TABLE OF CONTENTS

Experiment Page 1 Measurement 1

2 Density of Solids and Liquids 9

3 Verifying the Empirical Formula of Magnesium Oxide 11

4 The Hydrogen Emission Spectrum 13

5 Flame Spectra of Metal Cations 17

6 Covalent Bonding and Molecular Structure 19

7* Solutions and Reactions 25

8* The Use and Abuse Aluminum 27

9a Standardization of a Sodium Hydroxide Solution 29

9b Determining the Molarity of a Hydrochloric Acid Solution 31

10 Determining the Molar Volume of an ideal Gas and the Universal Gas Constant 33

11* Halogens and Their Compounds 37

12* Chemistry of Natural Waters 39

13 Specific Heat of Metals 41

14 Additivity of Heat of Reactions: Hess’s Law 43

Appx 1 Periodic Table

Appx 2 Chemistry Laboratory Conduct * Introduction and Procedure for these labs can be found in the Catalyst.

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Measurement Introduction Much of the work done in chemistry (and all science for that matter) involves measurements. Most scientific measurements involve the International System (SI system) that employs metric units. Metric units are convenient because they relate to one another by powers of ten. For example 1000 meters (m) are equivalent to 1 kilometer (km). Many of the measurements made in the United States not involved in science use non-metric measurements such as feet, inches, and pounds. In this laboratory exercise you will become familiar with the metric system, practice converting between metric and non-metric units of measurements, and be tortured with significant figures. Procedures Part I. Length Measurements 1. Measure the length and height in centimeters of the box shown below. Convert these

measurements to inches and to meters.

Part II. Temperature Measurements 1. Determine the temperature of the room using a Celsius thermometer 2. Convert this recorded temperature to Kelvin and Fahrenheit. 3. Prepare an ice water bath by filling either a 150 mL or 250 mL beaker with water and ice until

the volume is near the top of the beaker. 4. Immerse your thermometer in the ice bath and stir (don’t break the thermometer) and record

the temperature. 5. Convert this recorded temperature to Kelvin and Fahrenheit. Part III. Mass Measurements 1. Fill a 250 mL Erlenmeyer flask to the 100 mL mark with deionized water. 2. Place the flask on the analytical balance. The mass will be changing due to evaporation.

Record the mass. 3. Record the mass at one minute intervals for five minutes. 4. Fill a 250 mL beaker to the 100 mL mark with deionized water. 5. Place the beaker on the analytical balance and record the original mass. 6. Record the mass at one minute intervals for five minutes. 7. Produce two x-y scatter graphs. Plot time (recorded in minutes, starting at 0) on the x axis

versus mass (recorded in grams) on the y axis for both the Erlenmeyer flask and beaker data. Your graph should be labeled properly, contain the line of best fit, and the equation for the line of best fit.

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Part IV. Volume and Mass Measurements 1. Determine the mass of a dry 50 mL Erlenmeyer flask. 2. To the best of your ability place 30 mL of deionized water into the flask and determine the

mass. 3. Determine the mass of a dry 50 mL graduated cylinder. 4. To the best of your ability place 30 mL of deionized water into the cylinder and determine the

mass.

Part V. Precision and Accuracy Record the values for the mass of water placed into the Erlenmeyer flask as well as the graduated cylinder obtained by other class members. Data and Workup Part I (Data) Length of box in centimeters _______________ Height of box in centimeters _______________ Part I (Calculations, show all work and box your answers!) Calculate the length of the box in inches: Calculate the height of the box in inches: Calculate the length of the box in meters: Calculate the height of the box in meters: Calculate the area of box in square centimeters:

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Calculate the area of box in square inches Part II (Data) Temperature of room in degrees Celsius _______________ Temperature of ice water bath in degrees Celsius _______________ Part II (Calculations, show all work and box your answers!) Calculate the temperature of the room in Kelvin: Calculate the temperature of the room in Fahrenheit: Calculate the temperature of the ice bath in Kelvin: Calculate the temperature of the ice bath in Fahrenheit:

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Part III (Data)

Erlenmeyer Flask

Beaker

Time (minutes) Mass (g) Time (minutes) Mass (g)

0 0

1 1

2 2

3 3

4 4

5 5

Part III. (Calculations, show all work and box your answers!) Based on the slope of the graph produced, what is the rate of evaporation for water in the Erlenmeyer flask (g/min)? Based on the slope of the graph produced, what is the rate of evaporation for water in the beaker (g/min)? Which piece of equipment, an Erlenmeyer flask or beaker should be used if you were concerned about a liquid sample evaporating as you work during the lab?

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Part IV. (Data) Mass of empty Erlenmeyer flask _______________ Mass of Erlenmeyer flask containing 30 mL of water _______________ Mass of empty graduated cylinder _______________ Mass of graduated cylinder containing 30 mL of water _______________ Part IV. (Calculations, show all work and box your answers!) Calculate the mass of the water placed into the Erlenmeyer flask: Assuming the density of water is 0.997 g/mL, the expected mass placed into the flask would be 29.9 grams. Calculate the percent error between your experimentally measured value and the expected mass: Calculate the mass of the water placed into the graduated cylinder: Assuming the density of water is 0.997 g/mL, the expected mass placed into the graduated cylinder would be 29.9 grams. Calculate the percent error between your experimentally measured value and the expected mass: Based on your results, which piece of glassware did you find to be the best to measure volume accurately? Why?

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Part V (Data) Experimental values for the mass of water placed into Erlenmeyer flask for each student group.

Experimental values for the mass of water placed into graduated cylinder for each student group.

Part V (Calculations, show all work and box your answers!) Calculate the class average for the mass of water placed into the Erlenmeyer flask: Calculate the class average for the mass of water placed into the graduated cylinder:

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Calculate the percent difference for the data of the mass of water placed into the Erlenmeyer flask: Calculate the percent difference for the data of the mass of water placed into the graduated cylinder: Calculate the percent error (using the class average) for the data of the mass of water placed into the Erlenmeyer flask: Calculate the percent error (using the class average) for the data of the mass of water placed into the graduated cylinder: Based on the class results, which piece of glassware did you find to be the best to measure volume accurately? Why? Based on the class results, which piece of glassware did you find to be the best to measure volume precisely? Why?

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9

Density of Solids and Liquids Introduction In this experiment, physical properties of several different chemical samples will be studied. Among many physical properties you will observe, the main emphasis of this experiment will be upon a quantity called density. Density (d) of matter is defined as mass of matter (m) divided by its volume (V). Using a mathematical equation, we can define the density (d) as:

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Equation 1. d =mV

In other words, it is necessary to determine the mass and the volume of an object in order to calculate its density. Table 1 provides the accepted values for the density of common substances at 25 °C.

Table 1. Density values calculated at 25°C Substance Density (g/mL) Water, H2O(l) 0.998 NaCl(s) 2.165 Magnesium, Mg (s) 1.74 Lime stone, CaCO3(s) 2.710 Zinc, Zn(s) 7.14 Iron, Fe(s) 7.86 Nickel, Ni(s) 8.90 Copper, Cu(s) 8.92 Mercury, Hg(l) 13.5462

As you can see in the above table, each substance has its own characteristic density. Density can be used as one clue to determine the identity of an unknown substance. There are many applications associated with the density of a substance. For example, from the density and a known mass of a substance, the volume of the substance can be determined by rearranging equation 1:

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Equation 2. V =md

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Procedures Part I. Density of Solid Samples IA. Solid Cubes 1. From the reagent central, obtain a sample of two of the following solid cubes.

a. Aluminum (Al) b. Iron (Fe) c. Copper (Cu) d. Brass (Bu)

2. Determine the mass of each sample using the proper balance determined by your instructor. Record data in your notebook.

3. Measure the length, width, and height for each sample using the metric ruler from your lab kit. Record data in your lab notebook.

4. Calculate the density of each sample. These results should be placed into a results table. IB. Solid (Non Metal) Blocks 1. From Reagent central, obtain two grey block samples. Determine and record the density of each sample as described in the previous section. IC. Quartz Samples 1. From Reagent central, obtain two separate quartz samples. Results will be best if each

sample contains more than one stone. 2. Determine the mass of each sample using the proper balance determined by your instructor.

Record data in your notebook. 3. Determine the volume of each sample using water displacement in a graduated cylinder.

Record data in your notebook. 4. Determine and record the density of each sample as described is the previous section. 5. Return samples to Reagent Central dry. ID. Solid Cylinders 1. Your instructor will provide you with a solid cylinder. 2. Determine the density of the cylinder using a technique of your choice. 3. Record the all data and results in your notebook. Part II. Density of Liquid Sample 1. Obtain a liquid solution (aqueous sodium chloride) from the instructor. The liquid sample is

stored in a 125ml Erlenmeyer flask. To prevent loss of water through evaporation, the flask is sealed with a rubber stopper.

2. Record the unknown number. 3. Construct a data table in order to record both mass and volume for six measurements. 4. Measure the mass of an empty 100mL beaker (initial mass) with watch glass “lid” using an

analytical balance. Be sure to dry the beaker completely before the measurement. 5. Using a 10.00 mL volumetric pipet, transfer 10.00 mL of solution into the 100 mL beaker. 6. Measure and record the mass of the beaker (with lid) and the 10.00 mL of liquid sample. 7. Repeat step 6 until you have added 50.00 mL of solution to the beaker. 8. Using Excel, make a graph for your experimental data. This graph should have mass (in

grams) of solution as y-axis (vertical axis) and the volume of solution (in milliliters) as x-axis (horizontal axis). a. Give graph a title and label both axis properly (including units) b. Display the “line of best fit” along with the equation for this line on graph. What does the

slope represent for this experiment?

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Verifying the Empirical Formula of Magnesium Oxide

Introduction In this portion of the lab you will determine the empirical formula of a magnesium oxide. Looking at the positions of magnesium (column 2A) and oxygen (column 6A) on the periodic table, the empirical formula for magnesium oxide (MgO) can be predicted. In this experiment you will verify this formula and calculate the percent error for the experiment. The experiment starts by heating a known mass of magnesium ribbon in the presence of oxygen and nitrogen from the atmosphere. The magnesium reacts with the oxygen to form MgO according to reaction 1a below. The magnesium will also react with the nitrogen to produce magnesium nitride (Mg3N2) according to reaction 1b. All of the original magnesium must be incorporated into magnesium oxide for a successful experiment. In order to remove unwanted Mg3N2, water is added which will react with Mg3N2 and produce ammonia gas (NH3) and MgO according to reaction 2. Once this step is complete, all of the original magnesium exists in the form of MgO. Reaction 1a: 2 Mg(s) + O2(g) ® 2 MgO(s) Reaction 1b: 3 Mg(s) + N2(g) ® Mg3N2(s) Reaction 2: Mg3N2(s) + 3 H2O(l) ® 2 NH3(g) + 3 MgO(s) Procedure 1. From reagent central, obtain a crucible. CHECK CRUCIBLE FOR CRACKS. If the crucible

is cracked it will not withstand the heat in this experiment. 2. Using a Bunsen burner, clean crucible and the crucible cover while holding them individual

over the flame using metal tongs. BE EXTREMELY CAREFUL not touch the hot crucible or cover. Let the crucible and cover cool to room temperature (≈ 5 minutes).

3. Using metal tongs determine the mass of the crucible. Do not handle the crucible with your hands. This will affect the mass determination.

4. Using tweezers from your lab kit, obtain a piece of magnesium ribbon (≈5 cm in length) that has been prepared for you by the kind people in the stockroom.

5. Place the magnesium ribbon in the crucible. Weigh the crucible with the magnesium ribbon inside.

6. Place the crucible with lid properly positioned into a clay triangle. 7. Heat the crucible for approximately ten minutes to ensure a complete reaction. Make sure

the Bunsen burner is positioned properly to allow the hottest part of the flame to be at the bottom of the crucible.

8. After heating, allow crucible and contents to cool for approximately ten minutes. 9. Add approximately 10 drops of deionized water to the mixture inside crucible, repeat steps six

through eight. 10. Weigh crucible containing the magnesium oxide product. 11. Use the data collected in this experiment to verify the empirical formula 12. Calculate the percent error for you experimentally determined empirical formula.

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Extra Problems During this lab activity there will be “down time” while you wait for materials to either heat up or cool down. During this time complete the following two problems in your lab notwebook. Show all work. 1. An experiment was conducted in order to determine the empirical formula of a compound

made from copper and sulfur. The mass of a clean dry crucible was 10.443 grams. Copper wire was added to the crucible and the mass of the crucible and copper sample was 11.229 grams. The wire was then covered with excess sulfur. The crucible was heated until all of the excess sulfur burned off. After the crucible was allowed to cool to room temperature, its mass was determined. The crucible plus the product of the reaction between copper and sulfur was 11.432 grams. Determine the empirical formula of the product.

2. An experiment similar to that described above was conducted to determine the empirical

formula of an iron oxide sample. At the start of the experiment the mass of iron was 16.76 grams. At the end of the experiment the mass of the iron-oxygen compound was 23.96 grams. Determine the empirical formula of the iron oxide product.

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The Hydrogen Emission Spectrum Introduction Gas discharge lamps employ electricity to excite the particles in a tube filled with a particular substance. The added electronic energy is given off as light. By using a spectrophotometer, the emitted light can be separated into individual wavelengths. Each line corresponds to light with a specific amount of energy (E), which can be characterized by wavelength (l), and frequency (n). Using Planck’s constant (h, 6.626 x 10-34J×s) and the speed of light (c, 2.998 x 108 m/s), the energy for each line within the spectrum can be calculated according to equation 1.

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Equation 1. E = hν =hcλ

The measurement of the hydrogen emission spectrum was one of the pivotal experiments leading to the quantized description of energy. Equations describing this spectrum have been derived from both experiment and theory. The Rydberg equation (equation 2) comes from spectral data describing a relationship between wavelength and some whole number, n.

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Equation 2. 1λ

= R 1nf

2 −1ni

2

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% &

'

( )

Another approach to understanding the hydrogen emission spectrum was introduced by Niels Bohr. He assumed the energy states of an electron could only have certain allowed values. In other words, the energies of an electron are quantized. Another assumption made by Bohr, which contradicted classical physics, was that an electron could orbit the nucleus without losing energy. This would explain how the electron, a charged particle would not radiate energy and eventually collapse into the nucleus. In Bohr’s model, the quantum number n relates to the radius of the electron’s orbit. From his assumptions and theoretical treatment of an electron orbiting a nucleus, he theoretically (and successfully) derived an equation (equation 3) describing the energy levels in a hydrogen atom.

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Equation 3. En = −2.179x10−18 J 1n2#

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&

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In today’s experiment you will measure the wavelengths of three lines (red, turquoise, and purple) emitted from a hydrogen lamp. The measured wavelengths will be used to derive Rydberg’s constant (R). Once Rydberg’s constant has been obtained, you will use the Rydberg equation (equation 2) to calculate the value of nf for the turquoise line. You will then compare experimental results with theoretical calculations obtained using Bohr’s equation.

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Procedure 1. Calibrate your spectroscope using the Hg lamp available in the lab.

a. Line up the slit on spectroscope until you see the thin line spectrum. Thick lines are probably scattered light from above. Read your spectroscope to the limit of the device. Record the wavelength of the thin Kelly green line.

b. Calculate the calibration factor for your spectroscope using equation 4. This factor will be added to each of your measurements for the hydrogen emission spectrum.

Equation 4. Calibration factor = 545 nm – (wavelength of your green Hg line)

2. Measure and record the wavelength of the three observed narrow lines in the hydrogen

emission spectrum. The purple line is a combination of two lines very close together. The spectroscopes used in this lab cannot distinguish between these two lines.

3. Apply your calibration factor using equation 5 to calculate the corrected value for the three

wavelengths observed.

Equation 5. Corrected Wavelength = Observed Wavelength + Calibration Factor

About the spectroscope The spectroscope is a device used in analyzing light for its component wavelengths. Light enters the slit and is broken into the component wavelengths by the diffraction grating. The grating serves as a “prism”, but works on a different principle. The result is projected against a scale at the rear of the spectroscope. The scale gives the wavelength value for each component. The numbers on the scale represent hundreds of nanometers. Each of these major divisions is divided into ten minor units. Each of the smallest divisions is worth 10 nm. Look through the diffraction grating at the narrow end of the spectroscope while lining up the slit with the emission lamp to be analyzed. The spectrum will appear on the right side of the slit below the scale where the emission wavelengths can be read. If the light source is too bright, you may observe multiple overlapping spectra on the right. If this is the case, move the scope from side to side until you find a position that yields a single spectrum. The visibility of the spectrum can be improved by holding the hand around the narrow end of the scope using the thumb and index finger to keep stray light from entering your eye

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Data Calibration Factor _____

Data Red Line Turquoise Purple

Observed Wavelength (nm)

Corrected Wavelength (nm)

Frequency (Hz)

Energy (J)

Data Workup and Calculations 1. Using the Rydberg equation and your corrected experimental wavelength for the red line, the

assumption that the red line corresponds to the transition (quantum jump) from ni = 3 to nf = 2, calculate the value of Rydberg’s constant (R). Make sure you have the correct units.

2. Using the Rydberg equation and your corrected experimental wavelength for the turquoise line, the assumption that the turquoise line corresponds to the transition (quantum jump) from ni = 4 to nf = 2, calculate the value of Rydberg’s constant. Make sure you have the correct units.

3. Calculate the average value of Rydberg’s constant obtained in the previous two problems. 4. Calculate the percent error for your experimentally determined Rydberg constant. 5. Using your experimentally determined average value of Rydberg’s constant and the corrected

experimental wavelength of the purple line (lpurple), calculate the value of ni. The assumption is nf = 2.

6. Using the Bohr equation, calculate the theoretical value for the energy of an electron in n2 of

the hydrogen atom. 7. Using the Bohr equation, calculate the theoretical value for the energy of an electron in n5 of

the hydrogen atom. 8. Calculate the change in energy (DE) of the electron as it “falls” from n5 to n2 using the values

for E5 and E2 obtained in the two previous problems. 9. Using DE from the previous problem, determine the frequency of the light given off as an

electron moves from n5 to n2. 10. Using the frequency calculated in the previous problem, determine the wavelength of the light

given off as an electron moves from n5 to n2. 11. Calculate the percent error between your corrected experimental results for lpurple and the

theoretical value calculated in the previous problem.

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17

Flame Spectra of Metal Cations

Introduction Each element has a unique electronic structure. The specific energy levels of all electrons within an element will affect how the element absorbs or emits energy. These unique absorptions and emissions can be used to distinguish one element from another. In this experiment a variety of metal cations (Na+, K+, Ca2+, Ba2+, Sr2+) will be observed in order to distinguish them from one another. A Bünsen burner will be used to heat aqueous solutions of the different metal cations. Due to the quantized electronic energy levels of these metal cations, electromagnetic radiation of only certain frequencies will be emitted when the excited valence electrons “fall” to lower energy states. The four cations to be used in this experiment will emit distinctively different spectra of light due to the difference in their electronic structures. You will observe and record the flame spectra of each metal cation and will be asked to identify an unknown sample by comparing it to your observations. Procedure 1. Obtain a platinum wire to be used to place the various metal solutions into the flame. 2. Half fill a well of your 24 well tray with 6M HCl. 3. Turn on the Bünsen burner and adjust the flame properly. 4. To clean the platinum wire, dip it into the 6M HCl, and then place into the Bünsen burner

flame for 30 – 60 seconds. Another technique used to clean the wrie is to soak in deionized water and clean with a paper towel.

5. Place a small amount of the following samples into separate well of your 24 well tray. a. BaCl2 b. CaCl2 c. SrCl2 d. KCl e. NaCl

6. Obtain an unknown sample from the instructor. 7. Dip the platinum wire into the BaCl2(aq) solution. Heat the solution by placing the wire into

the flame. Observe and record you observations. 8. Clean the platinum wire according to step 4. Test the wire to ensure it is clean. 9. Repeat step 7 for the other samples (including your unknown). Test the NaCl solution last.

The spectrum of NaCl is very strong and tends to contaminate other tests. 10. Dispose of all solutions as instructed by your instructor.

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Covalent Bonding and Molecular Structure Introduction This exercise is designed to help students better understand Lewis dot structure, chemical bonding, and molecular geometry. For a detailed Introduction please visit c4.cabrillo.edu/chem1a/ex1. Procedure 1. Go to c4.cabrillo.edu/chem1a/ex1 and view each model (grouped by classification). 2. Answer questions for each geometric classification.

a. The first question for each geometric classification asks for the determination of Lewis dot structure, electronic geometry, molecular geometry, perspective drawing, and orbital hybridiaiton for each structure listed. A suggested table to keep these answers organized is provided at the bottom of page 23. The exapmle used is for CO2.

b. Bond angles should be determined based on ideal angles. For example, ideal bond angles in tetrahedral molecules are 109.5°. However, when a bond angle deviates from ideality, the bond angle should be listed as greater than (>) or less than (<) 109.5°. In the case of an ideal bond angle of 180°, report a change as ¹ 180°.

c. Certain molecules used in this exercise contain many of the same atoms. So, to distinguish them, individual atoms will be labeled. The following notations are used:

XE: atom in an equatorial position XA: atom in an axial position X#: atom given a number to distinguish from others of the same element. Geometric Classifications AX2: CO2 1. Determine the following for CO2

a. Lewis dot structure b. Electronic geometry c. Molecular geometry d. Perspective drawing e. Orbital hybidization for each atom

2. What is the O–C–O bond angle? 3. Which element is more electronegative, carbon, or oxygen? 4. Would you expect CO2 to be a polar or a non-polar molecule? Explain. AX3: BF3, NO3

-, CO32-

1. Determine the following for BF3, NO3-, CO3

2- a. Lewis dot structure b. Electronic geometry c. Molecular geometry d. Perspective drawing e. Orbital hybidization for each atom

2. What is the F–B–F bond angle in BF3? 3. What is the O–N–O bond angle in NO3

-? 4. What is the O–C–O bond angle in CO3

2-? 5. Is the electron density in the B-F bond of BF3 oriented closer to fluorine, or to boron? 6. What is the average bond order for the N-O bonds in the nitrate ion? 7. What is the average bond order for the C-O bonds in the carbonate ion?

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AX2E: SO2, SnCl2 1. Determine the following for SO2, SnCl2

a. Lewis dot structure b. Electronic geometry c. Molecular geometry d. Perspective drawing e. Orbital hybidization for each atom

2. What is the O–S–O bond angle in SO2? 3. What is the Cl–Sn–Cl bond angle in SnCl2? 4. Which bond is longer, the S–O bond in SO2 or the Sn–Cl bond in SnCl2? Explain. 5. Which bond angle would you expect to be smaller, O–S–O or Cl–Sn–Cl? Explain. AX4: CH4, NH4

+, CH3Cl 1. Determine the following for CH4, NH4

+, CH3Cl a. Lewis dot structure b. Electronic geometry c. Molecular geometry d. Perspective drawing e. Orbital hybidization for each atom

2. What is the H–C–H bond angle in CH4? 3. What is the H–N–H bond angle in NH4

+? 4. What is the H–C–Cl bond angle in CH3Cl? 5. What is the H–C–H bond angle in CH3Cl? 6. Which one of these molecules is polar? Using your perspective drawing, draw an arrow

toward the more negative side of the molecule. 7. In CH3Cl, which bond is longer, C–H or C–Cl? AX3E: NF3, NH3, H3O+ 1. Determine the following for NF3, NH3, H3O+

a. Lewis dot structure b. Electronic geometry c. Molecular geometry d. Perspective drawing e. Orbital hybidization for each atom

2. What is the F–N–F bond angle in NF3? 3. What is the H–N–H bond angle in NH3? 4. What is the H–O–H bond angle in H3O+? 5. Would you expect the electron density from the lone pair to be closer to the nitrogen in NF3,

or in NH3? Explain. 6. Which bond angle do you expect to be smaller, F–N–F in NF3 or H–O–H in H3O+? Explain. AX2E2: H2O, CH3–O–CH3 1. Determine the following for H2O, CH3–O–CH3

a. Lewis dot structure b. Electronic geometry (for each “central” atom) c. Molecular geometry (for each “central” atom) d. Perspective drawing e. Orbital hybidization for each atom

2. What is the H–O–H bond angle in H2O? 3. What is the C–O–C bond angle in CH3-O-CH3? 4. What is the H–C–H bond angle in CH3-O-CH3? 5. Do you expect the bond angle H–O–H to be smaller in H2O, or in H3O+? Explain.

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AX5: PF5, SOF4 1. Determine the following for PF5, SOF4

a. Lewis dot structure b. Electronic geometry c. Molecular geometry d. Perspective drawing e. Orbital hybidization for each atom

2. What is the FA–P– FA bond angle in PF5? 3. What is the FA–P– FE bond angle in PF5? 4. What is the FE–P– FE bond angle in PF5? 5. What is the FE–S– O bond angle in SOF4? 6. What is the FA–S– O bond angle in SOF4? 7. What is the FE–S– FE bond angle in SOF4? 8. What is the FE–S– FA bond angle in SOF4? 9. Is the oxygen in SOF4 located in an axial or an equatorial position? Why is this? 10. Which bond in SOF4 has the shortest bond length? AX4E: SF4, IF4

+ 1. Determine the following for SF4, IF4

+ a. Lewis dot structure b. Electronic geometry c. Molecular geometry d. Perspective drawing e. Orbital hybidization for each atom

2. What is the FA–S– FA bond angle in SF4? 3. What is the FA–S– FE bond angle in SF4? 4. What is the FE–S– FE bond angle in SF4? 5. What is the FA–I– FA bond angle in IF4

+? 6. What is the FA–I– FE bond angle in IF4

+? 7. What is the FE–I– FE bond angle in IF4

+? 8. Does the lone pair of electrons in SF4 occupy an axial or an equatorial position? Explain. AX3E2: ClF3 1. Determine the following for ClF3

a. Lewis dot structure b. Electronic geometry c. Molecular geometry d. Perspective drawing e. Orbital hybidization for each atom

2. What is the FA–Cl– FA bond angle in ClF3? 3. What is the FA–Cl– FE bond angle in ClF3? 4. Why are you not being asked to predict the FE-Cl-FE bond angle? 5. Which bond do you expect to be longer, FA–Cl or FE–Cl? Explain.

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AX2E3: I3-, IF2-

1. Determine the following for I3-, IF2-

a. Lewis dot structure b. Electronic geometry c. Molecular geometry d. Perspective drawing e. Orbital hybidization for each atom

2. What is the I–I–I bond angle in I3-? 3. What is the F–I–F bond angle in IF2

-? 4. Which bond length do you expect to be longer, the I–I bond in I3- or the F–I one in IF2

-? Explain.

5. Why doesn't either fluorine atom in IF2- occupy an equatorial position?

AX6: SF6, IOF5 1. Determine the following for SF6, IOF5

a. Lewis dot structure b. Electronic geometry c. Molecular geometry d. Perspective drawing e. Orbital hybidization for each atom

2. What is the F1–S–F2 bond angle in SF6? 3. What is the F1–S–F3 bond angle in SF6? 4. What is the O–I–F1 bond angle in IOF5? 5. What is the F1–I–F2 bond angle in IOF5? 6. What is the F1–I–F5 bond angle in IOF5? 7. What is the F2–I–F5 bond angle in IOF5? 8. Are all of the fluorine atoms in SF6 equivalent? Why? 9. Would you expect angles F1–I–F2 and F2–I–F3 in IOF5 to be equal? Why or why not? AX5E: BrF5, TeF5

- 1. Determine the following for BrF5, TeF5

- a. Lewis dot structure b. Electronic geometry c. Molecular geometry d. Perspective drawing e. Orbital hybidization for each atom

2. What is the F1–Br–F2 bond angle in BrF5? 3. What is the F1–Br–F3 bond angle in BrF5? 4. What is the F1–Br–F5 bond angle in BrF5? 5. What is the F1–Te–F2 bond angle in TeF5

-? 6. What is the F1–Te–F3 bond angle in TeF5

-? 7. What is the F1–Te–F5 bond angle in TeF5

-? 8. Which bond in BrF5, Br–F1 or Br–F5, do you expect to be longer? Explain. 9. Would you expect bond Br–F1 in BrF5 to have a bond length that is shorter, longer, or the

same as the length of bond Te–F1 in TeF5-? Explain.

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AX4E2: ICl4- 1. Determine the following for ICl4-

a. Lewis dot structure b. Electronic geometry c. Molecular geometry d. Perspective drawing e. Orbital hybidization for each atom

2. What is the Cl1–I–Cl2 bond angle in ICl4-? 3. What is the Cl1–I–Cl3 bond angle in ICl4-? 4. Why are the two lone pairs ICl4- on opposite vertices of the molecule? 5. Would you expect Cl1–I to have a shorter, longer, or the same bond length as Cl2–I? Explain. Table 1. Answers to question 1 for Carbon Dioxide

Sample CO2

Lewis dot structure O C O

Electronic geometry Linear

Molecular geometry Linear

Perspective Drawing O C O

Orbital hybridization (carbon) sp3

Orbital hybridization (oxygen) sp2

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Solutions and Reactions Introduction and Procedure in Catalyst Data and Workup Part A. Naming and Making Solutions Part B. Solubility and Solutions 1. Make a table in your notebook with fourteen rows (one for each solution you will be working

with) and five columns (name, formula, concentration, description, and observation when mixed with Universal Indicator)

Part C. Solutions and Reactions 1. Record observations in your notebook for this section. Pay particular attention to the

following steps. a. Step 2 (before adding water) b. Step 2 (after adding water) c. Step 5 (after pushing both solids into water)

2. Write the Molecular, Total Ionic, and Net Ionic equation for the reaction between aqueous

Pb(NO3)2 and KI.

Part D. Four Major Types of Chemical Reaction in Aqueous Solution. (OMIT) Part E. A Chemical Reaction Survey: The Ion Reaction Chart There is a chart at the end of the chapter in the Catalyst Lab Manual for use in recording all observations. You may need to use abbreviations, so make sure any abbreviations you use make sense to you. These notes will be your only source of information in part G, so take good notes. Part F. Identification of Three Unknowns 1. Record your unknown number. 2. Record observation in your lab notebook. 3. Make a results table (see below) in your notebook

Part G. Five Unknowns: Solo 1. Record your unknown number. 2. Record observation in your lab notebook. 3. Make a results table (see below) in your notebook

26

27

The Use and Abuse of Aluminum Introduction and Procedure in Catalyst Data and Workup Part A. The Aluminum Can The following diagrams (inside and outside of the can) should be drawn into your lab notebook in order to record your observations after treating the can with acid and base.

KOHObservationsH2SO4

Observations

KOHObservations

H2SO4 Observations

Outside of can

KOHObservationsH2SO4

Observations

KOHObservations

H2SO4 Observations

Inside of can

28

Part B. Recycling Aluminum: The Synthesis of Alum 1. Record the mass of the aluminum foil used in step 1. 2. Record your observations as the aluminum foil is treated with 1.4M KOH. 3. Write the chemical reaction that is taking place? 4. Record your observations following the addition of 9M H2SO4 in step 16.

a. There are two reactions that are responsible for your observations, provide both of these reactions.

5. Determine the yield of crystals. a. What is the theoretical yield of alum based on the initial amount of aluminum

used in the experiment? Provide all the appropriate calculations in your lab notebook.

b. Calculate the experimental percent yield? Part C. Qualitative Analysis of an Alum Sample

1. In your lab notebooks record all data in table format that lead to verification of sulfate ion (SO4

2-), potassium ion (K+), water, and aluminum ion (Al3+). Part D. A Practical Use for the Reaction for Aluminum with a Base

1. In your lab notebook make illustrations showing the different morphological shapes you see in the Crystal Drano®.

2. Record any changes that occur over time before the addition of water. 3. Record what happens when water is added to Crystal Drano®. Observe for a few

minutes.

29

Standardization of a Sodium Hydroxide Solution

Introduction Standard solutions are solutions with known concentrations. There are two different ways to make a standard solution. A standard solution can be prepared directly by dissolving a known mass of solid with enough solvent to make a specific volume of solution. An alternate and sometimes necessary technique is to dissolve an approximate amount of solid sample into a volume of solvent, then determining it’s exact concentration through the use of titration. This second technique is used when preparing a solution of sodium hydroxide (NaOH). This is necessary because NaOH is hygroscopic (absorbs moisture from the air). The absorption of moisture from the air makes it impossible to obtain an accurate mass determination using a balance since the mass of the sample changes. In an acid-base titration experiment, the number of moles of base (in this case NaOH) can be determined by neutralizing them with a known number of moles of acid (in this case oxalic acid, H2C2O4). In this experiment, the acid-base neutralization occurs according to the following reaction.

2 NaOH(aq) + H2C2O4(aq) ® 2 H2O(l) + Na2C2O4(aq) The number of moles of oxalic acid can be determined using a balance since oxalic acid is not hygroscopic and will not absorb water from the air. Notice the stoichiometry between NaOH and H2C2O4 is NOT 1:1. For every one mole of H2C2O4 two moles of NaOH will be neutralized. This will come into play in your calculations in the near future. One other important feature of H2C2O4 is that in the solid form it has two water molecules attached to it, and is referred to oxalic acid dihydrate H2C2O4•2H2O. This is important to recognize because for every one molecule of oxalic acid being weighed on the balance, two waters are also being weighed. Below is a sample calculation showing how to deal with this concept. Sample Calculation, assume we want to neutralize 0.050 moles of NaOH

0.050 moles NaOH 1 mole H2C2O4

2 mole NaOH!

"#

$

%&= 0.025 moles H2C2O4

0.025 moles H2C2O41 mole H2C2O4 ⋅2H2O

1 mole H2C2O4

!

"#

$

%&= 0.025 moles H2C2O4 ⋅2H2O

0.025 moles H2C2O4 ⋅2H2O126.068g H2C2O4 ⋅2H2O1 mole H2C2O4 ⋅2H2O

!

"#

$

%&= 3.2g H2C2O4 ⋅2H2O

so, if you want to neutralize 0.050 moles of NaOH, you will need to use 3.2 grams of H2C2O4×2H2O.

30

Procedure 1. Prepare approximately 500 mL of approximately 0.25M NaOH.

a. First, calculate the number of grams of NaOH needed to prepare this solution (show in “Data and Workup” section).

b. Weigh approximately this amount out and place it into a plastic 500 mL bottle. c. Add enough DI water to bring the volume close to the 500 mL mark. d. Since the solution being made will be standardized it is not necessary to be exact in your

measurements at this time. 2. Prepare three Erlenmeyer flasks each containing a precisely measured amount of

H2C2O4×2H2O. a. Calculate the amount of oxalic acid dihydrate needed to neutralize 25 mL of 0.25M NaOH

(show in “Data and Workup” section). b. Weigh approximately this amount into the three separate Erlenmeyer flasks. It is not

necessary to weigh out the exact amount calculated, but it is CRITICAL that you precisely record the values.

c. Add approximately 50 mL of DI water to each of the flasks and dissolve the acid. The amount of water is not critical because it is the acid molecules present involved in the reaction.

3. Prepare a buret to load your NaOH solution to be standardized. a. Rinse the buret with DI water 2-3 times. Do this by coating the walls of the buret with

water and then pouring the water out. b. Repeat previous step using your NaOH solution instead of water. c. Fill the buret with your NaOH solution. BE CAREFUL!

4. Place one of the Erlenmeyer flasks containing acid underneath the buret. 5. Add 2-3 drops of phenolphthalein indicator to the flask containing the acid. 6. Record the initial reading on the buret. 7. Dispense the NaOH solution into the flask until the acid in the flask is neutral. 8. Record the final reading on the buret 9. If necessary refill the buret with NaOH and repeat steps 4 – 8 using the other two Erlenmeyer

flasks. 10. Store your NaOH solution properly until next lab when it will be used to determine the

concentration of a solution of hydrochloric acid (HCl). 11. Calculate the [NaOH] for each of the three runs. 12. Determine the average concentration of your three runs.

31

Determine Molarity of a Hydrochloric Acid Solution

Introduction In this exercise using a titration you will determine the concentration of a hydrochloric acid (HCl) solution with an unknown molarity using a standardized sodium hydroxide (NaOH) solution prepared a previous lab. The neutralization reaction between HCl and NaOH proceeds according to the chemical equation shown below.

HCl(aq) + NaOH(aq) ® H2O(l) + NaCl(aq) This reaction is often written in one of the following forms showing the net ionic equation.

H3O+(aq) + OH-(aq) ® 2 H2O(l)

or

H+(aq) + OH-(aq) ® H2O(l) Since the reaction between HCl and NaOH is 1:1, the reaction is at the neutral point when there is one mole of NaOH for every one mole of HCl.

Equation 1. nNaOH = nHCl In this equation, “n” represents the number of moles of either NaOH or HCl. During the titration you will be carefully measuring the volume of NaOH (VNaOH) required to neutralize the unknown acid solution. Since the volume and concentration of the NaOH (CNaOH) are known, the number of moles of NaOH can be calculated using equation 2.

Equation 2. nNaOH = CNaOH( ) VNaOH( ) Equation 3 shows this relationship for HCl.

Equation 3. nHCl = CHCl( ) VHCl( ) Using equations 1 – 3 and some simple algebra, equation 4 can be derived, which allows for the calculation of the hydrochloric acid concentration.

Equation 4. CHCl = CNaOH( ) VNaOH( )VHCl( )

In the experiment the volume of HCl is known (set at 20.00 mL), the molarity of NaOH was determined in the previous lab period, and the volume of NaOH will be measured. The only unknown in equation 4 is the concentration of HCl.

32

Procedure 1. Clean and load a buret with the standardized NaOH solution from previous lab.

a. Rinse the buret with DI water 2-3 times. Do this by coating the walls of the buret with water and then pouring the water out.

b. Rinse the buret with your NaOH solution 2-3 times. c. Fill the buret with your NaOH solution. BE CAREFUL!

2. Dispense 20.00 mL of the HCl solution using a volumetric pipette into an Erlenmeyer flask. 3. Place the Erlenmeyer flasks containing acid underneath the buret. 4. Add 2-3 drops of phenolphthalein indicator to the flask containing the acid. 5. Record the initial reading on the buret. 6. Dispense the NaOH solution into the flask until the end point is reached. 7. Record the final reading on the buret 8. If necessary refill the buret with NaOH and repeat steps 2 – 7 using two different clean

Erlenmeyer flasks. 9. Calculate the concentration of your hydrochloric acid solution.

33

Determining the Molar Volume of an Ideal Gas and the Universal Gas Constant

Introduction According to the ideal gas laws, all gases (which behave ideal) will occupy the same volume under identical conditions (temperature, pressure, and moles of gas). Hydrogen gas is probably the most ideal gas we can study. The molar volume (L/mol) of hydrogen gas will be determined using a small chemical reaction producing hydrogen. In order to determine the molar volume of a gas it is necessary to determine the volume occupied by a given number of moles of gas. The volume will be measured using an inverted graduated cylinder. The number of moles of gas will be calculated using the balanced chemical equation shown below.

Mg(s) + 2 HCl(aq) ®® MgCl2(aq) + H2(g) If we start with a known amount of magnesium, we can calculate the theoretical number of moles of hydrogen gas that will be produced. This reaction is considered to go to completion, so it is assumed the theoretical yield will also be the actual yield. Procedure 1. Fill a 400 mL beaker to 75% capacity with tap water (DI water is not necessary). 2. Allow the water to reach room temperature 3. Obtain a piece of magnesium ribbon (mass between 0.005g – 0.009g). Two pieces may be

used if necessary. Record its mass. 4. Wrap a piece of copper wire around the magnesium (making a “cage”) so the magnesium will

not be separated from the copper if placed into a solution. Your instructor will demonstrate this.

5. With the other end of the copper wire, thread it through a stopper with a single hole and fasten it the to stopper.

6. Wash a 10 mL graduated cylinder and do two final rinses with DI water. 7. Very carefully pour approximately 3 mL of 3M HCl along the side of the graduated cylinder

being careful not to splash it around in the cylinder. 8. Using a plastic DI water bottle, gently add DI water to the top trying not to agitate the acid

layer from the previous step. This is for two reasons. a. First, it is not good lab procedure to add water on top of acid, but if we can prevent mixing

then we can also prevent heat being generated that accompanies this mixing. b. Second, we want the hydrochloric acid to remain at the bottom of the cylinder so we can

control the initiation of the reaction. 9. Place the stopper containing the copper wire and magnesium into the top of the graduated

cylinder. Water should come out through the opening in the stopper, and no air should be present inside the cylinder. If air exists inside the cylinder, add more water and try again. For best results, the magnesium should sit approximately 1 cm – 1.5 cm above the stopper.

10. As soon as you turn the graduated cylinder upside down the reaction will begin. 11. Place you finger over the opening in the stopper (to prevent water from coming out when

cylinder is flipped). Quickly (but carefully) invert the cylinder and place into the beaker containing the water so the stopper end of the graduated cylinder is completely submerged under water.

12. After three minutes the reaction should be complete. Make sure all the solid magnesium is gone, and that no more gas bubbles are being formed. If there are gas bubbles stuck on the copper wire, gently tap the cylinder to remove the bubbles.

34

13. Adjust the level of the liquid in the graduated cylinder to be even with the level of the water in the beaker. Record the volume of gas produced.

14. Record the temperature of the water. 15. Repeat the experiment with a second sample of magnesium ribbon.

35

Table 1. Vapor Pressure of Water at Various Temperatures Temperature

(°C) Vapor Pressure

(torr) Temperature

(°C) Vapor Pressure

(torr) Temperature

(°C) Vapor Pressure

(torr) 15.0 12.8 19.4 16.9 23.8 22.1 15.1 12.9 19.5 17.0 23.9 22.2 15.2 13.0 19.6 17.1 24.0 22.4 15.3 13.1 19.7 17.2 24.1 22.5 15.4 13.2 19.8 17.3 24.2 22.6 15.5 13.2 19.9 17.4 24.3 22.8 15.6 13.3 20.0 17.6 24.4 22.9 15.7 13.4 20.1 17.7 24.5 23.1 15.8 13.5 20.2 17.8 24.6 23.2 15.9 13.6 20.3 17.9 24.7 23.3 16.0 13.7 20.4 18.0 24.8 23.5 16.1 13.8 20.5 18.1 24.9 23.6 16.2 13.8 20.6 18.2 25.0 23.8 16.3 13.9 20.7 18.3 25.1 23.9 16.4 14.0 20.8 18.4 25.2 24.0 16.5 14.1 20.9 18.5 25.3 24.2 16.6 14.2 21.0 18.7 25.4 24.3 16.7 14.3 21.1 18.8 25.5 24.5 16.8 14.4 21.2 18.9 25.6 24.6 16.9 14.5 21.3 19.0 25.7 24.8 17.0 14.6 21.4 19.1 25.8 24.9 17.1 14.7 21.5 19.2 25.9 25.1 17.2 14.7 21.6 19.4 26.0 25.2 17.3 14.8 21.7 19.5 26.1 25.4 17.4 14.9 21.8 19.6 26.2 25.5 17.5 15.0 21.9 19.7 26.3 25.7 17.6 15.1 22.0 19.8 26.4 25.8 17.7 15.2 22.1 20.0 26.5 26.0 17.8 15.3 22.2 20.1 26.6 26.1 17.9 15.4 22.3 20.2 26.7 26.3 18.0 15.5 22.4 20.3 26.8 26.4 18.1 15.6 22.5 20.4 26.9 26.6 18.2 15.7 22.6 20.6 27.0 26.7 18.3 15.8 22.7 20.7 27.1 26.9 18.4 15.9 22.8 20.8 27.2 27.1 18.5 16.0 22.9 20.9 27.3 27.2 18.6 16.1 23.0 21.1 27.4 27.4 18.7 16.2 23.1 21.2 27.5 27.5 18.8 16.3 23.2 21.3 27.6 27.7 18.9 16.4 23.3 21.5 27.7 27.9 19.0 16.5 23.4 21.6 27.8 28.0 19.1 16.6 23.5 21.7 27.9 28.2 19.2 16.7 23.6 21.8 28.0 28.3 19.3 16.8 23.7 22.0 28.1 28.5

36

37

Halogens and Their Compounds

Introduction and Procedure in Catalyst Data and Workup Part A. From Fluorine to Astatine: A Basic Introduction to the Halogens There is no data in this section. Answer the following questions in your lab notebook.

1. Write the electron configuration for a fluorine atom. 2. Write the electron configuration for a fluoride ion. 3. Write the electron configuration for a chlorine atom.

a. Circle the valence electrons in the electron configuration written for chlorine. 4. Write the short hand electron configuration for a chlorine atom. 5. Write the short hand electron configuration for a bromine atom. 6. Draw the Lewis dot structure for an iodine atom. 7. Draw the Lewis dot structure for an iodine molecule. 8. Which of the halogens is the most electronegative? 9. Write a balanced net ionic equation for the reaction between bromide and chlorine. 10. Write a balanced net ionic equation for the reaction between chloride and bromine. 11. Write the reaction to show what would occur if aluminum metal was dropped into fluorine

gas. Part B. The Synthesis and Reactions of Chlorine The following diagram should be drawn into your lab notebook in order to record your observations and each chemical reaction occurring within the different drops. Write both the full chemical equation, and net ionic equations.

bleach

filter paper (0.1M KI)Before:

After:

0.1M KIBefore:

After:

starch/0.1M KIBefore:

After:

bromocrescol greenBefore:

After:

Part C. A Small-Scale Pilot Plant for the Manufacture of Chlorine by the Industrial Process

1. Record all observations at both the anode and cathode. 2. Write the half reactions occurring at the anode and cathode.

38

Part D. Electrochemical Writing with a Halogen 1. Record all observations. 2. Write the net ionic equation for both reactions occurring.

Part E. Precipitation Reactions and the Titrations of a Halide

1. Construct a table and record your observations and the net ionic equation for each of the mixtures in step 1.

2. Construct a table and record your observations for the AgNO3 titration of NaCl. 3. Calculate the [NaCl] of the unknown sample.

Part F. Redox Analysis of Commercial Bleach

1. Record observation and the number of drops of 0.01M Na2S2O3 required to reach the end point of the titration. The titration should be run twice.

2. Perform the following calculations in order to determine the percent of NaClO in bleach. a. Calculate the molarity of NaOCl in the diluted bleach sample. b. Calculate the molarity of NaOCl in the original bleach sample. c. Calculate the grams/milliter of NaClO in the original bleach sample. d. Calculate the percentage (grams/100mL) of NaClO in the original sample.

39

Chemistry of Natural Waters Introduction and Procedure in Catalyst Data and Workup Part A. The Evaporation of Water Samples to Give Total Dissolved Solids

1. Record observations for the TDS of multiple waters samples. Part B. Divalent Cation Analysis by EDTA

1. Construct a table to record your observations for the four serial titrations performed. Part C. The Dissolution of Rocks

1. Construct a table to record your observations for the pool titration of the limestone and granite pool.

2. Calculate the [divalent] for your limestone and granite pool samples.

3. Construct a table to record your observations for the pool titration of the cold water, hot water, and acidic limestone pool samples.

4. Calculate the [divalent] for each sample. Part D. Important Ways of Reporting the Hardness of Water

1. Convert 2.5 x 10-3M Ca2+ into ppm. Show all work. Part E. Determination of the Hardness of Groundwater, Spring Water, and Well Water

1. Construct a table to record your observations for the pool titration of three water samples. 2. Calculate the [divalent] for each water sample.

Part F. The Reaction of Divalent Cations with Soap; Soap Titration

1. Record observations in your laboratory notebook as you proceed through this section. 2. What is the chemical formula of the precipitate that forms with the addition of soap?

Part G. Water Softening with Commercial Water-Conditioning Agents

1. Record observations in your laboratory notebook as you proceed through this section.

Part H. Divalent Cation Removal by Ion Exchange 1. Record the amount of EDTA required to reach the end point after the well water was

treated with the ion exchange resin. 2. Calculate the hardness of water treated with the ion exchange resin.

40

41

Specific Heat of Metals Introduction The amount of heat energy required to raise the temperature of one gram of a substance by one degree Celsius is called the specific heat capacity, or simply the specific heat. Water for instance, has a specific heat (cp) of 4.184 joules per gram degree Celsius (4.184 J/g×°C). The amount of heat energy involved in changing the temperature of a sample of a particular substance depends on three parameters; the specific heat (cp) of the substance, the mass (m) of the sample, and the magnitude of the temperature change (DT). The Greek letter delta (D) is used to indicate a change.

€

Equation 1. ΔT = temperaturefinal −temperatureinitial The amount of heat energy transferred in the process of producing a temperature change can be calculated using equation 2.

€

Equation 2. q = (cp )(m)(ΔT)

In this experiment, the specific heat of a three known metals will be determined. A heated sample of metal will be placed into a coffee cup calorimeter containing room temperature water. Heat will be transferred from the metal to the water. Eventually, the water and the metal will reach the same temperature (reach equilibrium). Since styrofoam is a good insulator, heat cannot easily escape from the calorimeter to the surroundings. Therefore, the heat lost by the metal can be said, for the purpose of this experiment, to be equal to the heat gained by the water. The amount of heat energy gained by the water will be calculated using equation 3.

€

Equation 3. qwater = (cwater )(mwater)(ΔTwater ) The heat lost by the metal can be calculated using equation 4.

€

Equation 4. qmetal = (cmetal )(mmetal )(ΔTmetal ) Since the heat gained by the water must equal the heat lost by the metal, an equality can be established according to equation 5.

€

Equation 5. qwater = - qmetal Combing equations 3, 4, and 5 we arrive at Equation 6, which allows for the calculation of the specific heat of a metal using experimental measurements along with the specific heat of water.

€

Equation 6. (cwater )(mwater)(ΔT) = - (cmetal )(mmetal )(ΔTmetal ) The specific heat of water is known. The temperature changes of water and metal can be measured, as can the mass of the water and metal.

42

Table 1. Specific Heat for Common Metals

Metal

Specifc Heat (J/g×°C)

Metal

Specifc Heat (J/g×°C)

Lead 0.129 Iron 0.451

Copper 0.385 Nickel 0.44

Aluminum 0.900 Tin 0.22

Cobalt 0.46 Zinc 0.39 Procedure 1. Construct a data table in your laboratory notebook in order to record all necessary data for

the determination of the specific heat of three metals. 2. Prepare a hot water bath by placing a large beaker 75% full of water placed on a hot plate. 3. Obtain enough metal to fill a test tube about 1/4 full. 4. Weigh the metal and record the mass. 5. CAREFULLY transfer the metal shot to a large, DRY test tube. Be careful to pour the metal

shot into the tube slowly so the bottom of the test tube is not broken in this process. Place the test tube in the boiling water. Make sure the metal shot is below the level of water in the beaker. Allow the test tube to remain in the boiling water bath for at least 10 minutes. Proceed to step 6 while the metal shot is heating.

6. Using the balance, measure the mass of your empty coffee cup calorimeter. Pour approximately 100 mL of distilled water into the calorimeter. Re-weigh the calorimeter with the water to determine the mass of water placed into the calorimeter. Place the calorimeter into a 400 mL beaker for support.

7. After the metal shot has been heating for at least 10 minutes, using the thermometer provided in the boiling water baths, measure the temperature of the hot water. The temperature of the metal shot can be assumed to be the same as the water bath. Record this temperature to the nearest 0.5 °C as the initial temperature of the metal sample.

8. Using a temperature probe begin collecting temperature data of the water in the calorimeter. 9. Carefully remove the test tube from the bath. Carefully, but quickly, pour the metal shot into

the water in the styrofoam cup (use a paper towel to keep any hot water on the tube from dropping into the calorimeter).

10. Gently stir the calorimeter and its contents noting the temperature frequently. 11. Data collection can be stopped once the temperature of the water begins to decline. 12. Determine the maximum temperature of the water in the calorimeter and record this as the

final temperature of both water metal. 13. Repeat steps 3 – 12 for two other metals.

43

Additivity of Heats of Reaction: Hess’s Law

Introduction Hess’s law states “the change in enthalpy (DH) that occurs when reactants are converted to products in a reaction is the same whether the reaction takes place in one step or in a series of steps.” The change in enthalpy for the dissolution of solid sodium hydroxide will be determined using two techniques. First, by directly measuring the heat released when NaOH(s) is dissolved in water (Reaction 1). Second, by determining the heat of two separate reactions (Reaction 2 and Reaction 3), then applying Hess’s law.

1. Solid NaOH dissolves in water to form aqueous NaOH.

NaOH(s) ® NaOH(aq) DH1 = ?

2. Solid NaOH reacts with aqueous HCl to form water and an aqueous NaCl. NaOH(s) + HCl(aq) ® H2O(l) + NaCl(aq) DH2 = ?

3. Aqueous NaOH and aqueous HCl react to form water and aqueous NaCl.

NaOH(aq) + HCl(aq) ® H2O(l) + NaCl(aq) DH3 = ? Using Hess’s law, it can be shown that reaction 2 and reaction 3 can be arranged to sum to reaction 1. Determine how reaction 2 and reaction 3 should be arranged to sum to reaction 1. Procedure 1. Prepare computer to collect temperature data for 1000 seconds. Reaction 1 2. Make a coffee cup calorimeter by placing two styrofoam cups into a 250-mL beaker. 3. Place calorimeter on stir plate. Place a small stir bar into calorimeter. READ STEPS 4 – 7 COMPLETLY BEFORE PROCEEDING. THESE STEPS MUST BE DONE WITHOUT MUCH DELAY. 4. Measure approximately 55 mL of deionized water the styrofoam cup. Determine the mass of

water used and record in data table. 5. Use a utility clamp to suspend the temperature probe from a ring stand. Place the

temperature probe into water. 6. Weigh out about 1 gram of NaOH(s). Since NaOH(s) is hygroscopic it will be difficult to

obtain the precision we are used to, so record the mass to the nearest 0.01 g. It is necessary to weigh it and proceed to the next step without delay.

7. Turn the stirring mechanism on low and begin collecting temperature data. After roughly ten seconds, add NaOH(s) to the water. Continue to collect date until a maximum temperature has been reached and the temperature starts to drop. Click “stop” to end data collection.

8. Record necessary data into you laboratory notebook. 9. Discard the solution into the aqueous waste.

44

Reaction 2 10. Make approximately 55 mL of 0.50 M HCl by diluting 1.00 M HCl. Allow the diluted

solution to cool down for 5 minutes prior to moving onto the next step. 11. Repeat Steps 1-9 using the 55 mL of 0.50 M hydrochloric acid (HCl) instead of water. Reaction 3 12. Repeat Steps 1-9, initially measuring approximately 55 mL of 1.0 M HCl (instead of water)

into the styrofoam calorimeter. In Step 6, instead of solid NaOH, measure 50.0 mL of 1.0 M NaOH solution into a graduated cylinder. Make sure to determine the mass of both solutions used. Assume the density of NaOH(aq) = 1.00 g/mL

Workup 1. Using data collected from Reaction 1, calculate DH1 (per mole NaOH). 2. Using data collected from Reaction 2, calculate DH2 (per mole NaOH). 3. Using data collected from Reaction 3, calculate DH3 (per mole NaOH). 4. Calculate the change in enthalpy for the dissolution of NaOH(s) using Reaction 2 and

Reaction 3.

!

Pe

rio

dic

Ta

ble

of E

lem

en

ts

1

1A

18

8A

1

H

1.0

08

2

2A

13

3A

14

4A

15

5A

16

6A

17

7A

2

He

4.0

03

3

Li

6.9

41

4

Be

9.0

12

5

B

10

.81

6

C

12

.01

7

N

14

.01

8

O

16

.00

9

F

19

.00

10

Ne

20

.18

11

Na

22

.99

12

Mg

24

.31

3

3B

4

4B

5

5B

6

6B

7

7B

8

8B

9

8B

10

8B

11

1B

12

2B

13

Al

26

.98

14

Si

28

.09

15

P

30

.97

16

S

32

.07

17

Cl

35

.45

18

Ar

39

.95

19

K

39

.10

20

Ca

40

.08

21

Sc

44

.96

22

Ti

47

.88

23

V

50

.94

24

Cr

52

.00

25

Mn

5

4.9

4

26

Fe

55

.85

27

Co

58

.93

28

Ni

58

.69

29

Cu

63

.55

30

Zn

6

5.3

9

31

Ga

69

.72

32

Ge

72

.59

33

As

74

.92

34

Se

78

.96

35

Br

79

.90

36

Kr

83

.80

37

Rb

85

.47

38

Sr

87

.62

39

Y

88

.91

40

Zr

91

.22

41

Nb

92

.91

42

Mo

95

.94

43

Tc

(98

)

44

Ru

10

1.0

7

45

Rh

10

2.9

1

46

Pd

10

6.4

2

47

Ag

10

7.8

7

48

Cd

11

2.4

1

49

In

11

4.8

2

50

Sn

11

8.7

1

51

Sb

12

1.7

6

52

Te

12

7.6

0

53

I 1

26

.90

54

Xe

13

1.2

9

55

Cs

13

2.9

1

56

Ba

1

37

.33

71

Lu

17

4.9

7

72

Hf

17

8.4

9

73

Ta

18

0.9

5

74

W

18

3.8

4

75

Re

18

6.2

1

76

Os

19

0.2

3

77

Ir 1

92

.22

78

Pt

19

5.0

8

79

Au

19

6.9

7

80

Hg

20

0.5

9

81

Tl

20

4.3

8

82

Pb

20

7.2

83

Bi

20

8.9

8

84

Po

(20

9)

85

At

(21

0)

86

Rn

(22

2)

87

Fr

(22

3)

88

Ra

(22

6)

10

3

Lr

(26

2)

10

4

Rf

(26

7)

10

5

Db

(26

8)

10

6

Sg

(27

1)

10

7

Bh

(2

72

)

10

8

Hs

(27

0)

10

9

Mt

(27

6)

11

0

Ds

(28

1)

11

1

Rg

(28

0)

57

La

13

8.9

1

58

Ce

14

0.1

2

59

Pr

14

0.9

1

60

Nd

14

4.2

4

61

Pm

(1

45

)

62

Sm

1

50

.36

63

Eu

15

1.9

6

64

Gd

15

7.2

5

65

Tb

1

58

.93

66

Dy

16

2.5

0

67

Ho

16

4.9

3

68

Er

16

7.2

6

69

Tm

1

68

.93

70

Yb

1

73

.05

8

9

Ac

(22

7)

90

Th

23

2.0

4

91

Pa

23

1.0

4

92

U

23

8.0

3

93

Np

(23

7)

94

Pu

(24

4)

95

Am

(2

43

)

96

Cm

(2

47

)

97

Bk

(24

7)

98

Cf

(25

1)

99

Es

(25

2)

10

0

Fm

(2

57

)

10

1

Md

(25

8)

10

2

No

(25

9)

!

CHEMISTRY LABORATORY CONDUCT

The following regulations have been compiled to promote safety and efficiency. Failure to observe them scrupulously can place you and others at risk.

GENERAL REGULATIONS

1. Work is permitted in the laboratory only when there is an instructor in charge. You are urged to perform your experiments during regular laboratory sessions. Consult your instructor if you want to work in the laboratory at other times. 2. Carry out all experiments independently unless directed otherwise. 3. Record your individual observations directly on the data sheets. 4. Follow directions carefully. This is important for your own safety and the safety of others. 5. Read the assigned experiment before coming to lab. Be prepared to ask questions about any part of the experiment that is not clear to you. If it becomes apparent that you are not prepared to perform the experiment, you may be asked to leave.

SAFETY AND FIRST AID 1. Safety goggles are to be worn at all times in the lab. It is not only what you are doing but also what your neighbors are doing that could threaten your eyes. If you get chemicals in your eyes, flush your eyes out immediately with one of the emergency eyewashes at either end of the lab. Notify your instructor and rinse eyes for a FULL FIFTEEN MINUTES. Hold eyelids open with a thumb and forefinger and roll eyes while rinsing. Have another person help you keep track of the time to ensure full rinsing. 2. Familiarize yourself with the locations of fire extinguishers, eyewashes/drench hoses, and spill kits. 3. If you have long hair, you should tie it back. Hair burns easily. 4. Dress Code/Personal Protective Equipment (PPE):Wear long pants and closed shoes to lab at all times. Long sleeves are also recommended. If you do not have proper coverage, you will be asked to wear one of the lab aprons, lab coats, scrub pants, or a pair of shoe covers and may lose points for lab. 5. Avoid direct contact with all chemicals. Disposable gloves are provided on the counter by the windows. 6. If you spill a chemical, notify your instructor, who will advise you how to clean up the spill. 7. Do not bring food or drinks into the laboratory (this includes water bottles). There is a table outside the lab door for consumable items to be stored during lab time. 8. In case of fire to clothing, use the drench hose or shower. 9. In case of fire in containers or on the bench, step back, warn your instructor and others, and allow your instructor to take the appropriate action. There are fire extinguishers available. 10. Report all accidents, including all cuts and burns, promptly to your instructor. 11. In case of evacuation: as directed by your instructor or lab technician, turn off any gas, water and electrical equipment (hot plates, etc.) and proceed to the West end of the building to parking lot F. Stay together as a class and your instructor will join you to take attendance.

HANDLING OF REAGENTS AND APPARATUS 1. Chemicals will be found either: at the center of the benches, in the fume hoods, by the balances, or on the reagent bench. Please leave the chemical containers where you find them at the beginning of the lab so others can find them easily. 2. Read labels twice before taking any chemical to be sure it is both the correct substance and at the required concentration. 3. Never insert an instrument of your own into a chemical container. Use the droppers or scoops provided. 4. Never return chemicals to their original container. 5. When pouring liquids, hold the palm of your hand over the label. If some of the liquid runs down the outside of the bottle, wash it with water and dry it. 6. Always replace the lids on the chemical containers immediately after use. 7. Set stoppers and lids on the bench in such a way as to avoid contamination. 8. Be economical in the use of reagents, detergent, deionized water, and acetone. 9. Graduated cylinders, burettes and other glassware with fine and/or etched graduations should not be heated. Flasks and beakers with enameled graduations can be heated. 10. Be careful with hot objects. Hot objects should never be placed directly on the bench. Put them on your ceramic tile. 11. Some equipment, such as hot beaker tongs, 50 mL graduated cylinders, test tube racks, goggles, wash bottles, etc. is set out for common use by everyone. Do not place these items in your locker. 12. Treat the balances with special respect. Do not move the balance from its location. Never place a hot object on a balance. Carefully clean up anything you spill in, on, or around the balance.

CLEANLINESS 1. Wash your equipment as soon as you finish using it. Be sure to keep your both your locker equipment and common equipment clean, dry, and in good condition. 2. The cleaning procedure usually involves (1) cleaning first with tap water, detergent and a brush, (2) rinsing with tap water, (3) rinsing with deionized water, and (4) drying the outside with a paper towel. Do not use acetone unless directed to do so by your instructor. Consult your instructor if something won’t come clean. 3. It is your responsibility to immediately clean up any chemicals you spill in the lab. If you are not certain how to do this, consult your instructor. 4. Dispose of waste in the appropriate place. If you are unsure where to dispose of something, ask your instructor. a. WASTE BASKET: for paper and gloves only, no glass or chemicals. b. SINK: Only water-soluble, nontoxic substances. NO SOLIDS!!! Flush with tap water. c. SPECIAL DISPOSAL CONTAINERS: Your instructor will tell you which substances to place here. d. GLASS DISPOSAL CONTAINER: For all glass waste, including Pasteur pipets. No trash here! 5. Before leaving the laboratory, clean the sink nearest you and wash and dry your table top. Be sure that you have not left any of your own apparatus out. Please also see that you have not misappropriated any equipment that should have been returned [in good condition] to its original location.