat for ahout five minutes and then remove the tube from the heating bath and add two ml of distilled water to decompose the anhydride.

water and chill thoroughly in an ice bath.

t h e product until it is dry.

extremely flammable and must be handled carefully) and theii add dri c-qx--l c\zQur,t of petroleum ether(bp 30-600) and then cool in an ice Lac:-. Again collect by vacuum filtration and determine the melting point of the dry crystal.

The material is then crystallized by adding an additional 10 ml of

Collect the crude product by vacuum filtration and draw air thru

Tc purify the crystals, dissolve them in 10 ml of ether (ether is

Requirement for report

your melting point data and yield amounts. Follow the established guidelines for report writing and include

Qu e s t i o n s 1. How is the commercial aspirin prepared? 2 . Are there any other salicylic acid derivatives used in medicine?

[Reference: Laboratory experiments in Organic Chemistry by Adams and Johnson published by Macmillan Company]

inorganic chem

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