2013 - 2014

Table of Contents

Topic 1: Quantitative Chemistry ......................................................................... 1

Topics 2 and 3: Atomic Theory and Periodicity ................................................ 37

Topic 4: Bonding ............................................................................................... 61

Topic 10: Organic Chemistry ............................................................................ 93

Option F: Food Chemistry .............................................................................. 117

Periodic Table ...................................................................................... Back Cover

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Quantitative Chemistry

1.1 – The Mole and Avogadro’s Constant 1. Apply the mole concept to substances

2. Determine the number of particles and the amount of substance (in moles)

1.2 – Formulas 1. Define the terms relative atomic mass (Ar) and relative molecular mass (Mr)

2. Calculate the mass of one mole of a species from its formula

3. Solve problems involving the relationship between the amount of substance in moles, mass and molar

mass

4. Distinguish between the terms empirical formula and molecular formula

5. Determine the empirical formula from the percent composition or from other experimental data

6. Determine molecular formula when given both the empirical formula and experimental data

1.3 – Chemical Equations 1. Deduce chemical equations when all reactants and products are given (balancing)

2. Identify the mole ratio of any two species in a chemical equation

3. Apply the state symbols (s), (l), (g), and (aq)

1.4 – Mass and Gaseous Volume Relationships in Chemical Reactions 1. Calculate theoretical yields from chemical equations

2. Determine the limiting reactant and the reactant in excess when quantities of reacting substances are

given

3. Solve problems involving theoretical, experimental and percentage yield

4. Apply Avogadro’s law to calculate reacting volumes of gases

5. Apply the concept of molar volume at standard temperature and pressure in calculations

6. Solve problems involving the relationship between temperature, pressure and volume for a fixed

mass of an ideal gas

7. Describe the characteristics of an ideal gas, and solve problems using the ideal gas equation, PV=nRT

8. Analyze graphs relating to the ideal gas equation

1.5 – Solutions 1. Distinguish between the terms solute, solvent, solution, and concentration (g dm-3, mol dm-3)

2. Solve problems involving concentration, amount solute and volume of solution

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3

PPREPARINGREPARING SSOAPOAP THROUGHTHROUGH THETHE SSAPONIFICATIONAPONIFICATION RREACTIONEACTION

Laboratory Goals

In this lab, you will: Learn how soap is prepared

Test some properties of soap

Safety Note

The potassium hydroxide solution used in this lab is extremely concentrated. Be sure to avoid any contact with skin and

especially eyes as it can cause serious burns. All spills MUST be immediately reported to the instructor and cleaned.

Introduction

The process of soap-making goes far back in history. Most people who have made soap throughout the centuries have

had no idea what is occurring; they simply made soap through trial and error, lots of luck and governing superstitions.

The process (similar to what we will be doing in lab) involved combining some form of fat with an alkali (basic) materi-

al. Most commonly the alkali was in the form of potash and pearlash, which contain KOH. Potash and pearlash soaps

were used by everyone from the reigning monarchs to the peasant or cottager, who made their own soap from the waste

fats and ashes they saved.

The First Soap

It is recorded that the Babylonians were making soap around 2800 B.C. and that it was known to the Phoenicians around

600 B.C. These early references to soap and soap-making were apropos the use of soap in the cleaning of textile fibers

such as wool and cotton in preparation for weaving into cloth.

The Romans and Celtics

The first definite and tangible proofs of soap-making are found in the history of ancient Rome. Pliny, the Roman histori-

an, described soap being made from goat's tallow and causticized wood ashes. He also wrote of common salt being add-

ed to make the soap hard. The ruins at Pompeii revealed a soap factory complete with finished bars.

While the Romans are well known for their public baths, generally soap was not used for personal hygiene. To clean the

body the Greeks (and later the Romans) would rub the body with olive oil and sand. A scraper, called a strigil, was then

used to scrape off the sand and olive oil, also removing dirt, grease, and dead cells from the skin, which was left clean.

Afterwards the skin was rubbed down with salves prepared from herbs.

Throughout history, people have taken baths in a variety of bathing mediums, with herbs and other ostensibly beneficial

additives. It is well known that Cleopatra, who captivated the leaders of the Roman world, attributed her beauty to her

baths in mare's milk. During the early centuries of the common era, soap was used by physicians in the treatment of dis-

ease. Galen, a second century physician, recommended bathing with soap for the amelioration of some skin conditions.

Soap for personal cleansing became popular during the later centuries of the Roman era.

The Celtic peoples are also thought by some historians to have discovered soap-making; their soaps were used for bath-

ing and washing. Perhaps due to increased contact with the Celtics by the Romans, using soap for personal cleasing be-

came popular.

There is an interesting legend surrounding the discovery of soap-making. This legend accords the discovery of soap to

the Romans, so it might have been fabricated to confront the Celtic claim to soap-making. Probably both of these in-

ventive peoples discovered soap-making independently. The legend asserts that soap was first discovered by women

washing clothes along the Tiber River at the bottom of Sapo Hill. The women noticed their wash became cleaner with

far less effort at that particular location. What was happening? The ashes and the grease of animals from the sacrificial

fires of the temples situated on the top of Sapo Hill mixed with the rain; the resulting soap--which ran down the slope in

the streams of rain water--give the women a wash day bonus. One can see at a glance that “saponification”, the chemical

name for the soap-making reaction, bears the name of that hill in Rome long ago, which caused one washer-women to

comment to another, "My wash is cleaner than yours".

The Chemistry of Soapmaking

As stated earlier, the chemistry behind soap-making was not understood for many years. It is now known that saponifica-

tion of soaps proceeds by the conversion of the triglycerides, which are the components of fats and oils, to fatty acid salts

and glycerol as show in Figure 1. The R groups in the figure represent long carbon chains with the accompanying hydro-

gens. For each specific triglyceride, these specific R groups can be determined. For example tristearin gives the reaction

shown in Figure 2.

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Typically, fats and oils have different more than one different R group in the same molecule so a variety of sodium salts

are produced. In order to separate out the salts from the rest of the reaction products, a saturated NaCl solution is added.

This forces the soap to coagulate without dissolving in the water. It can then be collected by filtration and washed to re-

move the excess base.

In this lab you will both make soap as well as test some of its properties. You can either use the provided fats and oils, or

you can bring in your own sample to make a different soap from everyone else in lab.

Prelab

Here are some terms we may use in our discussions during the activity. We may diagram and discuss these terms togeth-

er before going into the lab. Some space has been left if you’d like to add diagrams. Add these to your working chemis-

try vocabulary.

triglyceride: Triglycerides are the main type of fat transported by your body. The fat gets its name from the chemical

structure: triesters of glycerol and three carboxylic acids. Triglycerides are not water soluble. Fats are solid triglycerides,

and oils are liquid triglycerides. We will discuss these more during our unit on organic chemistry.

fatty acid: A fatty acid is the carboxylic acid portion of the triglyceride. They range in size from 10 – 20 carbon atoms,

and usually have an even number of carbons.

A fatty acid is said to be saturated if it does not contain any alkene groups (carbon-carbon double bonds), and unsaturat-

ed if it does contain an alkene group.

hydrophobic: Hydrophobic means “water-hating,” meaning it will not mix with water. The long carbon chain portion

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of the fatty acid is hydrophobic.

hydrophilic: Hydrophilic means “water-loving,” meaning it will mix with water. The carboxylic acid portion of the fat-

ty acid is hydrophilic.

mixture: A mixture is an impure substance in which the components can be separated from one another by physical

means. If the mixture is uniform throughout, it is said to be homogeneous, whereas heterogeneous mixtures are not uni-

form throughout. A chemical solution is an example of a homogeneous mixture.

emulsion: An emulsion is created when small globules of one liquid are suspended throughout a second liquid in which

the first will not mix to form a homogeneous solution. Emulsions have a milky white appearance due to light being

reflected off the small, suspended globules of oil.

Many salad dressings are examples of emulsions, in which globules of hydrophobic oil are suspended throughout hydro-

philic vinegar. Some salad dressings will “separate out” if not shaken consistently. Mayonnaise is also an emulsion in

which egg is mixed with oil.

Emulsions are a subset of a larger class of two-phase mixtures called colloids.

micelle: A micelle is a spherical arrangement of amphipathic molecules (molecules containing both a hydrophobic and

hydrophilic portion) in which the hydrophilic ends are arranged outwardly (contacting the water) and the hydrophobic

ends are arranged inwardly.

Soap molecules form micelles in water when the hydrophilic head of the molecules face outward (exposed to the water)

and the hydrophobic tails face inward, towards each other. Dirt, which is a mixture of grease, oils, and other hydropho-

bic substances, can easily associate with the hydrophobic “insides” of the soap micelles. Water is then able to carry the

“dirty” micelles away by dissolving the hydrophilic outside of the micelles.

Answer the following questions. They will help you prepare for the lab.

1. What reaction is described by the word saponification? What are the products from this reaction and what are the

required reactants?

2. What is the most dangerous part of the lab? Why is it critical to clean up all spills immediately?

Procedure

1. First you will have to select the type of fat or oil that you wish to use to make soap. Olive oil, vegetable oil, and

lard will be provided. You are welcome to use any other type of oil or fat, but you will have to provide it

(looking at other oils can give you drastically different soaps that provide an interesting comparison.)

2. Start with 20 mL of the selected oil (or about 16 g of the lard) and put it in a clean 400 mL beaker.

3. Add in 20 mL of ethanol followed by 25 mL of 20 % potassium hydroxide (remember this is very concentrated.

If any of it is spilled it should be reported to immediately to the instructor and then cleaned up. It is concentrated

enough to cause serious chemical burns and irreversible eye damage if it comes in contact with you.) Stir the

mixture with a glass rod.

4. Turn on a hot plate to level 5 and place the beaker atop it. Periodically stir the mixture during heating. During the

heating the mixture may foam up. Stirring will help prevent this, but if the foam climbs up the beaker you will

need to remove it from the heat (using beaker tongs) momentarily until the foaming subsides.

5. The mixture should be heated until all of the ethanol is removed. When is that? It will be when ethanol vapors

are no longer being released from the heated mixture. You can detect this by smell, as the ethanol vapor has a

characteristic odor. When smelling, be sure and use the wafting technique, rather than placing your nose directly

over the boiling mixture. The loss of the ethanol will likely coincide with the increase in the amount of foaming

that is occurring in your beaker.

6. Once the ethanol is gone remove the beaker from the heat and turn off the hot plate.

7. After allowing the mixture to cool most of the way to room temperature, add in 100 mL of saturated sodium

chloride and mix thoroughly. The soap should coagulate into a solid mass and can now be filtered to remove the

by-products. Our filtering will be done using cheesecloth. To do this put the cheesecloth on top of a large beaker

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and pour the soap containing liquid onto the gauze (some soap may get through, but the majority will collect on

the gauze.) The soap should then be washed at least twice with 10 mL of ice water (ice can be obtained from the

biology lab) to remove the excess KOH. It may be more effective to put the soap back into the original beaker

and add the ice water and filter again. Once the rinsed soap has drip dried, move it to a paper towel to finish dry-

ing.

Testing the Soap

1. Take a pea sized piece of the newly formed soap and put it in a 125 mL Erlenmeyer flask and add about 50 mL

of distilled water. Repeat this in a second and third flask using either a commercial soap or soap made by another

lab group starting from a different fat or oil. Stopper the flasks and shake them vigorously for 20-30 seconds and

observe the results paying attention to the solubility of the soaps and the foaming action. The ability to create a

foam indicates the presence of the soap. The amount of foaming and the length of time until deflation both relate

to the surface tension of the solution.

2. Once you have been able to compare the rate of deflation of the foam, take a clean glass rod and dip it into a so-

lution and then touch a piece of pH paper. Test the other solutions as well as the pH of distilled water. This will

give a feeling of how well you were able to remove the excess NaOH.

3. Next add two drops of mineral oil to the solutions. To a separate, clean flask add 50 mL of distilled water and 2

drops of mineral oil. Cover and shake the flasks for about 10 seconds and compare how well the soaps were able

to emulsify the oil (prevent it from immediately separating out.)

Chemical Disposal

All liquid chemicals may be disposed of down the drain. KOH should be washed down the drain with lots of

water. Any solid waste (excess lard or soap that you don’t want to keep) should be discarded in the trash.

Data Collection: Use this space to neatly record data collected during the lab.

Questions

1. How is a pure substance different than an impure substance? What are some examples of pure substances and

impure substances used in this lab?

2. How is a physical change different than a chemical change? At what point during the lab was there evidence that

a chemical change was taking place?

3. What is an emulsion? Which mixture was the better at maintaining an emulsion, oil and water or soap, oil and

water? Explain why on a molecular level.

4. How can soap remove oil or dirt from clothes?

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NNAMINGAMING IIONICONIC CCOMPOUNDSOMPOUNDS

What are the structural units that make up ionic compounds and how are they named?

Why?

When working in chemistry, it is often convenient to write a chemical in symbols. For example we might write down the substance table salt as NaCl. In talking about chemistry however, it is a bit tacky to say “en-ay see-ell” when we want to refer to a substance. Also, in formal writing we should use the name of the compound rather than its symbols. Therefore we need to learn how to say the proper names of ionic substances.

Model 1 – Ion Charges for Selected Elements

1. Based on the information in Model 1:

a. Identify three elements that form only one cation.

b. Identify three elements that form only one anion.

c. Identify three elements that form more than one cation.

d. In what region of the periodic table are these “multiple ion” elements usually located?

2. Consider the ions of potassium (K) and sulfur (S). Write chemical formulas for all possible ionic compounds in-volving these ions, using the simplest ratio(s) of potassium (K) and sulfur (S). Keep in mind that the sum of the charges in an ionic compound must equal zero.

3. Consider the ions of iron (Fe) and sulfur (S). Write chemical formulas for all possible ionic compounds involving these ions, using the simplest ratio(s) of iron (Fe) and sulfur (S). Keep in mind that the sum of the charges in an ionic compound must equal zero.

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Model 2 – Ionic Compound Names (Metals that form one ion)

4. Circle the symbol for the metal in each of the compounds in Model 2.

5. Which element comes first in the name and formula of the compounds in Model 2-the metal or the nonmetal?

6. Use the table of ions in Model 1 to answer the following questions: a. In the compound zinc phosphide, what is the charge on the zinc ion?

b. In the compound zinc phosphide, what is the charge on the phosphide ion?

7. Explain why a 3 to 2 ratio of ions is necessary for the compound zinc phosphide.

8. The compound carbon dioxide has a name that gives you a hint as to how many oxygen atoms are in the com-pound. Is there anything in the name “zinc phosphide” that indicates there are three zinc and two phosphorus ions in the formula unit?

9. Is there any other ratio of zinc and phosphorus ions that could exist? For instance, could you have Zn2P or

ZnP2? Explain your answer. 10. Explain why you don’t need to specify the number of ions in the compound when you are naming ionic sub-

stances like those in Model 2. 11. Model 2 is labeled “Metals that form one ion.” What other metals that also form only one ion could be includ-

ed in the Model 2 list? Model 1 may be helpful in this regard. 12. Describe how the names of the nonmetal elements in Model 2 are changed when they are in their anion

forms.

NaCl Sodium chloride Zn3P2 Zinc phosphide

CaS Calcium sulfide Al2O3 Aluminum oxide

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14. Provide the chemical formula for each of the following ionic compounds. Barium chloride Magnesium oxide 15. Consider the two chemical formulas you wrote in Question 3 for compounds of iron and sulfur. Would the

name “iron sulfide” be sufficient to uniquely identify either of those compounds? Explain.

Read This!

When the metal in an ionic compound always forms an ion with the same charge, you need not indicate that charge as part of the compound name. However, some atoms have the ability to form more than one type of ion. This can make naming confusing. You can’t simply refer to a compound of copper and oxygen as “copper oxide.” People won’t know which compound you are referring to—CuO or Cu2O.

Model 3 – Ionic Compound Names (Metals that form multiple ions) 16. Model 3 is labeled “Metals that form multiple ions.” What other metals that form multiple ions could be in-

cluded in Model 3? Model 1 may be helpful in this regard. 17. Describe the most obvious difference between the names in Model 3 and those in Model 2. 18. Do the Roman numerals in the names in Model 3 relate to the number of cations or number of anions in the

formula unit? Support your answer by citing two specific examples. 19. Keeping in mind that the sum of the charges in an ionic compound must equal zero, use the chemical formulas

in Model 3 to answer the following questions:

a. Identify the charge on the copper cations in copper(I) oxide and copper(II) oxide, respectively. b. Identify the charge on the iron cations in iron(II) chloride and iron(III) chloride, respectively. 20. What do the Roman numerals in the compounds described in Question 19 indicate?

Cu2O Copper(I) oxide PbO Lead(II) oxide

CuO Copper(II) oxide PbO2 Lead(IV) oxide

SnF2 Tin(II) fluoride FeCl2 Iron(II) chloride

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21. Fill in the table below using what you’ve learned from Model 3.

22. For each of the compounds in the table below, determine the type of metal in the compound and then name the compound using the correct naming method.

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TTOPICOPIC 1 P1 PROBLEMROBLEM SSETET –– QQUANTITATIVEUANTITATIVE CCHEMISTRYHEMISTRY

Many of these questions involve concepts you should be familiar with from MYP chemistry. They progress from simple to more challenging. After completing 1 – 58, a selection of prior exam questions follows. It is important that you devel-op a feel for what prior exam questions are like, as being familiar with exam questions will be quite helpful when you take your IB chemistry exam.

1. Practice your fluency in inorganic nomenclature with these charts.

Ionic Compounds:

Molecular (Covalent) Compounds:

2. Give the name for the following ions:

3.

formula # and type of atoms and Ar of each atom Formula mass (Mr)

cations anions formula name Formula mass (Mr)

Al3+

S2-

Potassium oxide

Ca2+

NO3-

Zinc (II) carbonate

CuF2

Name Formula Formula mass (Mr)

SO3

P2O5

Carbon tetrachloride

Trinitrogen hexasulfide

NO2

a) NH4+ e) OH- i) PO4

3-

b) SO42- f) Cr2O7

2- j) CH3COO-

c) CrO42- g) CO3

2-

d) NO3- h) HCO3

-

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3. How many moles are there in 4.2 g aluminum chloride? 4. How many moles are there in 42.5 g copper (II) fluoride? 5. What is the mass of 2.25 moles zinc (II) phosphate? 6. How many moles are there in 27.5 g aluminum nitrate? 7. How many molecules of sulfur dioxide are there in 1.25 g of sulfur dioxide? 8. What is the mass of 3.29 x 10

25 molecules of water?

9. How many moles of hydrogen are there in 64.9 g water? 10. What mass of zinc is in 14.4 g zinc (II) fluoride? 11. How many atoms of bromine are there in 0.155 g magnesium bromide? 12. How many atoms are there in 12 g of K2O? 13. Ionic compounds form solid crystals that are held together by the attractive force between positive cations

and negative anions. A hydrated ionic compound contains small amounts of water distributed throughout the crystal. A hydrated ionic compound has a unique formula, in which the number of moles of water is de-noted after the formula of the compound. For example, in calcium chloride dihydrate, there are two moles of

water to every one mole of calcium chloride, and the formula is written: CaCl2 ∙ 2H2O.

A hydrated copper (II) sulfate crystal is dehydrated by heating above a flame. Hydrated copper (II) sulfate has a characteristic light blue color, while dehydrated copper (II) sulfate is gray/white.

Consider the data

Calculate the formula of the hydrate.

Object Mass (± 0.1 g)

evaporating dish 47.4 g

evaporating dish and CuSO4 * X H2O 52.6 g

mass of dish and CuSO4 remaining after heating 50.9 g

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For each of the following (14 – 22):

a) classify the reaction b) write chemical equation (include state symbols: s, l, g, aq) c) balance the equation

14.liquid ethyl alcohol ( CH3CH2OH) burns in the presence of oxygen

15.sodium reacts with aqueous aluminum carbonate

16.zinc reacts with oxygen

17.aqueous sodium sulfide reacts with aqueous barium nitrate

18.copper (II) carbonate reacts to form copper (II) oxide and a gas

19.potassium reacts with iodine

20.ethane (C2H6) burns in air

21.barium reacts with aqueous zinc (II) chloride

22.aqueous sodium carbonate reacts with aqueous zinc (II) fluoride

23.lead (II) oxide decomposes

For each of the following (24 – 27), determine the empirical formula from the percentage composition or mass data giv-en. An empirical formula displays the lowest whole number mole ratio of elements in a compound. A molecular formula shows this same ratio, only it isn’t always in the most simplified form (for example, glucose has a molecu-lar formula of C6H12O6, but when simplified, its empirical formula is CH2O).

24. A 3.00 gram sample of an unknown compound that contains 1.71 g carbon and 2.29 g oxygen

25. A 7.95 gram sample of an unknown compound that contains 2.03 g magnesium and 5.92 g chlorine

26. An unknown sample that contains 31.9 % potassium, 29.0 % chlorine, and 39.1 % oxygen, by mass.

27. Deduce the empirical formula of a compound that contains 25.9 % nitrogen and 74.1% oxygen, by mass.

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28. A compound is 80% carbon and 20% hydrogen, by mass. a) What is the empirical formula?

b) If the molar mass is about 30, what is the molecular formula?

29. Determine the empirical and molecular formulas of a compound with a molecular weight of 42 that contains 85.9% C

and 14.1% H.

30. A compound is analyzed and found to contain: 4.80g carbon, 0.80g hydrogen, 3.20g oxygen and 2.80g nitrogen. What is the empirical formula? If the molar mass is about 174, what is the molecular formula?

31. A sample of a hydrocarbon burns completely in oxygen to form 13.2g carbon dioxide and 5.4g water. What is the

empirical formula? For 32 - 37:

Write correct formulas for products and reactants.

Write the equation

Balance the equation

Perform the conversion, using the following T-chart template as a guide if you like:

mass given mol given unknown from equation formula mass unknown

Formula mass given given from equation mol unknown

32. How many grams of oxygen react with excess hydrogen to form 13.5 g water?

33. How many grams of silver are formed as 125 g zinc reacts with excess aqueous silver nitrate?

34. What mass of aqueous lithium carbonate reacts with excess calcium nitrate to form 87 g of aqueous calcium car-bonate?

35. How many grams of oxygen form as 50.0 g of Hg2O decomposes?

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36. How many grams of potassium oxide form as 15.0g potassium burns in the presence of oxygen?

37. Calculate the amount (in grams) of potassium and the amount (in grams) of silver (I) oxide required to produce 4.6 grams of silver.

38. 45.0 g of potassium carbonate react with 45.0 g silver nitrate. The products of the reaction are silver carbonate and potassium nitrate.

a) What is the theoretical yield of silver (I) carbonate (i.e. what is the maximum mass of silver (I) carbonate that can be produced)?

b) What is the limiting reactant? c) What is the reactant present in excess? d) Determine the percentage yield if only 30.0 grams of silver (I) carbonate are produced experimentally.

39. 15.0 g of potassium reacts with 5.0 g sulfur. a) What is the theoretical yield of potassium sulfide (i.e. what is the maximum mass of potassium sulfide

that can be produced)?

b) What is the limiting reactant? c) What is the reactant present in excess?

d) Determine the percentage yield if only 10.9 grams of potassium sulfide are produced experimentally.

16

40. What is the maximum mass of iron (III) oxide produced when 10.0 grams of iron are allowed to completely oxidize?

4Fe(s) + 3O2(g) 2Fe2O3(s)

41. What volume of carbon dioxide is released when 10 dm

3 of carbon monoxide are combined with 10 dm

3 of oxygen?

2CO(g) + O2(g) 2CO2(g)

42. 1.00 moles of NaCl are dissolved in 500 ml of solution. What is the concentration in mol dm

-3?

43. 3.5 moles of KF are dissolved in 2.0 L of solution. What is the concentration in mol dm-3

?

44. 133.0 g of NaCl are dissolved in 1.00 L of solution. What is the concentration in mol dm-3

?

45. How many moles of sodium carbonate are dissolved in 500 ml of 0.055 mol dm-3

solution?

46. What is the concentration in mol dm-3

of 38 g lead (II) iodide dissolved in 500 ml of solution?

47. How many grams of NaCl are there in 500 ml of a 1.3 mol dm-3

solution NaCl?

48. What mass of copper (II) chloride will react with 500 mL of 2 mol dm-3

magnesium chloride?

49. Methanol, CH3OH, can be prepared according the equation: CO + 2 H2 CH3OH

When 175 g of carbon monoxide is mixed with 36.5 g hydrogen gas and allowed to react, 112 g of product forms. Determine a) theoretical yield b) percent yield c) limiting reactant

17

50. 150 mL of 1.5 mol dm-3

BaCl2 and 30.0g Na2SO4 are placed in a beaker to react.

a) How many grams of BaSO4 form?

b) Which reactant is the limiting reactant?

c) Which reactant is present in excess?

d) If 20 g of BaSO4 is isolated, what is the percent yield?

51. 12 mL of 3.5 mol dm-3

calcium nitrate and 27.5 ml of 1.2 mol dm-3

aluminum sulfate are placed in a beaker to react.

a) What is the theoretical yield of CaSO4?

b) Which reactant is the limiting reactant?

c) Which reactant is present in excess? How much will remain unreacted?

52. A mixture of MgSO4 and MgSO4 ∙ 7 H2O with a mass of 11.08 g is heated until the water is removed. The total mass

of the anhydrous MgSO4 after heating is 7.25g. Determine the percentage of the hydrate in the original mixture.

53. If 12.56 g of a solid dibasic acid reacts completely with 24.4 cm

3 of 1.8 mol dm

-3 NaOH, what is the molar mass of

the acid, in g/mol?

54. Menthol is a compound that contains only carbon, hydrogen and oxygen. When a 0.0956g sample of menthol burns in air, 0.269 g CO2 and 0.110 g H2O are formed. What is the empirical formula of menthol?

18

55. 7.321 mg of an organic compound containing carbon, hydrogen, and oxygen was analyzed by combustion. The

amount of carbon dioxide produced was 17.873 mg and the amount of water produced was 7.316 mg. Determine

the empirical formula of the compound.

56. 0.1101 gram of an organic compound containing carbon, hydrogen, and oxygen was analyzed by combustion. The

amount of carbon dioxide produced was 0.2503 gram and the amount of water produced was 0.1025 gram. A de-

termination of the molar mass of the compound indicated a value of approximately 115 grams/mol. Determine the

empirical formula and the molecular formula of the compound.

19

UUSINGSING CHEMICALCHEMICAL PRECIPITATESPRECIPITATES TOTO NAMENAME IONICIONIC COMPOUNDSCOMPOUNDS In the context of chemistry, precipitation refers to the formation of a solid phase product from an aqueous solution. An

aqueous solution is created when a water soluble compound is dissolved in water. Take a solution of salt water, for

example. When crystals of sodium chloride are dissolved in water to form salt water, the solution of salt water can tech-

nically be called an aqueous solution. All aqueous solutions are homogeneous mixtures, because they are composed of

more than one type of pure substance and have the same composition throughout.

In this investigation, we will be forming precipitates of ionic compounds, which are also called salts. Sodium chloride

(common table salt) is obviously a “salt,” but so is any other compound formed from two or more oppositely charged

ions. A salt is created when positively charged cations stick to negatively charged anions. Magnesium oxide is a salt

(magnesium cations, Mg2+

, and oxide anions, O2-

). Copper (II) sulfate is a salt (copper (II) cations, Cu2+

, and sulfate ani-

ons, SO42-

). Lithium chloride is a salt (lithium cations, Li+, and chloride anions, Cl

-). We could go on for hours. ALL ION-

IC COMPOUNDS ARE SALTS! Positive cations (usually metal ions), and negative anions (usually nonmetals, or polya-

tomic ions), held together by electrostatic attraction – that makes a salt!

The electrostatic attraction between the cations and anions in a salt is relatively high; it isn’t real easy to break them

apart. It can be done however, even by ordinary water, for some salts. Water has the ability to separate the oppositely

charged ions in some salts, but not others. Those salts that can be “dissolved” (literally, broken up) by water are water

soluble. Those that cannot are water insoluble. There are several factors that dictate whether a salt is water soluble or

not, but we won’t go into that just yet. It’s actually pretty fascinating though!

Let’s take common table salt, or sodium chloride for example (sodium cations, Na

+, and chloride anions, Cl

-):

We won’t get too involved with what happened to create the posi-

tive sodium ions and negative chloride ions, but what basically

happened was that a neutral atom of sodium, with 11 protons

and 11 electrons (Na = atomic number 11), gave up an electron

to a neutral atom of chlorine that had 17 protons and 17 elec-

trons (Cl = atomic number 17). This left the sodium atom with 11

protons and 10 electrons, and since protons are positive and

electrons are negative, the sodium atom acquired a charge of +1.

This made it a cation! The chlorine atom was left with 17 protons

and 18 electrons. This made it a chloride anion with a -1 charge!

Since exchanging an electron made two oppositely charged ions,

the oppositely charged ions will stick together, forming an ionic

compound (a salt!). However, this chemical bond, called an ion-

ic bond, to be specific, is not irreversible. It can be broken by

ordinary water.

We can write a chemical equation for this process:

H2O(l) NaCl (s) Na

+(aq) + Cl

-(aq)

In the equation, the (s) subscripts refer to the solid state of mat-

ter, and the (aq) subscripts refer to the aqueous state (meaning dissolved in water). The (l) subscript means liquid.

(g) subscripts would indicate the gaseous state of matter, although no gases are present in this simple process. Ba-

sically, when we add water to solid sodium chloride, we make an aqueous solution of sodium chloride.

20

The equation says that solid sodium chloride is converted into sodium ions and chloride ions by liquid water. Since we

had sodium ions and chloride ions both before and after the process occurred, this is just a physical change; the chemi-

cal makeup hasn’t changed. Before the process occurred, the sodium and chloride ions were stuck together. After the

process occurred, they weren’t. Physical change. Not chemical.

Some salts are NOT water soluble. For example, silver (I) chloride:

H2O(l) AgCl (s) AgCl (s)

Even after water was added to solid silver (I) chloride, the ions remained stuck together. No change took place, not even

the physical separation of the ions. Silver ions, Ag+, and chloride ions, Cl

-, are just too darn attracted to each other.

Kind of like Romeo and Juliet. Nothing could keep them apart.

So what would happen if we mixed a solution of sodium ions and chloride ions (an aqueous solution of sodium chloride)

with a solution of silver ions? Chloride ions may not be too interested in sticking to sodium ions, but when silver ions are

there, well, hey, whole different story!

Ag+

(aq) + Cl-(aq) AgCl(s)

Silver ions and chloride ions stick together so well that water cannot separate them. Silver (I) chloride is not water solu-

ble. Some silver salts are water soluble, like silver (I) nitrate for example. But many silver salts are not soluble.

So, if we added water to two different salts that are water soluble, we would have two aqueous solutions. Let’s say one

was sodium chloride, the other was silver (I) nitrate:

H2O(l) Solution 1: NaCl (s) Na

+(aq) + Cl

-(aq)

H2O(l) Solution 2: AgNO3 (s) Ag

+(aq) + NO3

-(aq)

One solution has sodium and chloride ions, who aren’t all that keen to stick back together. The other solution has silver

and nitrate ions, and the same is true with them. But . . . . . . . what if the solutions were mixed??! Remember how well

silver and chloride ions stick together?

Ag+

(aq) + Cl-(aq) AgCl(s)

This is what happens when these two aqueous solutions are mixed. A precipitate of silver chloride, a solid product from

two solutions, is formed. In this lab, you will create several precipitates and use your own knowledge to predict their

names. In the example above, the ions combined in a 1:1 ratio. This is not the case for all cation/anion combinations!

Generally, this chart can be used to determine whether or not an ionic compound will be water soluble:

21

Method: There were be several aqueous solutions of ions available:

*******All solutions are to be treated as potentially hazardous – under no circumstances should you allow the solutions to come into contact with your skin, eyes, or mucous membranes!************

1. Predict: There are 16 possible cation/anion combinations. Which cation/anion combinations will form precipitates when the solutions are combined?

2. Start by obtaining a well plate from the lab counter in the back. Be sure it is washed and dried thoroughly before beginning. Be sure it has at least an 4 x 4 section of wells.

3. Take note of the color and appearance of all of the ion solutions before being

combined with other solutions. 4. While carefully observing for signs of precipitate formation, add a few drops of

silver cations (Ag+) to four wells, all in one column. Add iron ions to four wells in another column, copper ions to another column, and finally potassium ions in another column:

5. Add a few drops of hydroxide ions (OH-) to each of the four cations, Ag

+, Fe

3+, Cu

2+, and K

+. In another row,

add a few drops of carbonate ions (CO32-) to each of the four cations. Repeat for the final two anions (phosphate and chloride) so that you have combined each of the positive cations with each of the anions.

6. It is safe to dispose of these chemicals down the drain with plenty of water. Please take care not to expose

your skin to the solutions during rinsing/cleaning your well plate.

Observations: Appearance of ion solutions before being combined: Appearance of ion solutions after being combined:

Ag+

PO43-

Cu2+

OH-

K+

Cl-

CO32-

Fe3+

22

While some solutions combinations may not have formed precipitates, others will have formed visible precipitates (did

your predictions match the results?). Write the chemical formula, as well as the formula name, for all precipitates

formed during the laboratory below:

23

FFORMULAORMULA FFOROR A HA HYDRATEYDRATE Problem:

What is the percentage of water in a hydrated salt? What is the formula for the hydrate. Introduction:

A hydrate is a compound that contains water in its crystal structure. The water may be removed from the salt in the la-boratory by heating the salt. The salt without the water is called an anhydrous salt. Here are some examples of hy-drates: CaSO4 · 2H2O CoCl2 · 6 H2O Mg SO4 · 7 H2O In a hydrate the water molecules are a distinct part of the compound but are joined to the salt by connections that are weaker than the connections in the salt or the connections in the water molecules. Notice we use a dot to connect the water units to the salt formula. In this experiment you will measure the mass of a hydrate. Then remove the water from the crystals and measure the mass of the anhydrous salt. The data gathered will allow you to determine the percent water in the hydrated salt and also the empirical formula for the hydrated salt. Pre-Laboratory Assignment:

(Work with your lab partner(s) on these questions before beginning). Read the Problem and Introduction before you begin.

1. What information is necessary to determine the percentage of water in your hydrate sample?

2. How will the water be removed from the hydrate in this experiment?

3. A hydrate has the formula Mg SO4 · 7 H2O. What is the percent water in this hydrate. Method:

Write a brief method for dehydrating the copper(II) sulfate hydrate you are given. Observations and Data:

mass of empty beaker ______g mass of beaker + hydrate ______g mass of beaker + anhydrous salt ______g

24

Analysis and Conclusions: 1. Determine the percentage of water in the hydrated salt. Think about the fact that the difference between the mass of

the hydrated salt and the mass of the anhydrous salt is the mass of the water in the hydrate. Remember the definition of percent.

2. Obtain the formula for your anhydrous salt and write it here: ___________________ 3. Determine the number of moles of water evaporated from your hydrate. Think about how mass is related to moles. 4. Determine the number of moles of anhydrous salt. 5. Calculate the empirical formula for your hydrated salt. The empirical formula for a hydrate is always written in the

form of 1 mole anhydrous salt · moles H2O. The 1 mole in front of the anhydrous salt formula is understood and is not written as part of the formula.

6. Why should the beaker and contents be cooled before finding its mass? 7. Why must at least 2 successive mass readings be equal before finishing the experiment? 8. List several possible sources of error for this experiment.

25

AANN EXPERIMENTEXPERIMENT TOTO DETERMINEDETERMINE THETHE IDENTITYIDENTITY OFOF ANAN UNKNOWNUNKNOWN GASGAS BYBY

CALCULATINGCALCULATING ITSITS MOLARMOLAR MASSMASS

Research question:

What is the molar mass, and thus the probable identity, of the gas dispensed in the chemistry laboratory?

Hypothesis:

The gas dispensed in the laboratory is methane, and thus the calculated molar mass of the gas should be approximately 16 grams per mole.

Materials:

500 cm3 Erlenmeyer flask

No. 7 rubber stopper

40 cm hollow rubber tubing

Digital balance (± 0.01 g)

Unknown gas from the installed nozzle on lab bench

Digital thermometer (± 0.1 °C)

Mercury Barometer (± 0.1 cm Hg)

100 cm3 graduated cylinder (± 1 cm

3)

Method:

1. Find and record the mass of the empty (air-filled) 500 cm3 flask and the No. 7 rubber stopper.

2. Attach the hollow rubber tubing to the gas nozzle on the lab bench. Feed the free end of the tubing into the flask

and carefully open the valve to allow the gas to flow into the flask. Allow the gas to flow for approximately five

seconds.

3. Close the valve and quickly stopper the flask. Find and record the mass of the flask with the unknown gas and

the stopper.

4. Repeat steps 2 and 3 until the mass of the flask no longer changes.

5. To find the volume of the flask, completely fill the flask with water and replace the stopper (which will displace

some of the water). Remove the stopper and pour the water into a graduated cylinder to determine the volume.

6. Record the ambient temperature and pressure in the room using a thermometer and a barometer.

Data Collection:

Air pressure: __________ cm Hg (± 0.1 cm Hg)

Temperature: __________ °C (± 0.1 °C)

Volume of water in flask: _______ cm3 (± 1 cm

3) + _______ cm

3 (± 1 cm

3) + _______ cm

3 (± 1 cm

3)

+ _______ cm3 (± 1 cm

3) + _______ cm

3 (± 1 cm

3)

Object Mass (± 0.01g)

Flask/Stopper/Air

Flask/Stopper/Gas (1)

Flask/Stopper/Gas (2)

Flask/Stopper/Gas (3)

Flask/Stopper/Gas (4)

26

Data Processing:

1. Determine the percent uncertainty for the measured pressure ([absolute uncertainty/measurement] x 100).

Convert the pressure from the measured unit (cm Hg) to kPa. The percent uncertainty will be the same for

both units.

2. Determine the percent uncertainty for the measured temperature. Convert the temperature from the meas-

ured units (°C) to Kelvins. The percent uncertainty will be the same for both units.

3. Add up each separate volume measurement to find the total volume of the flask. Add the absolute uncertain-

ties for each measurement to determine the overall absolute uncertainty for the volume measurement. Use

this absolute uncertainty to determine the percent uncertainty for the volume. Convert the volume to dm3.

The percent uncertainty will be the same for both units.

4. Determine the moles of gas that occupy the flask. This will be the same for both air and the unknown gas.

Add the percent uncertainties for the measured pressure, volume, and temperature. This is the percent un-

certainty for the number of moles of gas.

5. Convert the moles of air in the flask to grams of air using the molecular mass for air (28.96 g/mole). Subtract

the mass of air from the mass of the flask/stopper/air to determine the mass of the evacuated flask/stopper.

Add the absolute uncertainties for these two raw measurements; this is the absolute uncertainty for the mass

of the evacuated flask/stopper.

6. Subtract the mass of the evacuated flask/stopper from the final mass of the flask/stopper/gas to determine the

mass of the unknown gas. Add the absolute uncertainties for these two measurements; this is the absolute

uncertainty for the mass of the unknown gas. Determine the percent uncertainty for the mass of the unknown

gas.

7. Determine the molar mass of the gas by dividing the mass of the unknown gas by the moles of the unknown

gas. Add the percent uncertainty for these measurements; this is the percent uncertainty for the molar mass

of the unknown gas.

8. The true identity of the gas is methane (CH4), with a molar mass of 16.05 g/mole. Determine the percent error

for your measurement:

| (experimentally-determined molar mass – true molar mass) / true molar mass | x 100

Evaluation/Conclusion:

Report the final, calculated molar mass for the unknown gas, including the percent uncertainty. Compare your

calculated value to the expected value, reporting your percent error. State if the percent error of your measure-

ment lies within the percent uncertainty. If your measured value lies within the margin of percent uncertainty, it

can be stated that the amount of systematic error in this experiment (method-based error) was minimal, and that

the majority of error was from random uncertainty (measurement-based error).

Identify potential weaknesses or error sources in the experimental method (i.e. is the gas pure methane or are

there other components?). Suggest improvements to these weaknesses that would improve the accuracy of

your measurement if you were to repeat the experiment.

27

PPASTAST IB TIB TESTEST QQUESTIONSUESTIONS FORFOR TTOPICOPIC 11

1. What amount of oxygen, O2, (in moles) contains 1.8 ×1022

molecules?

A. 0.0030 C. 0.30

B. 0.030 D. 3.0

2. Which compound has the empirical formula with the greatest mass?

A. C2H6 C. C5H10

B. C4H10 D. C6H6

3. __C2H2(g) + __O2(g) → __ CO2(g) + __ H2O(g)

When the equation above is balanced, what is the coefficient for oxygen?

A. 2 C. 4

B. 3 D. 5

4. 3.0 dm

3 of sulfur dioxide is reacted with 2.0 dm

3 of oxygen according to the equation below.

2SO2(g) +O2(g) → 2SO3(g)

What volume of sulfur trioxide (in dm3)is formed? (Assume the reaction goes to completion and all gases are

measured at the same temperature and pressure.)

A. 5.0 C. 3.0

B. 4.0 D. 2.0

5. The relative molecular mass of aluminium chloride is 267 and its composition by mass is 20.3% Al and 79.7% chlorine. Determine the empirical and molecular formulas of aluminium chloride.

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

6. Sodium reacts with water as follows.

2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)

1.15 g of sodium is allowed to react completely with water. The resulting solution is diluted to 250 cm3. Cal-

culate the concentration, in mol dm–3

, of the resulting sodium hydroxide solution.

…………………………………………………………………………………………………

…………………………………………………………………………………………………

…………………………………………………………………………………………………

…………………………………………………………………………………………………

28

7. (i) Calcium carbonate is added to separate solutions of hydrochloric acid and ethanoic acid of the same concentration. State one similarity and one difference in the observations you could make.

(ii) Write an equation for the reaction between hydrochloric acid and calcium carbonate.

…………………………………………………………………………………………..

(iii) Determine the volume of 1.50 mol dm–3

of hydrochloric acid that would react with exactly 1.25 g of calcium carbonate.

…………………………………………………………………………………………..

…………………………………………………………………………………………..

…………………………………………………………………………………………..

…………………………………………………………………………………………..

(iv) Calculate the volume of carbon dioxide, measured at 273 K and 1.01 × 10

5 Pa, which would be produced when 1.25 g of calcium carbonate reacts completely with the

hydrochloric acid.

…………………………………………………………………………………………..

…………………………………………………………………………………………..

…………………………………………………………………………………………..

…………………………………………………………………………………………..

…………………………………………………………………………………………..

8. What volume (in dm3) of 0.30 mol dm

–3 NaCl solution can be prepared from 0.060 mol of solute?

A. 0.018 C. 0.50

B. 0.20 D. 5.0

9. What amount (in moles) is present in 2.0 g of sodium hydroxide, NaOH?

A. 0.050 C. 20

B. 0.10 D. 8

10. A hydrocarbon contains 90 % by mass of carbon. What is its empirical formula?

A. CH2 C. C7H10

B. C3H4 D. C9H10

11. Copper can react with nitric acid as follows.

3Cu +_HNO3 → _Cu(NO3)2 +_H2O + _NO

What is the coefficient for HNO3 when the equation is balanced?

A. 4 C. 8

B. 6 D. 10

29

12. Lithium hydroxide reacts with carbon dioxide as follows.

2LiOH + CO2 → Li2CO3 +H2O

What mass (in grams) of lithium hydroxide is needed to react with 11 g of carbon dioxide?

A. 6 C. 24

B. 12 D. 48

13. Which solution contains the smallest amount of H+ ions?

A. 10.0 cm3 of 0.250 mol dm

–3 HCl C. 10.0 cm

3 of 0.250 mol dm

–3 H2SO4

B. 20.0 cm3 of 0.250 mol dm

–3 HCl D. 10.0 cm

3 of 0.500 mol dm

–3 HCl

14. How many hydrogen atoms are contained in one mole of ethanol, C2H5OH?

A. 5 C. 1.0 × 1023

B. 6 D. 3.6 × 1024

15. The percentage by mass of the elements in a compound is C = 72%, H = 12 %, O = 16%. What is the mole ratio of C : H in the empirical formula of this compound?

A. 1 : 1 C. 1 : 6

B. 1 : 2 D. 6 : 1

16. What is the coefficient for O2 (g) when the equation below is balanced?

__C3H8(g) + __O2(g) → __CO2(g) + __H2O(g)

A. 2 C. 5

B. 3 D. 7

17. What amount of NaCl (in moles) is required to prepare 250 cm3 of a 0.200 mol dm

−3 solution?

A. 50.0 C. 0.800

B. 1.25 D. 0.0500

18. 100 cm3 of ethene, C2H4, is burned in 400 cm

3 of oxygen, producing carbon dioxide and some liquid water.

Some oxygen remains unreacted.

(a) Write the equation for the complete combustion of ethene.

..................................................................................................................................... (b) Calculate the volume of carbon dioxide produced and the volume of oxygen remaining.

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

30

19. (a) Write an equation for the formation of zinc iodide from zinc and iodine.

.....................................................................................................................................

(b) 100.0 g of zinc is allowed to react with 100.0 g of iodine producing zinc iodide. Calculate the amount (in moles) of zinc and iodine, and hence determine which reactant is in excess.

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

(c) Calculate the mass of zinc iodide that will be produced.

.....................................................................................................................................

.....................................................................................................................................

20. Which of the following contains the greatest number of molecules?

A. 1 g of CH3Cl C. 1 g of CHCl3

B. 1 g of CH2Cl2 D. 1 g of CCl4

21. Which of the following compounds has/have the empirical formula CH2O?

I. CH3COOH

II. C6H12O6

III. C12H22O11

A. II only C. I and II only

B. III only D. II and III only

22. Consider the equation below.

Fe(s) + S(s) → FeS(s)

If 10.0 g of iron is heated with 10.0 g of sulfur to form iron (II) sulfide, what is the theoretical yield of FeS in grams?

A. 10.0 + 10.0 C.

B. D.

23. Assuming complete reaction, what volume of 0.200 mol dm

–3 HCl(aq) is required to neutralize 25.0 cm

3 of

0.200 mol dm–3

Ba(OH)2(aq)?

A. 12.5 cm3 C. 50.0 cm

3

B. 25.0 cm3 D. 75.0 cm

3

06.320.1091.87

06.320.1085.55

85.550.1091.87

31

24. An oxide of copper was reduced in a stream of hydrogen as shown below.

After heating, the stream of hydrogen gas was maintained until the apparatus had cooled.

The following results were obtained.

Mass of empty dish = 13.80 g Mass of dish and contents before heating = 21.75 g Mass of dish and contents after heating and leaving to cool = 20.15 g

(a) Explain why the stream of hydrogen gas was maintained until the apparatus cooled.

.....................................................................................................................................

.....................................................................................................................................

(b) Calculate the empirical formula of the oxide of copper using the data above, assuming complete re-duction of the oxide.

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

(c) Write an equation for the reaction that occurred.

.....................................................................................................................................

(d) State two changes that would be observed inside the tube as it was heated.

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

25. Consider the following equation.

2C4H10(g) + 13O2(g) → 8CO2(g) + 10H2O(1)

How many moles of CO2(g) are produced by the complete combustion of 58 g of butane, C4H10 (g)?

A. 4 C. 12

B. 8 D. 16

32

26. 6.0 moles of Fe2O3(s) reacts with 9.0 moles of carbon in a blast furnace according to the equation below.

Fe2O3(s) + 3C(s) → 2Fe(s) + 3CO(g)

What is the limiting reagent and hence the theoretical yield of iron?

27. What volume of 0.500 mol dm

−3 HCl(aq) is required to react completely with 10.0 g of calcium carbonate

according to the equation below?

CaCO3(s) + 2HCl(aq) → CaCl2(aq) + H2O(l) + CO2(g)

A. 100 cm3

C. 300 cm3

B. 200 cm3

D. 400 cm3

28. 27.82 g of hydrated sodium carbonate crystals, Na2CO3. xH2O, was dissolved in water and made up to 1.000 dm

3. 25.00 cm

3 of this solution was neutralized by 48.80 cm

3 of hydrochloric acid of concentration 0.1000 mol

dm−3

.

a) Write an equation for the reaction between sodium carbonate and hydrochloric acid.

..................................................................................................................................... b) Calculate the molar concentration of the sodium carbonate solution neutralized by the hydrochloric acid.

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

..................................................................................................................................... c) Determine the mass of sodium carbonate neutralized by the hydrochloric acid and hence the mass of sodium carbonate present in the1.000dm3 of solution.

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

..................................................................................................................................... d) Calculate the mass of water in the hydrated crystals and hence find the value of x.

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

Limiting reagent Theoretical yield of iron

A. Fe2O3 6.0 mol

B. Fe2O3 12.0 mol

C. carbon 9.0 mol

D. carbon 6.0 mol

33

29. Which is a correct definition of the term empirical formula?

A. formula showing the numbers of atoms present in a compound

B. formula showing the numbers of elements present in a compound

C. formula showing the actual numbers of atoms of each element in a compound

D. formula showing the simplest ratio of numbers of atoms of each element in a compound 30. The reaction of ethanal and oxygen can be represented by the unbalanced equation below.

__ CH3CHO + __ O2 → __ CO2 + __ H2O

When the equation is balanced using the smallest possible integers, what is the coefficient for O2?

A. 3 C. 5

B. 4 D. 6 31. The equation for the complete combustion of butane is

2C4H10 + 13O2 → 8CO2 + 10H2O

What is the amount (in mol) of carbon dioxide formed by the complete combustion of three moles of butane?

A. 4 C. 12

B. 8 D. 24 32. Which solution contains the greatest amount (in mol) of solute?

A. 10.0 cm3 of 0.500 mol dm–3

NaCl

B. 20.0 cm3 of 0.400 mol dm–3

NaCl

C. 30.0 cm3 of 0.300 mol dm–3

NaCl

D. 40.0 cm3 of 0.200 mol dm–3

NaCl 33. The percentage composition by mass of a hydrocarbon is C = 85.6 % and H = 14.4 %.

a) Calculate the empirical formula of the hydrocarbon.

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

..................................................................................................................................... b) A 1.00 g sample of the hydrocarbon at a temperature of 273 K and a pressure of 1.01 × 105 Pa (1.00 atm) has a volume of 0.399 dm3.

i) Calculate the molar mass of the hydrocarbon.

..........................................................................................................................

..........................................................................................................................

ii) Deduce the molecular formula of the hydrocarbon.

..........................................................................................................................

34

34. How many oxygen atoms are present in 0.0500 mol carbon dioxide?

A. 3.01 × 1022

C. 6.02 × 1023

B. 6.02 × 1022

D. 1.20 × 1024

35. The molar mass of an acid can be found by titrating it with a standard solution of sodium hydroxide.

2.04 g of an insoluble, dibasic organic acid (H2A) were dissolved in 20.0 cm3 of 2.00 mol dm-3 NaOH (an excess). At the conclusion of the reaction, the excess alkali (NaOH) was titrated with hydrochloric acid. The excess NaOH required 17.6 cm3 of 0.50 mol dm-3 HCl to neutralize it. a) Calculate the number of moles of sodium hydroxide that were added to the dibasic, organic acid.

...................................................................................................................................

...................................................................................................................................

................................................................................................................................... b) Write and balance an equation for the reaction of HCl with sodium hydroxide. ...................................................................................................................................

c) Determine the number of moles of HCl that reacted with the excess sodium hydroxide.

...................................................................................................................................

...................................................................................................................................

...................................................................................................................................

d) Using your answers from parts (b) and (c), determine the number of moles of sodium hydroxide that were in excess following the reaction with the dibasic acid.

...................................................................................................................................

e) Using your answers from parts (a) and (d), determine the number of moles of sodium hydroxide that re acted with the dibasic, organic acid.

...................................................................................................................................

...................................................................................................................................

f) Write and balance an equation for the reaction of the dibasic, organic acid (H2A) with sodium hydroxide.

...................................................................................................................................

g) Use your answers from parts (e) and (f) to determine the number of moles of the dibasic, organic acid that reacted.

...................................................................................................................................

h) Calculate the molar mass of the 2.04 gram sample of the dibasic, organic acid.

...................................................................................................................................

...................................................................................................................................

...................................................................................................................................

35

36. Two 1.00 dm3 containers A and B each contain 2.00 g of the gas indicated at 25.0C.

A B

a) Calculate the pressure, in kPa, in container B.

...................................................................................................................................

...................................................................................................................................

...................................................................................................................................

b) Deduce without calculation whether the pressure in A is higher or lower than container B and explain your answer.

...................................................................................................................................

...................................................................................................................................

H2

O2

36

37

Atomic Structure

2.1 – The Atom 1. State the position of protons, neutrons and electrons in the atom.

2. State the relative masses and relative charges of protons, neutrons and electrons.

3. Define the terms mass number (A), atomic number (Z) and isotopes of an element.

4. Deduce the symbol for an isotope given its mass number and atomic number.

5. Calculate the number of protons, neutrons and electrons in atoms and ions from the mass number,

atomic number and charge.

6. Compare the properties of the isotopes of an element.

7. Discuss the uses of radioisotopes.

2.2 – The Mass Spectrometer 1. Describe and explain the operation of a mass-spectrometer.

2. Describe how the mass spectrometer may be used to determine relative isotopic, atomic, and

molecular masses using the 12C scale.

3. Calculate non-integer relative atomic masses and abundance of isotopes from given data.

2.3 – Electron Arrangement 1. Describe the electromagnetic spectrum.

2. Distinguish between a continuous spectrum and a line spectrum.

3. Explain how lines in the emission spectrum of hydrogen are related to electron energy levels

4. Determine the electron arrangement up to Z=20.

12.4 – Electron Configuration (AHL) 1. Explain how evidence from first ionization energies across periods accounts for the existence of main

energy levels and sub-levels in atoms.

2. Explain how successive ionization energy data is related to the electron configuration of an atom.

3. State the relative energies of s,p,d, and f orbitals in a single energy level.

4. State the maximum number of orbitals in a given energy level.

5. Draw the shape of an s orbital and the shapes of the px, py, and pz orbitals.

6. Apply the Aufbau principle to electron configurations, Hund’s rule and the Pauli exclusion principle to

write the electron configurations for atoms up to Z=54.

7. Relate the electron configuration of an atom to its position in the periodic table.

38

Periodicity

3.1 – The Periodic Table 1. Describe the arrangement of elements in the periodic table in order of increasing atomic number.

2. Distinguish between the terms group and period.

3. Apply the relationship between the electron arrangement of elements and their position in the

periodic table up to Z = 20.

4. Apply the relationship between the number of electrons in the highest occupied energy level for

an element and it position in the periodic table.

3.2 – Physical Properties 1. Define the terms first ionization energy and electronegativity.

2. Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronega-

tivities and melting points for the alkali metals (Li to Cs) and halogens (F to I).

3. Describe and explain the trends in atomic radii, ionic radii, first ionization energies and electro-

negativities for elements across period 3.

4. Compare the relative electronegativity values of two or more elements based on their positions

in the periodic table.

3.3 – Chemical Properties 1. Discuss the similarities and differences in the chemical properties of elements in the same group.

2. Discuss the changes in nature, from ionic to covalent and from basic to acidic, of the oxides across

period 3.

13.1 – Trends Across Period 3 (AHL) 1. Explain the physical states (under standard conditions) and electrical conductivity (in the molten

state) of the chlorides and oxides of the elements in period 3 in terms of their bonding and

structure.

2. Describe the reactions of chlorine and the chlorides referred to in 13.1.1 with water.

13.2 – First Row d-block Elements (AHL) 1. List the characteristic properties of the transition elements.

2. Explain why Sc and Zn are not considered to be transition elements.

3. Explain the existence of variable oxidation number in ions of transition elements.

4. Define the term ligand.

5. Describe and explain the formation of complexes of d-block elements.

6. Explain why some complexes of d-block elements are colored.

7. State examples of the catalytic action of transition elements and their compounds.

8. Outline the economic significance of catalysts in the Contact and Haber processes.

39

TTOPICOPIC 2 2 ANDAND 3 P3 PROBLEMROBLEM SSETET –– AATOMICTOMIC TTHEORYHEORY ANDAND PPERIODICITYERIODICITY Atomic Theory

1. What is the atomic number? What information does it provide? 2. What is the mass number? How is this different from atomic mass?

3. According to the most current model of the atom, describe the location, charge, and relative mass of protons, neu-

trons and electrons. A labeled diagram may help.

4. Define isotope.

5. Four elements are described below:

A) Which elements are isotopes of each other? (2 pairs)

B) How will elements A and D compare in terms of chemical/physical properties?

C) How will elements A and D compare in terms of atomic mass?

D) How will elements B and C compare in terms of charge?

E) Draw atomic symbols for all four elements. 6. Fill in the chart:

Element # of protons # of neutrons # of electrons

A 5 6 2

B 1 1 1

C 1 0 0

D 5 5 5

Isotope # of

protons # of

neutrons # of

electrons

40

7. Fill in the chart:

8. Sketch a simple diagram of a mass spectrometer. Label each phase of the mass spectrometer’s operation on the diagram.

9. What specific piece of a mass spectrometer is adjusted in order to detect particles with varying masses?

10. Of the following particles, 16

O+,15

O+,

16O

2+, (16

O –16

O)+ which experiences the greatest amount of deflection in a mass

spectrometer? Why are all ions formed in a mass spectrometer positively charged?

11. The following graph is produced when a pure sample of boron is passed through a mass spectrometer. Use the

data to calculate the relative atomic mass for boron.

12. Lithium occurs naturally as two isotopes,

6Li and

7Li. The relative atomic mass of

lithium is 6.941 g/mol. Determine the percent abundance of each of lithium’s isotopes.

13. State the name and the mass number of the isotope relative to which all atomic masses are measured.

14. Write the symbol for the species with 17 protons, 19 neutrons, and 18 electrons.

15. Write the symbol for the species with 6 protons, 8 neutrons, and 6 electrons.

16. Write the symbol for the species with 3 protons, 3 neutrons, and 2 electrons.

Atomic # Symbol # protons Mass # # neutrons Charge # electrons

As 42 -3

64 155 61

12 25 +2

I-

74 54

48 115 48

58 82 +4

16 32 -2

41

17. Explain why most atomic masses are not whole numbers.

18. In the context of electromagnetic radiation, what is wavelength? How does wavelength relate to energy?

19. In the context of electromagnetic radiation, what is frequency? How does frequency relate to energy?

20. List the colors of visible light in order of increasing energy.

21. a) Indicate the wavelength of each wave in the following wave diagrams.

b) Which of the waves has a higher frequency?

c) Which of the waves has a lower energy?

Wave A

Wave B

22. As wavelength gets shorter, frequency ______________. 23. As wavelength gets shorter, energy _______________. 24. Describe the Bohr model of the atom, including the evidence collected that led to its development. 25. What is the difference between a bright line spectrum and a continuous spectrum? How do energy levels account

for the appearance of bright line spectra? 26. What is the electromagnetic (EM) spectrum? What is the highest energy wave? Lowest? 27. Consider this diagram of an atom with arrows representing electron movement.

a) Which two arrows correspond to energy absorption by the atom? _________ b) Which two arrows correspond to energy emission by the atom? _________ c) If violet and green light are produced by the movement illustrated here, which

arrow represents emission of violet light? _____ green light? _____ nucleus

42

28. Consider the spectroscopy lab. How did the flame tests you performed on different metals relate to firework produc-tion? Why did each metal produce a different color flame when excited?

29. Consider the spectroscopy lab. What caused the gas in the tubes to glow? How does this topic relate to “neon”

sign production? Why is the term “neon sign” a bit misleading? 31. How do scientists use bright line spectra to identify unknown elements in far off stars or other samples of matter? 32. Consider only the first four primary energy levels of a hydrogen atom. The transition that would result in photon of

the shortest wavelength would be a) from n = 4 to n = 1 b) from n = 4 to n = 3 c) from n = 2 to n = 1 d) from n = 1 to n = 4 33. Consider only the first four primary energy levels of a hydrogen atom. The transition that would result in photon of

the longest wavelength would be a) from n = 4 to n = 1 b) from n = 4 to n = 3 c) from n = 2 to n = 1 d) from n = 1 to n = 4 34. The Lyman series of bright lines in the hydrogen atom are due to electrons dropping to the first energy level, the

Balmer series are due to electrons dropping to the second energy level. Which one of these series is U.V. and which is visible? Explain.

35. How do waves of red light and blue light differ with respect to frequency? wavelength? energy? 36. How many electrons can a single atomic orbital hold? How many orbitals can be found in an s sublevel? p? d? f? 37. “s” sublevels can hold a total of ___ electrons. p sublevels can hold ___ electrons, while d sublevels can hold ___,

and f sublevels can hold ___ electrons. 38. Which is bigger, the 3s sublevel or the 5s sublevel? How many electrons can each hold? 39. What are valence electrons and why are they important? What is the octet rule? 40. Every element wants a full outer energy level (valence level). This is normally ___ electrons, although in the case of

helium it is ____ electrons. 41. An atom is in the _______________________________state when the electrons in an atom are in the lowest

possible energy levels. 42. An atom is in the ____________________________ state when one or more electrons moves to a higher than

normal energy level. 43. Which of the following has the highest energy? A) 4d B) 5s C) 5p D) 3p 44. Which of the following has the lowest energy? A) 6s B) 5p C) 4f D) 6p

43

45. Give electron complete configurations for: Na+, Fe, Br, Ar, Al

+3, O

-2, He, Ni

+2, K, Ne, Cu

+.

46. Give shorthand electron configurations for: Ag, I, Rb, Au, Cu, S

-2

47. What does isoelectronic mean? Give three elements that are isoelectronic with Kr. 48. 1s

22s

22p

63s

23p

4 is the electron configuration for which element? How many valence electrons does the element

have? 49. 1s

22s

22p

63s

23p

64s

23d

104p

6 is the electron configuration for which noble gas? How can you use its electron

configuration to confirm it is a noble gas? Periodicity

50. What is meant by effective nuclear charge? 51. Compare the effective nuclear charge on the valence electrons of sodium, magnesium and aluminum. 52. How can effective nuclear charge explain the observed difference in first ionization energy between sodium and

magnesium? Consult table 7 of the data booklet for data regarding first ionization energies. 53. Despite the trend in effective nuclear charge, magnesium has a higher first ionization than aluminum. Explain this in

terms of electron configuration.

44

54. Sketch a graph of successive ionization energies for Mg.

Ionization Energy 1 2 3 4 Successive ionizations 55. Sketch a graph of successive ionization energies for Al. Ionization Energy 1 2 3 4 Successive ionizations 56. Write a chemical equation that shows which is more reactive: a) Na or K b) Cl2 of F2 57. Consider a reaction where gaseous chlorine is bubbled through a solution of potassium bromide. a) Write and balance an equation for the reaction, including state symbols. b) Write the net ionic equation for the reaction c) Describe the color change that would be observed as the reaction progresses. d) Explain why the reaction would not proceed in the reverse direction.

45

58. A solution of silver (I) nitrate is added to an unknown solution of halide ions. Upon mixing, a faint yellow precipitate is produced. What is the likely identity of the halide ions?

59. Consider the 3

rd period elements.

a) Write the formula of their common oxides b) Write a chemical equation (or equations) to show if sodium oxide is acidic, basic, or amphoteric

c) Write a chemical equation (or equations) to show that sulfur dioxide is acidic, basic, or amphoteric

d) Write a chemical equation (or equations) to show that aluminum oxide is acidic, basic, or amphoteric.

e) Describe how the pH of a solution of sodium oxide would differ from the pH of a solution of sodium chloride (acidic solutions have a pH < 7, basic solutions have a pH > 7).

60. Consider the complex ion [Cu(H2O)6]

2+

a) State the oxidation state (charge) of the transition element. b) Hydrated Cu

2+ ions have a light blue color. Explain why the species is colored.

c) Explain why complex ions of zinc are colorless. d) What is the coordination number of the ion?

e) What is the ligand? If the water molecules were replaced with ammonia (NH3) in a ligand exchange reaction, would the solution appear to be the same color, or would the color be changed?

61. Consider the following compounds containing complex ions of iron:

Na[FeCl4] [Fe(H2O)6]SO4 [Fe(H2O)6]SO4 [Fe(H2O)4]PO4

a) State the oxidation state (charge) for iron in each of the four compounds. b) State the coordination number for each of the complex ions. c) Identify three factors that affect the color of complex ions containing transition metals. 62. Write balanced equations, including state symbols, for the Contact Process and the Haber Process (include the

catalyst). Describe these as examples of homogeneous catalysis or heterogeneous catalysis.

46

47

MMASSASS SSPECTROMETRYPECTROMETRY SSIMULATIONIMULATION

Problem:

Determine how various part of a mass spectrometer function.

Method:

1. Label your mass spectrometer with the following main steps: vaporized sample introduction, ionization chamber,

acceleration, deflection, detection.

2. Collect one set of three “isotopes.” Each isotope should have a different mass. You should have one tennis ball,

one billiard ball, and one polystyrene ball.

3. Determine the mass of each isotope in grams and record this in the table.

4. Have one group member roll each isotope through the mass spectrometer while another member acts as a

“deflector.” Another group member should record the time required for the isotope to move from the beginning

of the “acceleration” step to the “detection” step.

5. As the isotope passes the deflection area, the “deflector” should use only the force of their breath to alter the

path of the isotope. No hands allowed! The isotope must not touch the walls of the spectrometer, only the end

of the detector. The “deflector” should judge the strength of their “deflection force” on a scale of 1 – 10; 1 being

a very light breeze and 10 being gale force hurricane breath! Record this number in the table, as well as the

time required for the isotope to move through the instrument.

6. It is very important that the person introducing the isotopes into the machine attempt to move them at a

consistent speed! There shouldn’t be much variation in the time required.

7. Successfully pass each of the three isotopes through the mass spectrometer a total of five times, for a total of

fifteen data points.

Data table:

Observations -

Isotope mass (± .01 g)

Force required for deflection (1 -10)

Time required to pass through instrument (sec.)

Trial 1 Trial 2 Trial 3 Avg. Trial 1 Trial 2 Trial 3 Avg.

48

Questions:

1. Which isotope required the greatest force to deflect? Does this match your prediction?

2. What physical concept is related to an object’s resistance to changes in motion state? How is that related to

this activity?

3. Suppose you were to change the speed at which the isotopes passed through the spectrometer. How would

altering the speed, both increasing it and decreasing it, affect the ease with which the isotopes could be de-

flected? Experiment with each scenario and report your findings here.

4. Explain why the deflection force has to be “calibrated” to allow for isotopes of different masses to be detected.

49

Name _______________________ Per _____

SPECTROSCOPY LAB

STATION 1:

Description: __________________________________________________________________________

_____________________________________________________________________________________

Bright Line Spectrum

STATION 2:

Description: __________________________________________________________________________

_____________________________________________________________________________________

Bright Line Spectrum

STATION 3:

Description: __________________________________________________________________________

_____________________________________________________________________________________

Bright Line Spectrum

STATION 4:

Description: __________________________________________________________________________

_____________________________________________________________________________________

Bright Line Spectrum

STATION 5:

Description: __________________________________________________________________________

_____________________________________________________________________________________

Bright Line Spectrum

50

Questions:

1. Which line is emitting photons traveling in the longest wavelength for the bright line spectrum at station 1?

2. Are the spectra you viewed in this experiment continuous spectra or bright line spectra? Explain the differences be-

tween the two types of spectra.

3. How is visible light created by an atom? Be specific about the role of energy levels in this process.

4. Which line is emitting photons at the highest energy (highest frequency) at station 4?

5. Which station is the element hydrogen? Explain how you deduced your answer.

6. Outline the electromagnetic spectrum.

7. Draw a diagram to show how red light might be created in an atom. Be sure to use the appropriate energy levels for

creating visible light.

8. Why do some elements have more lines in their spectra? Be sure to use energy levels in your explanation.

51

Name _______________________ Per _____

Graphing Periodic Trends

Purpose: Observe the patterns in periodic trends.

Graphing:

Use the table below to graph atomic radius. See table 8 in you data booklet for atomic radius values. Label

your x-axis with atomic number and the y-axis with atomic size.

Use the table below to graph first ionization energy (energy required to remove an electron from a gaseous

atom). See table 7 for values. Label your x-axis with atomic number and the y-axis with ionization energy.

Element Atomic

Radius

Element Atomic

Radius

Hydrogen Neon

Helium Sodium

Lithium Magnesium

Beryllium Aluminum

Boron Silicon

Carbon Phosphorus

Nitrogen Sulfur

Oxygen Chlorine

Fluorine Argon

Element Ionization

Energy

(kJ/mol)

Element Ionization

Energy

(kJ/mol)

Hydrogen Neon

Helium Sodium

Lithium Magnesium

Beryllium Aluminum

Boron Silicon

Carbon Phosphorus

Nitrogen Sulfur

Oxygen Chlorine

Fluorine Argon

52

Use the table below to graph electronegativity (pull of an atom on an electron in a neighboring atom). See table

7 in your data booklet for values. Label your x-axis with atomic number and the y-axis with electronegativity.

Questions:

1. Predict the values for the atomic radius, first ionization energy, and electronegativity of potassium (Z = 19).

2. Which group of elements appears to occupy the major peaks for atomic radius? _________________________

3. Which group of elements appears to occupy the major peaks for ionization energy? Why might that be?

4. Why wouldn’t noble gases have electronegativity?

5. Do metals or non-metals have the highest electronegativities? Why might that be?

6. What might cause the decrease in ionization energy for elements in group 3A compared to those in 2A? (Hint:

think about electron configuration)

7. Compare and contrast the trends for atomic radius and ionization energy.

8. Compare and contrast the trends for atomic radius and electronegativity.

9. Compare and contrast the trends for ionization energy and electronegativity.

Element Electro-

negativity

Element Electro-

negativity

Hydrogen Neon n/a

Helium n/a Sodium

Lithium Magnesium

Beryllium Aluminum

Boron Silicon

Carbon Phosphorus

Nitrogen Sulfur

Oxygen Chlorine

Fluorine Argon n/a

53

PPASTAST IB TIB TESTEST QQUESTIONSUESTIONS -- TTOPICSOPICS 2 2 ANDAND 33 1. Consider the composition of the species W, X, Y and Z below. Which species is an anion?

A. W C. Y

B. X D. Z 2. Energy levels for an electron in a hydrogen atom are

A. evenly spaced. C. closer together near the nucleus.

B. farther apart near the nucleus. D. arranged randomly.

3. Which is related to the number of electrons in the outer main energy level of the elements from the alkali

metals to the halogens?

I. Group number

II. Period number

A. I only C. Both I and II

B. II only D. Neither I nor II

4. The element vanadium has two isotopes, and , and a relative atomic mass of 50.94.

a) Define the term isotope.

. ..................................................................................................................................

b) State the number of protons, electrons and neutrons in .

...................................................................................................................................

c) State and explain which is the more abundant isotope.

...................................................................................................................................

...................................................................................................................................

d) State the name and the mass number of the isotope relative to which all atomic masses are measured.

...................................................................................................................................

V50

23 V51

23

V50

23

Species Number of protons Number of neutrons Number of electrons

W 9 10 10

X 11 12 11

Y 12 12 12

Z 13 14 10

54

5. The diagram below (not to scale) represents some of the electron energy levels in the hydrogen atom.

________________________________ n = ∞ ________________________________ n = 6 ________________________________ n = 5

________________________________ n = 4

________________________________ n = 3

________________________________ n = 2

________________________________ n = l i) Draw an arrow on the diagram to represent the electron transition for the ionization of hydro-

gen. Label this arrow A.

ii) Draw an arrow on the diagram to represent the lowest energy transition in the visible emission spectrum. Label this arrow B.

6. What is the correct number of each particle in a fluoride ion, 19F–?

7. Which statement is correct for the emission spectrum of the hydrogen atom?

A. The lines converge at lower energies.

B. The lines are produced when electrons move from lower to higher energy levels.

C. The lines in the visible region involve electron transitions into the energy level closest to the nucleus.

D. The line corresponding to the greatest emission of energy is in the ultraviolet region. 8. Electrons are directed into an electric field from left to right as indicated by the arrow in the diagram below. Which

path is most probable for these electrons?

A. 1 C. 3

B. 2 D. 4

protons neutrons electrons

A. 9 10 8

B. 9 10 9

C. 9 10 10

D. 9 19 10

1

2

34

+

–

55

9. How many valence electrons are present in an atom of an element with atomic number 16?

A. 2 C. 6

B. 4 D. 8

10. a) Evidence for the existence of energy levels in atoms is provided by line spectra. State how a line spectrum differs from a continuous spectrum.

...................................................................................................................................

...................................................................................................................................

b) On the diagram below draw four lines in the visible line spectrum of hydrogen.

(c) Explain how the formation of lines indicates the presence of energy levels.

...................................................................................................................................

...................................................................................................................................

11. A certain sample of element Z contains 60% of 69

Z and 40% of 71

Z. What is the relative atomic mass of ele-ment Z in this sample?

A. 69.2 C. 70.0

B. 69.8 D. 70.2

12. What is the difference between two neutral atoms represented by the symbols Co and Ni?

A. The number of neutrons only.

B. The number of protons and electrons only.

C. The number of protons and neutrons only.

D. The number of protons, neutrons and electrons.

13. Which pair of elements reacts most readily?

A. Li + Br2 C. K + Br2

B. Li + Cl2 D. K + Cl2

14. For which element are the group number and the period number the same?

A. Li C. B

B. Be D. Mg

Low energy High energy

27

59

28

59

56

15. Explain the following statements.

a) The first ionization energy of sodium is

i) less than that of magnesium.

…………………………………………………………………………………

…………………………………………………………………………………

ii) greater than that of potassium.

…………………………………………………………………………………

…………………………………………………………………………………

b) The electronegativity of chlorine is higher than that of sulfur.

……………………………………………………………………………………….

……………………………………………………………………………………….

……………………………………………………………………………………….

16. The following table shows values that appear in the Data Booklet.

Table 1 Covalent (atomic) radii / 10–12 m

Explain why

i) the magnesium ion is much smaller than the magnesium atom.

…………………………………………………………………………………………..

…………………………………………………………………………………………..

…………………………………………………………………………………………..

ii) there is a large increase in ionic radius from silicon to phosphorus.

…………………………………………………………………………………………..

…………………………………………………………………………………………..

…………………………………………………………………………………………..

N 0 F

70 66 58

Na Mg Al Si P S Cl

186 160 143 117 110 104 99

Table 2 Ionic radii/10–12 m

N3–

O2–

F–

171 146 133

Na+ Mg

2+ Al

3+ Si

4+ P

3– S

2– Cl

–

98 65 45 42 212 190 181

57

iii) the ionic radius of Na+ is less than that of F–.

…………………………………………………………………………………………..

…………………………………………………………………………………………..

…………………………………………………………………………………………..

17. Which of the following properties of the halogens increase from F to I?

I. Atomic radius

II. Melting point

III. Electronegativity

A. I only C. I and III only

B. I and II only D. I, II and III

18. Table 6 of the Data Booklet lists melting points of the elements. Explain the trend in the melting points of the alkali metals, halogens and period 3 elements.

........................................................................................................................................................

........................................................................................................................................................

........................................................................................................................................................

........................................................................................................................................................

........................................................................................................................................................

........................................................................................................................................................

........................................................................................................................................................

........................................................................................................................................................

19. i) Explain how the first ionization energy of K compares with that of Na and Ar.

…………………………………………………………………………………………..

…………………………………………………………………………………………..

…………………………………………………………………………………………..

ii) Explain the difference between the first ionization energies of Na and Mg.

…………………………………………………………………………………………..

…………………………………………………………………………………………..

…………………………………………………………………………………………..

…………………………………………………………………………………………..

58

iii) Suggest why much more energy is needed to remove an electron from Na+ than from Mg+.

…………………………………………………………………………………………..

…………………………………………………………………………………………..

…………………………………………………………………………………………..

20. Rubidium is an element in the same group of the periodic table as lithium and sodium. It is likely to be a metal which has a

A. high melting point and reacts slowly with water.

B. high melting point and reacts vigorously with water.

C. low melting point and reacts vigorously with water.

D. low melting point and reacts slowly with water.

21. When the following species are arranged in order of increasing radius, what is the correct order?

A. Cl–, Ar, K

+ C. Cl

–, K

+, Ar

B. K+, Ar , Cl

– D. Ar, Cl

–, K

+

22. Table 8 of the Data Booklet gives the atomic and ionic radii of elements. State and explain the difference between

i) the atomic radius of nitrogen and oxygen.

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

ii) the atomic radius of nitrogen and phosphorus.

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

…………………………………………………………………………………………… iii) the atomic and ionic radius of nitrogen.

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

59

24. What increases in equal steps of one from left to right in the periodic table for the elements lithium to ne-on?

A. the number of occupied electron energy C. the number of electrons in the atom levels

B. the number of neutrons in the most D. the atomic mass common isotope

25. State and explain the trends in the atomic radius and the ionization energy

i) for the alkali metals Li to Cs.

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

ii) for the period 3 elements Na to Cl.

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

26. i) Describe three similarities and one difference in the reactions of lithium and potassium with water.

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

ii) Give an equation for one of these reactions. Suggest a pH value for the resulting solution, and give a reason for your answer.

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

60

61

Bonding

4.1 – Ionic bonding 1. Describe the ionic bond as the electrostatic attraction between oppositely charged ions.

2. Describe how ions can be formed as a result of electron transfer.

3. Deduce which ions will be formed when elements in groups 1, 2 and 3 lose electrons.

4. Determine which ions will be formed when elements in groups 5, 6 and 7 gain electrons.

5. State that transition elements can form more than one ion.

6. Predict whether a compound of two elements would be ionic from the position of the elements in

the periodic table or from their electronegativity values.

7. State the formula of common polyatomic ions formed by non-metals in periods 2 and 3.

8. Describe the lattice structure of ionic compounds.

4.2 – Covalent Bonding 1. Describe the covalent bond as the electrostatic attraction between a pair of electrons and

positively charged nuclei.

2. Describe how the covalent bond is formed as a result of electron sharing.

3. Deduce the Lewis (electron dot) structure of molecules and ions for up to six electron pairs on

each atom.

4. State and explain the relationship between the number of bonds, bond length and bond strength.

5. Predict whether a compound of two elements would be covalent from the position of the

elements in the periodic table or from their electronegativity values.

6. Predict the relative polarity of bonds from electronegativity values.

7. Predict the shape and bond angles for species with four, three and two negative charge centers

on the central atom using the valence shell electron pair repulsion theory (VSEPR).

8. Predict whether or not a molecule is polar from its molecular shape and bond polarities.

9. Describe and compare the structure and bonding in three allotropes of carbon (diamond, graph-

ite and C60 fullerene).

10. Describe the structure of and bonding in silicon and silicon dioxide.

4.3 – Intermolecular Forces 1. Describe the types of intermolecular forces (attractions between molecules that have temporary

dipoles, permanent dipoles or hydrogen bonding) and explain how they arise from the structural

features of molecules.

2. Describe and explain how intermolecular forces affect the boiling points of substances.

4.4 – Metallic Bonding 1. Describe the metallic bond as the electrostatic attraction between a lattice of positive ions and

delocalized electrons.

2. Explain the electrical conductivity and malleability of metals.

62

4.5 – Physical Properties 1. Compare and explain the properties of substances resulting from different types of bonding

14.1 – Shapes of Molecules and Ions (AHL) 1. Predict the shape and bond angles for species with five and six negative charge centers using the

VSEPR theory.

14.2 – Hybridization (AHL) 1. Describe σ and π bonds.

2. Explain hybridization in terms of the mixing of atomic orbitals to form new orbitals for bonding.

3. Identify and explain relationships between Lewis structures, molecular shapes, and types of

hybridization (sp, sp2, and sp3).

14.3 – Delocalization of Electrons (AHL) 1. Describe the delocalization of π electrons and explain how this can account for the structures of

some substances.

63

TTOPICOPIC 4 P4 PROBLEMROBLEM SSETET –– BBONDINGONDING

Ionic Bonding

1. An ionic bond is characterized as the ___________ of electrons between an element with ______ electronegativity

and an element with ________ electronegativity (typically a metallic element and a nonmetallic element).

2. An element has 19 protons. What will be the charge of ions it forms?

3. What will be the charge of the ions formed by an element with an electron structure of 1s22s

22p

63s

23p

5?

4. Valence electrons are only found in the ____ and ____ sublevels, not in the ____ or ____.

5. The table below shows the electron structures for 5 pairs of elements. Predict the compound that will form. The

first one is completed as an example.

6. Describe the crystal structure that occurs in solid ionic compounds, and why this is not considered a molecular solid.

Covalent Bonding

7. Covalent bonding is defined as the _____________ of electrons between an element with ________

electronegativity and another element with ________ electronegativity (typically two nonmetallic elements).

8. An element forms a covalently bonded compound with hydrogen, and has the formula XH4. In which group of the

periodic table would X be found? What is the name of this group?

9. Two atoms each provide 3 electrons that are shared by the two atoms. This is an example of a :

A. single covalent bond C. triple covalent bond

B. double covalent bond D. quadruple covalent bond

10. Two atoms each provide 1 electron that are shared by the two atoms. This is an example of a :

A. single covalent bond C. triple covalent bond

B. double covalent bond D. quadruple covalent bond

11. A certain nonmetal usually forms two covalent bonds in its compounds. How many valence electrons does this

element have?

12. Noble gases do not form chemical compounds because:

Element A Element B Formula

1s22s

22p

63s

2 1s

22s

22p

63s

23p

64s

23d

104p

5 AB2

1s22s

1 1s

22s

22p

5

1s22s

22p

1 1s

22s

22p

63s

23p

4

1s22s

22p

63s

1 1s

22s

22p

4

64

13. Which of the following compounds contains ionic and covalent bonds?

A. SiO2 C. Na2CO3

B. BaF2 D. Cl2O

14. Predict the number of covalent bonds that would form and then predict a formula for:

a. nitrogen and chlorine

b. chlorine and hydrogen

c. silicon and fluorine

Shapes of Molecules

a) Draw a Lewis structure

b) Predict the orbital geometry and molecular geometry

c) Give the bond angles

d) tell whether the molecule is polar or nonpolar:

15. NH4+

16. H2O 17. CH4

18. CO2 19. NH3 20. SO2

21. BCl3 22. C2H4 23. C2H2

24. CH3Cl 25. CH2O 26. HCl

27. F2

65

28. A molecule that has a tetrahedral shape would have _____ nonbonding or “lone” pairs of electrons.

29. A molecule that has a trigonal pyramidal shape would have _____ nonbonding or “lone” pairs of electrons.

30. Carbon and chlorine form a series of compounds: CH4, CH3Cl, CH2Cl2, CH3Cl, CCl4. Which of these will be polar

molecules?

A. CCl4 only C. CH3Cl, CH2Cl2, and CH3Cl only

B. CH3Cl and CH3Cl only D. CH3Cl, CH3Cl, CH2Cl2, and CCl4 only

31. For each of the following pairs, which atom (if any) in the following bonds carries a partial negative charge?

a. H---H b. O---P c. C---F

d. S---S e. B---O

32. The proper Lewis structure for the NO3- ion contains resonance structures. What is the purpose of showing

resonance structures?

33. What is the N-O bond order in NO3-? Explain.

34. Which would have the longest N-O bond, NO2- or NO3

-? Explain.

35. Draw a Lewis structure for C2F4.

Sketch a bonding diagram of the molecule (including 3-dimensional representations of the bonding orbitals).

Label σ and π bonds.

66

36. Consider C2H6 (ethane) and C2H4 (ethene).

a) Draw a Lewis diagram for each.

b) Count the total number of sigma and pi bonds in each:

C2H6: σ _____ π _____ C2H4: σ _____ π _____

c) Describe the formation of sigma and pi bonds.

d) Compare the C – H bond angles in each.

e) Compare the carbon-carbon bond order, bond length, and bond strength in each.

37. Complete the chart:

Formu-la

Lewis Structure Hybridization of Central atom

Orbital geom-etry

3- D picture with polarity vectors

molecular geometry

po-lar?

SiF4

BrF5

ClF3

NO2-

NO3-

SBr4

67

Formula Lewis Structure Hybrid. of Central atom

Orbital geometry

3- D picture with polarity vectors

Molecular geometry

Polar?

CN-

XeF4

BF3

OF2

CO32-

PF5

PI3

SF6

XeF2

CS2

68

Properties and Bonding

38. In which of the following compounds would hydrogen bonding occur?

A. COCl2 C. H2CO

B. PH3 D. CH3OH

39. Which of the following molecules would have the highest melting point?

A. CH3-CH2-CH2-CH2-CH2-CH3 C. CH3-CH2CH(CH3)-CH2-CH2-CH3

B. CH3-CH(CH3)-CH2-CH2-CH3 D. CH3-C(CH3)2-CH2-CH3

40. In which of the following substances would there be the strongest forces between the molecules?

A. SiH4 C. CH3-CH3

B. H2C=O D. O2

41. Explain why at room temperature chlorine is a gas, bromine is a liquid, and iodine is a solid.

42. Explain why water (H2O) is liquid at room temperature but hydrogen sulfide (H2S) is a gas.

43. Explain why covalently bonded substances are poor conductors of heat and electricity.

44. Why do molecular solids have low melting points?

45. Why do NH3, H2O, and HF have abnormally high boiling points when compared to their analogs PH3, H2S, and HCl?

46. Describe the bonding that occurs in a metallic solid.

47. Use the commonly accepted model of metallic bonding to explain why:

a. the melting/boiling points of metals in period 3 increase from sodium to magnesium to aluminum

b. metals are malleable

c. metals conduct electricity in the solid state

69

48. Classify these substances as ionic, molecular, network, or metallic solids:

a. structure composed of atoms covalently bonded to neighboring atoms

b. a solid only at extremely low temperatures

c. a good conductor of heat and electricity

d. a good electric conductor only in solution

49. Which of the following substances would be soluble in water?

A. CH3-CH2-CH2-CH2-CH2-CH2-CH3 C. Cl-CH2-CH2-CH2-CH2-Cl

B. H2N-CH2-CH2-CH2-CH2-NH2 D. CH3-CH2-CH2-O-CH2-CH2-CH2-CH3

50. When ethanol (C2H5OH) boils, the gas consists of

A. a mixture of carbon dioxide and water C. water and ethanol

B. carbon, hydrogen, and oxygen D. ethanol only

51. Elements A,B, C, and D have consecutive atomic numbers. Element D is a monatomic gas with low melting and

boiling points. All efforts to form compounds of D in the laboratory have failed.

a. Which of the remaining elements, A, B, or C has the strongest affinity for an additional electron?

b. A compound of an alkali metal M with element C has a formula of MC. Does this compound have ionic or

covalent bonds? Predict other properties of MC such as melting point and solubility in water.

c. Write the formulas for hydrides of elements A, B, and C.

d. Predict the shape of each hydride molecule from above.

e. Predict the conductivity of solid B.

52. For each of the following liquids, list the type of intermolecular forces you would expect to find.

a. water, H2O

b. bromine, Br2

c. carbon tetrachloride, CCl4

70

53. Explain the reasons for the difference in boiling points between

HF (20 C) and HCl (-85 C)

HCl (-85 C) and (LiCl (1360 C)

CH2CH2CH2OH (78.4 C) and CH2CH2CH2CH2CH2CH2OH (157 C)

54. Match each of the solids in the first column with two properties in the second column. Try to use each property at

least once.

a. metallic solid I. low melting point

b. covalent network solid II. high melting point

c. ionic solid III. conducts electricity in solution

d. molecular solid IV. brittle

V. hard

VI. malleable

55. Describe and compare the structure and bonding present in the three allotropes of carbon (diamond, graphite, and

C60 fullerene).

For each of the following pairs in 54 – 60, circle the compound that would have the higher boiling point. Briefly explain.

56. RbF or NaF

57. CH3-CH2-CH2-CH2-CH2-CH3 or CH3-C(CH3)2-CH2-CH3

58. Mg or Na

59. PCl3 or MgCl2

60. I2 or Cl2

61. CH3-CH2 -OH or CH3-O-CH3

62. H2O or H2S

71

63. Circle the formula of the compound that would be more soluble in water. Explain.

CH3-CH2- CH2- CH2-CH2-OH or CH3-CH2-OH

CH3-CH2 -OH or CH3-O-CH3

64. ______substances are malleable and conduct electricity

A. metallic B. ionic C. network covalent D. molecular covalent

65. ______substances have low boiling points and do not conduct electricity in any state

A. metallic B. ionic C. network covalent D. molecular covalent

66. ______substances have high melting points and do not conduct electricity in any state

A. metallic B. ionic C. network covalent D. molecular covalent

67. ______substances have high melting points and conduct electricity in liquid (not solid) state

A. metallic B. ionic C. network covalent D. molecular covalent

68. Compare the three allotropes of carbon in terms of the hybridization of the carbon atoms, the molecular geometry

around the carbon atoms, their electrical conductivities, and their hardness. Describe the 3-dimensional structure

shown by each of these allotropes. Identify the allotrope of carbon that is most like silicon dioxide.

72

73

VVISCOSITYISCOSITY OFOF AA LLIQUIDIQUID -- DDESIGNESIGN LLABAB

There are several factors at work that affect the relative “thickness” of a liquid, or more specifically, its resistance to flow.

For molecular covalent substances, the viscosity of a liquid results from the internal friction between the molecules re-

sulting from intermolecular forces. In this investigation, you will design an experiment to investigate a single factor af-

fecting a liquid’s viscosity.

Objective:

Design and carry out an experiment to investigate a factor affecting the viscosity of a liquid.

Phase 1: Designing the Investigation

Research Question:

Begin by defining a research question that will guide your investigation.

________________________________________________________________________________________________

________________________________________________________________________________________________

________________________________________________________________________________________________

Hypothesis:

Formulate a hypothesis addressing your research question. For this investigation, your hypothesis should provide a pre-

diction of how you believe your independent variable will affect the dependent variable. Your hypothesis should be sup-

ported by a well developed rationale based on sound chemical principles.

________________________________________________________________________________________________

________________________________________________________________________________________________

________________________________________________________________________________________________

________________________________________________________________________________________________

________________________________________________________________________________________________

Dependent Variable:

Identify the dependent variable that will be measured. Describe how it will be measured.

________________________________________________________________________________________________

________________________________________________________________________________________________

________________________________________________________________________________________________

Independent Variable:

Identify the independent variable that will be manipulated. Describe how it will be manipulated.

________________________________________________________________________________________________

________________________________________________________________________________________________

________________________________________________________________________________________________

74

Controls:

Provide a list of controls you will employ to ensure that the only factor that is altered between successive trials is the in-

dependent variable. Explain why each control is selected by providing a brief description of what effect changing this

factor would have on the dependent variable.

________________________________________________________________________________________________

________________________________________________________________________________________________

________________________________________________________________________________________________

________________________________________________________________________________________________

________________________________________________________________________________________________

________________________________________________________________________________________________

________________________________________________________________________________________________

________________________________________________________________________________________________

________________________________________________________________________________________________

________________________________________________________________________________________________

Method:

Briefly describe the method you will use to test your variables. A minimum of five data points should be collected to

meet the requirement of collecting sufficient data.

________________________________________________________________________________________________

________________________________________________________________________________________________

________________________________________________________________________________________________

________________________________________________________________________________________________

________________________________________________________________________________________________

________________________________________________________________________________________________

________________________________________________________________________________________________

________________________________________________________________________________________________

75

Phase 2: Collecting and Processing Data

Carry out your approved method and collect all relevant data, both quantitative and qualitative. Remember to include

units in your presentation as well as uncertainties. Keep in mind that any raw measurement you take has a random un-

certainty associated with it that must be recorded. Processed measurements have a random uncertainty equal to the

sum of the random uncertainties of the raw measurements incorporated in the processed measurement. Consult the

chapter 11 excerpt for guidance. A graph displaying the trend observed is encouraged.

Phase 3: Analyzing the data and Evaluating the Investigation

Formulate a conclusion that directly addresses your research question by interpreting the processed data. Evaluate your

procedure by addressing weaknesses and limitations and suggesting realistic improvements.

76

77

PPASTAST IB TIB TESTEST QQUESTIONSUESTIONS -- TTOPICOPIC 44

1. What is the formula for the compound formed by calcium and nitrogen?

A. CaN C. Ca2N3

B. Ca2N D. Ca3N2

2. What is the best description of the carbon-oxygen bond lengths in CO32-

?

A. One short and two long bonds

B. One long and two short bonds

C. Three bonds of the same length

D. Three bonds of different lengths

3. The boiling points of the hydrides of the group 6 elements are

shown at right.

i) Explain the trend in boiling points from to H2S to H2Te.

…………………………………………………………………….

…………………………………………………………………….

…………………………………………………………………….

……………………………………………………………………………………….

ii) Explain why the boiling point of water is higher than would be expected from the group trend.

……………………………………………………………………………………….

……………………………………………………………………………………….

……………………………………………………………………………………….

……………………………………………………………………………………….

4. i) State the shape of the electron distribution around the oxygen atom in the water molecule and state

the shape of the molecule.

………………………………………………………………………………………….

………………………………………………………………………………………….

ii) State and explain the value of the HOH bond angle.

…………………………………………………………………………………………..

…………………………………………………………………………………………..

400

300

200

100

0

Boiling point / K

H O H S H Se H Te2 2 2 2

78

5. Explain why the bonds in silicon tetrachloride, SiCl4, are polar, but the molecule is not.

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

6. The diagrams below represent the structures of iodine, sodium and sodium iodide.

A B C

a) i) Identify which of the structures (A, B and C) correspond to iodine, sodium and sodium iodide.

……………………………………………………………………………………….

ii) State the type of bonding in each structure.

……………………………………………………………………………………….

b) i) Sodium and sodium iodide can both conduct electricity when molten, but only sodium can con-

duct electricity when solid. Explain this difference in conductivity in terms of the structures of

sodium and sodium iodide.

……………………………………………………………………………………….

……………………………………………………………………………………….

……………………………………………………………………………………….

ii) Explain the high volatility of iodine compared to sodium and sodium iodide.

……………………………………………………………………………………….

……………………………………………………………………………………….

7. What is the best description of the carbon-oxygen bond lengths in CO32-

?

A. One short and two long bonds C. Three bonds of the same length

B. One long and two short bonds D. Three bonds of different lengths

79

8. What is the number of sigma (s) and pi (p) bonds and the hybridization of the carbon atom in

9. Element X is in group 2, and element Y in group 7, of the periodic table. Which ions will be present in the compound formed when X and Y react together?

A. X+ and Y

– C. X

+ and Y

2–

B. X2+

and Y–

D. X2–

and Y+

10. a) Draw the Lewis structure of methanoic acid, HCOOH.

b) In methanoic acid, predict the bond angle around the

(i) carbon atom. .....................................................................................................

(ii) oxygen atom bonded to the hydrogen atom. ...................................................

(c) State and explain the relationship between the length and strength of the bonds between the carbon

atom and the two oxygen atoms in methanoic acid.

......................................................................................................................................

......................................................................................................................................

......................................................................................................................................

11. Which of the following increase(s) for the bonding between carbon atoms in the sequence of molecules C2H6, C2H4 and C2H2?

I. Number of bonds

II. Length of bonds

III. Strength of bonding

A. I only C. III only

B. I and III only D. I, II and III

O

C HH O

Sigma Pi Hybridization

A. 4 1 sp2

B. 4 1 sp3

C. 3 2 sp3

D. 3 1 sp2

80

12. Which of the following contain a bond angle of 90° ?

I. PCl4+ II. PCl5 III. PCl6

-

A. I and II only C. II and III only

B. I and III only D. I, II and III

13. Which allotropes contain carbon atoms with sp2 hybridization?

I. Diamond II. Graphite III. C60 fullerene

A. I and II only C. II and III only

B. I and III only D. I, II and III

14. State the complete electronic configuration of bromine, Br and the iron(III) ion, Fe3+

.

Br: ………………………………………………………………………………….

Fe3+

: ………………………………………………………………………………….

15. The boiling points of the hydrides of group 6 elements increase in the order

H2S < H2Se < H2Te < H2O.

Explain the trend in the boiling points in terms of bonding.

................................................................................................................................................

................................................................................................................................................

................................................................................................................................................

16. Identify which of the compounds butane, chloroethane, propanone and propan-1-ol are

(i) insoluble in water and give your reasoning.

................................................................................................................................................

................................................................................................................................................

................................................................................................................................................

................................................................................................................................................

(ii) water soluble and give your reasoning.

................................................................................................................................................

................................................................................................................................................

................................................................................................................................................

81

17. i) Draw the Lewis structures for carbon monoxide, carbon dioxide and the carbonate ion.

................................................................................................................................................

................................................................................................................................................

................................................................................................................................................

................................................................................................................................................

ii) Identify the species with the longest carbon-oxygen bond and explain your answer.

................................................................................................................................................

................................................................................................................................................

................................................................................................................................................

................................................................................................................................................

18. Hydrazoic acid, N3H, can be represented by two possible Lewis structures in which the atoms can be ar-

ranged as NNNH.

(i) Draw the two possible Lewis structures of N3H.

................................................................................................................................................

................................................................................................................................................

................................................................................................................................................

................................................................................................................................................

(ii) Predict the N––N––N and H––N––N bond angles in each case and give your reasoning.

................................................................................................................................................

................................................................................................................................................

................................................................................................................................................

(iii) Predict the hybridization of the N atom bonded to the hydrogen atom in each case.

................................................................................................................................................

................................................................................................................................................

19. Based on electronegativity values, which bond is the most polar?

A. B–C C. N–O

B. C–O D. O–F

82

20. What is the Lewis (electron dot) structure for sulfur dioxide?

A. C.

B. D.

21. Which substance is most soluble in water (in mol dm−3

) at 298 K?

A. CH3CH3 C. CH3CH2OH

B. CH3OCH3 D. CH3CH2CH2CH2OH

22. i) Draw Lewis (electron dot) structures for CO2 and H2S showing all valence electrons.

ii) State the shape of each molecule and explain your answer in terms of VSEPR theory.

CO2 .............................................................................................................................

H2S .............................................................................................................................

iii) State and explain whether each molecule is polar or non-polar.

.....................................................................................................................................

.....................................................................................................................................

23. Identify the strongest type of intermolecular force in each of the following compounds.

CH3Cl ...................................................................................................................................

CH4 .......................................................................................................................................

CH3OH .................................................................................................................................

24. Which of the following species is (are) planar (has (have) all the atoms in one plane)?

I. CO32-

II. NO3-

III. SO32-

A. I only C. I and II only

B. II only D. II and III only

O S O O S O

O S O O S O

83

25. What is the molecular shape and the hybridization of the nitrogen atom in NH3?

26. Which statement about sigma and pi bonds is correct?

A. Sigma bonds are formed only by s orbitals and pi bonds are formed only by p orbitals.

B. Sigma bonds are formed only by p orbitals and pi bonds are formed only by s orbitals.

C. Sigma bonds are formed by either s or p orbitals, pi bonds are formed only by p orbitals.

D. Sigma and pi bonds are formed by either s or p orbitals.

27. Draw Lewis (electron dot) structures for the following ions.

NO2– NO2+

Determine and explain the shape of each ion.

……………………………………………………………………………………………………

……………………………………………………………………………………………………

……………………………………………………………………………………………………

……………………………………………………………………………………………………

……………………………………………………………………………………………………

28. (i) List the following substances in order of increasing boiling point (lowest first).

CH3CHO C2H6 CH3COOH C2H5OH

………………………………………………………………………………………………

(ii) State whether each compound is polar or non-polar, and explain the order of boiling points in (c)(i).

………………………………………………………………………………………………

………………………………………………………………………………………………

………………………………………………………………………………………………

Molecular shape Hybridization

A. tetrahedral sp3

B. trigonal planar sp2

C. trigonal pyramidal sp2

D. trigonal pyramidal sp3

84

29. According to VSEPR theory, repulsion between electron pairs in a valence shell decreases in the order

A. lone pair-lone pair > lone pair-bond pair > bond pair-bond pair.

B. bond pair-bond pair > lone pair-bond pair > lone pair-lone pair.

C. lone pair-lone pair > bond pair-bond pair > bond pair-lone pair.

D. bond pair-bond pair > lone pair-lone pair > lone pair-bond pair.

30. Which molecule is linear?

A. SO2 C. H2S

B. CO2 D. Cl2O

31. Why is the boiling point of PH3 lower than that of NH3?

A. PH3 is non-polar whereas NH3 is polar.

B. PH3 is not hydrogen bonded whereas NH3 is hydrogen bonded.

C. Van der Waals’ forces are weaker in PH3 than in NH3

D. The molar mass of PH3 is greater than that of NH3

32. Which molecule is non-polar?

A. H2CO C. NF3

B. SO3 D. CHCl3

33. a) An important compound of nitrogen is ammonia, NH3. The chemistry of ammonia is influenced by its

polarity and its ability to form hydrogen bonds. Polarity can be explained in terms of electronegativity.

i) Explain the term electronegativity.

……………………………………………………………………………………

……………………………………………………………………………………

ii) Draw a diagram to show hydrogen bonding between two molecules of NH3.

The diagram should include any dipoles and/or lone pairs of electrons

……………………………………………………………………………………

……………………………………………………………………………………

……………………………………………………………………………………

85

iii) State the H–N–H bond angle in an ammonia molecule.

………………………………………………………………………………………

iv) Explain why the ammonia molecule is polar.

……………………………………………………………………………………

……………………………………………………………………………………

b) Ammonia reacts with hydrogen ions forming ammonium ions, NH4+.

i) State the H–N–H bond angle in an ammonium ion.

……………………………………………………………………………………

ii) Explain why the H–N–H bond angle of NH3 is different from the H–N–H bond angle of NH4+;

referring to both species in your answer.

……………………………………………………………………………………

……………………………………………………………………………………

……………………………………………………………………………………

……………………………………………………………………………………

34. NO3- is trigonal planar and NH3 is trigonal pyramidal. What is the type of hybridization of N in each of these

species?

35. Consider the following statements. Which statements are correct?

I. All carbon-oxygen bond lengths are equal in CO32-

.

II. All carbon-oxygen bond lengths are equal in CH3COOH.

III. All carbon-oxygen bond lengths are equal in CH3COO−.

A. I and II only C. II and III only

B. I and III only D. I, II and III

N in NO3- N in NH3

A. sp

2 sp

3

B. sp

2 sp

2

C. sp

3 sp

2

D. sp

3 sp

3

86

36. In 1954 Linus Pauling was awarded the Chemistry Nobel Prize for his work on the nature of the chemical

bond. Covalent bonds are one example of intramolecular bonding. Explain the formation of the following.

(i) σ bonding

……………………………………………………………………………………………

……………………………………………………………………………………………

(ii) π bonding

……………………………………………………………………………………………

……………………………………………………………………………………………

(iii) double bonds

……………………………………………………………………………………………

……………………………………………………………………………………………

(iv) triple bonds

……………………………………………………………………………………………

……………………………………………………………………………………………

37. Atomic orbitals can mix by hybridization to form new orbitals for bonding. Identify the type of hybridization

present in each of the three following molecules. Deduce and explain their shapes.

(i) OF2

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

(ii) H2CO

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

(iii) C2H2

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

87

38. Three scientists shared the Chemistry Nobel Prize in 1996 for the discovery of fullerenes.

Fullerenes, like diamond and graphite, are allotropes of the element carbon.

i) State the structures of and the bonding in diamond and graphite.

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

ii) Compare and explain the hardness and electrical conductivity of diamond and graphite.

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

iii) Predict and explain how the hardness and electrical conductivity of C60 fullerene would compare with

that of diamond and graphite.

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

39. What happens when sodium and oxygen combine together?

A. Each sodium atom gains one electron.

B. Each sodium atom loses one electron.

C. Each oxygen atom gains one electron.

D. Each oxygen atom loses one electron.

40. Which statement is correct about two elements whose atoms form a covalent bond with each other?

A. The elements are metals.

B. The elements are non-metals.

C. The elements have very low electronegativity values.

D. The elements have very different electronegativity values.

88

41. In ethanol, C2H5OH (l), there are covalent bonds, hydrogen bonds and van der Waals’ forces. Which bonds

or forces are broken when ethanol is vaporized?

A. only hydrogen bonds C. covalent bonds and van der Waals’ forces

B. covalent bonds and hydrogen bonds D. hydrogen bonds and van der Waals’ forces

42. Which substance has the lowest electrical conductivity?

A. Cu(s) C. H2(g)

B. Hg(l) D. LiOH(aq)

43. The letters W, X, Y and Z represent four consecutive elements in the periodic table.

The number of electrons in the highest occupied energy levels are:

W: 3, X: 4, Y: 5, Z: 6

Write the formula for

(i) an ionic compound formed from W and Y, showing the charges.

……………………………………………………………………………………………

……………………………………………………………………………………………

(ii) a covalent compound containing X and Z.

……………………………………………………………………………………………

……………………………………………………………………………………………

44. State the type of bonding in the compound SiCl4. Draw the Lewis structure for this compound.

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

45. Outline the principles of the valence shell electron pair repulsion (VSEPR) theory.

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

89

46. i) Use the VSEPR theory to predict and explain the shape and the bond angle of each of the molecules

SCl2 and C2Cl2

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

ii) Deduce whether or not each molecule is polar, giving a reason for your answer.

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

47. Which statement best describes the attraction present in metallic bonding?

A. the attraction between nuclei and electrons C. the attraction between positive ions and negative ions

B. the attraction between positive ions and electrons D. the attraction between protons and electrons

48. Which statement is correct about multiple bonding between carbon atoms?

A. Double bonds are formed by two π bonds. C. π bonds are formed by overlap between s orbitals.

B. Double bonds are weaker than single bonds. D. π bonds are weaker than sigma bonds.

49. Which statements are correct about diamond, graphite and a C60 fullerene?

I. The poorest electrical conductor of the three is diamond.

II. The atoms in graphite and C60 fullerene are sp2 hybridized.

III. The atoms in diamond and C60 fullerene are arranged in hexagons.

A. I and II only C. II and III only

B. I and III only D. I, II and III

90

50. For the following compounds

PCl3 , PCl5 , POCl3

i) Draw a Lewis structure for each molecule in the gas phase. (Show all non-bonding electron pairs.)

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

ii) State the shape of each molecule and predict the bond angles.

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

iii) Deduce whether or not each molecule is polar, giving a reason for your answer.

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

51. i) Explain the meaning of the term hybridization.

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

91

ii) Discuss the bonding in the molecule CH3CHCH2 with reference to

· the formation of σ and π bonds

· the length and strength of the carbon-carbon bonds

· the types of hybridization shown by the carbon atoms

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

.....................................................................................................................................

52. How do bond length and bond strength change as the number of bonds between two atoms increases?

53. Which of the following is true for CO2?

54. The molar masses of C2H6, CH3OH and CH3F are very similar. How do their boiling points compare?

A. C2H6 < CH3OH < CH3F

B. CH3F < CH3OH < C2H6

C. CH3OH < CH3F < C2H6

D. C2H6 < CH3F < CH3OH

55. Which is the correct description of polarity in F2 and HF molecules?

A. Both molecules contain a polar bond. C. Both molecules are polar.

B. Neither molecule contains a polar bond. D. Only one of the molecules is polar.

Bond length Bond strength

A. increases increases

B. increases decreases

C. decreases increases

D. decreases decreases

C==O bond CO2 molecule

A. polar non-polar

B. non-polar polar

C. polar polar

D. non-polar non-polar

92

56. Which types of bonding are present in CH3CHO in the liquid state?

I. Single covalent bonding

II. Double covalent bonding

III. Hydrogen bonding

A. I and II only C. II and III only

B. I and III only D. I, II and III

57. Which statement(s) is/are generally true about the melting points of substances?

I. Melting points are higher for compounds containing ions than for compounds containing mole-

cules.

II. A compound with a low melting point is less volatile than a compound with a high melting point.

III. The melting point of a compound is decreased by the presence of impurities.

A. I only C. II and III only

B. I and III only D. I, II and III

93

Organic 10.1 – Introduction

1. Describe the features of a homologous series.

2. Predict and explain the trends in boiling points of members of homologous series

3. Distinguish between empirical, molecular, and structural formulas.

4. Describe structural isomers as compounds with the same molecular formula but with different arrange-

ments of atoms.

5. Deduce structures formulas for the isomers of the non-cyclic alkanes up to C6.

6. Apply the IUPAC rules for naming the isomers of the non-cyclic alkanes up to C6.

7. Deduce structural formulas for the isomers of the straight-chain alkenes up to C6.

8. Apply the IUPAC rules for naming the isomers of the straight-chain alkenes up to C6.

9. Deduce structural formulas for compounds containing up to six carbon atoms with one of the following

functional groups: alcohol, aldehyde, ketone, carboxylic acid and halide.

10. Apply the IUPAC rules for naming compounds containing up to six carbon atoms with one of the following

functional groups: alcohol, aldehyde, ketone, carboxylic acid and halide.

11. Identify the following functional groups when present in structural formulas: amino (NH2), benzene ring,

and esters (RCOOR).

12. Identify primary, secondary and tertiary carbon atoms in alcohols and halogenoalkanes.

13. Discuss the volatility and solubility in water of the compounds containing the functional groups listed in

10.1.9.

20.1 – Introduction (AHL) 1. Deduce structural formulas for compounds containing up to six carbon atoms with one of the following

functional groups: amine, amide, ester and nitrile.

2. Apply the IUPAC rules for naming compounds containing up to six carbon atoms with one of the following

functional groups: amine, amide, ester and nitrile.

10.2 – Alkanes 1. Explain the low reactivity of alkanes in terms of bond enthalpies and bond polarity.

2. Describe, using equations, the complete and incomplete combustion of alkanes.

3. Describe, using equations, the reactions of methane and ethane with chlorine and bromine.

4. Explain the reactions of methane and ethane with chlorine and bromine in terms of a free-radical mecha-

nism.

20.2 – Nucleophilic Substitution Reactions (AHL) 1. Explain why the hydroxide ion is a better nucleophile than water.

2. Describe and explain how the rate of nucleophilic substitution in halogenoalkanes by the hydroxide ion

depends on the identity of the halogen.

3. Describe and explain how the rate of nucleophilic substitution in halogenoalkanes by the hydroxide ion

depends on whether the halogenoalkane is primary, secondary or tertiary.

4. Describe, using equations, the substitution reactions of halogenoalkanes with ammonia and potassium

cyanide.

5. Explain the reactions of primary halogenoalkanes with ammonia and potassium cyanide in terms of the

SN2 mechanism.

6. Describe, using equations, the reduction of nitriles using hydrogen and a nickel catalyst.

94

10.3 – Alkenes 1. Describe, using equations, the reactions of alkenes with hydrogen and halogens.

2. Describe, using equations, the reactions of symmetrical alkenes with hydrogen halides and water.

3. Distinguish between alkanes and alkenes using bromine and water.

4. Outline the polymerization of alkenes.

5. Outline the economic importance of the reactions of alkenes.

20.3 – Elimination Reactions (AHL) 1. Describe, using equations, the elimination of HBr from bromoalkanes.

2. Describe and explain the mechanism for the elimination of HBr from bromoalkanes.

10.4 – Alcohols 1. Describe, using equations, the complete combustion of alcohols.

2. Describe, using equations, the oxidation reactions of alcohols.

3. Determine the products formed by the oxidation of primary and secondary alcohols.

20.4 – Condensation Reactions (AHL) 1. Describe, using equations, the reactions of alcohols with carboxylic acids to form esters, and state the

uses of esters.

2. Describe, using equations, the reactions of amines with carboxylic acids.

3. Deduce the structures of the polymers formed in the reactions of alcohols with carboxylic acids.

4. Deduce the structures of the polymers formed in the reactions of amines with carboxylic acids.

5. Outline the economic importance of condensation reactions.

10.5 – Halogenoalkanes 1. Describe, using equations, the substitution reactions of halogenoalkanes with sodium hydroxide.

2. Explain the substitution reactions of halogenoalkanes with sodium hydroxide in terms of SN1 and SN2

mechanisms.

10.6 – Reaction Pathways 1. Deduce reaction pathways given the starting materials and the product.

20.6 – Stereoisomerism (AHL) 1. Describe stereoisomers as compounds with the same structural formulas but with different arrange-

ments of atoms in space.

2. Describe and explain geometrical isomerism in non-cyclic alkenes.

3. Describe and explain geometrical isomerism in C3 and C4 cycloalkanes.

4. Explain the differences in the physical and chemical properties of geometrical isomers.

5. Describe and explain optical isomerism in simple organic molecules.

6. Outline the use of a polarimeter in distinguishing between optical isomers.

7. Compare the physical and chemical properties of enantiomers.

95

Intro to Organic Chemistry Name__________________________________

1. Write the structural formulas and names for all of the isomers of pentane.

2. Which of the isomers you showed above would have the lowest boiling point? Explain.

3. Write the structural formula and name for all members of the homologous series containing pentane that have lower boiling points.

4. Consider the combustion of octane.

a. Write a balanced equation.

b. Is the combustion of hydrocarbons an exothermic or endothermic process?

c. What are the products of incomplete combustion of octane?

5. Consider the following amino acids.

a. Use the data booklet to name amino acids A and B.

b. Identify the chiral carbon in each amino acid by circling it.

c. Draw the enantiomer (optical isomer) for amino acid A.

d. What device would allow one to identify the presence of optical isomers (enantiomers) in a sam-ple?

NH2

CH

CCH

3

O

OH NH2

CH

CCH

2

O

OH

SH

A B

96

6. Draw molecules containing two carbon atoms, each with a different one of the following functional

groups: alcohol, aldehyde, ketone, carboxylic acid and halide. Name each molecule.

7. Draw one large molecule that contains ALL of the following functional groups: amine, amide, ester and nitrile.

8. Explain the differences between each of the following pairs of functional groups that are often mistaken for

one another:

a. Amines and amides

b. Alcohols and carboxylic acids

c. Esters and ketones

d. Aldehydes and ketones

e. Amines and nitriles

9. Explain what is meant by the term alkene and how that is different than an alkane.

10. Why is it necessary to include the prefix “di” in some names, such as dichloroethane or pent-1,3-diene?

97

Organic Naming Practice Name__________________________________

Write the name of each of the following organic compounds:

1) 2)

Is this molecule cis or trans?

3) 4) CH3-CH2-CH3

5) 6)

Is this molecule cis or trans?

Write the structures for the following organic molecules:

7) butane 8) 1-butene

9) 3-pentene 10) methane

11) 2-butyne 12) cyclopentane

For each molecule in 7-12, give the chemical formula (ex. C2H6)

7) __________ 8) __________ 9) __________

10) __________ 11) __________ 12) __________

98

13. First identify the function group(s) present in each molecule. Then name the following compounds:

a) b) c)

d) e) f)

g) h) CH3CH2CHClCHClCH3 i) CH3CH(CH3)CH(CH3)CH2CH3

14. Draw structures, or write structural formulas, for:

a) 3-fluoro pentane Bonus Challenge: d) 1-hydroxy but-2-ene

b) propanamide e) trans 1,2 dibromoethene

c) ethanoic acid

15. Write the structural formulas and names for all of the possible isomers of C5H12.

CO

O

CH

2

C

H2

CH

2

CH3

HO

CH3

C

H2

CH

2

C

H2

CH

2

C

H2

HCH

3

CH

CH

CH3

CH3

C

H2

CH

2

CH

O

CH3

C

H2

NH2

CH3

C

H2

CH

2

CCH

2

C

H2

CH

2

CH3

O

CH3

C

H2

CH C

H2

CH3

CH3

99

Investigating Odiferous Isomers – Olfactory fun for you

and your friends! Your tongue can only detect four distinct flavors – salty, sweet, bitter, and sour. The other “flavors” are

actually an artifact of olfaction. “Olfaction” is the technical term for your sense of smell. In humans, the

olfactory system consists of the olfactory epithelium, located in your nasal passages, and the olfactory

nerve, part of the central nervous system. The olfactory epithelium in your nasal passages contains re-

ceptors responsible for binding the odorant molecules. This binding then transmits a signal through the

olfactory nerve, where it is interpreted by your brain. Olfactory receptors are different than other pro-

tein-based receptors in the body, in that they do not necessarily bind to only one specific ligand. Rather,

olfactory receptors show affinity for a range of odorant molecules, although the interaction between a

specific odorant and a receptor results in the transmission of a specific impulse through the olfactory

nerve. The interpretation of this specific impulse varies for the multitude of “smells” we are capable of

sensing.

In order to smell something, first some of the molecules must evaporate into the air. These “odorants”

are the volatile components released by a particular substance. These vapors travel through the air,

where they are free to be inhaled and exposed to the olfactory epithelium. Once bound to the olfactory

receptors, these receptors experience a specific change in shape. This change in shape transmits an im-

pulse, via the olfactory nerve, to the brain, where it is interpreted as a specific scent.

100

Scents are unique to specific types of compounds. For instance, esters tend to be interpreted by your brain as

having a sweet or sometimes fruity aroma. Carboxylic acids often smell rancid or vinegary (vinegar, itself, is a

carboxylic acid). Amines often smell fishy. The molecular weight, as well as the presence of aromatic rings (not

aromatic as in fragrant, but aromatic as in containing a benzene ring), can also affect the smell of compounds. For

instance, aldehydes that are low in molecular weight have characteristically unpleasant, sometimes rotten smells,

whereas high molecular weight aldehydes are often pleasant and fruity smelling. What is important to recognize

is that the presence of these functional groups is what gives the odorants their characteristic shape, and thus

what affects the smell that your brain interprets.

What makes olfaction so interesting is that many times, despite the fact that two molecules may have identical

molecular formulas, because the atoms can be arranged differently, these two molecules can have very different

structures, and thus very different smells. This is called “isomerism,” which refers to a situation where two sub-

stances have identical molecular formulas but a different arrangement of atoms in space. For instance, butanoic

acid (C4H8O2) has a very unpleasant, rancid odor and is common in spoiled food. However, ethyl ethanoate (also

C4H8O2) has a very pleasant, sweet smell and is common in many fruits. This is because the atoms, though com-

mon to each substance, are arranged differently. Butanoic acid has a structural formula of CH3CH2CH2COOH,

while ethyl ethanoate is CH3CH2COOCH3. Butanoic acid is classified as a carboxylic acid, whereas ethyl ethanoate

is classified as an ester. Different shapes, different smells, though not necessarily different atoms!

Most of the compounds which give flavor and odor to foods, flowers, and perfumes are organic compounds. Some-

times it is just one compound, but often it is a mixture. If you think about it, you can distinguish hundreds of sub-

stances by their smell. Your ability to detect these molecules has to do with the shape of their hydrocarbon skele-

ton and the position of their functional groups.

In this investigation, you will be exposed to 13 distinct odorants. First, you will examine the molecular formula

and identify distinct functional groups. Then, you will make a prediction regarding the smell you expect the odor-

ant to have. Finally, you will actually smell the odorant and test your prediction.

In general:

ester: fruity, sweet

amine: fishy

aldehyde: depends on relative molecular weight and/or presence of aromatic ring (often imparts a sweet aro-

ma)

carboxylic acid: tart, rancid, putrid

ketone: sweet, medicinal

alcohol: depends on molecular weight and/or presence of aromatic ring

These are only generalizations, as actual smell interpretation depends on the combination of functional groups,

their position in the molecule, and the relative size of the molecule. Olfaction is indeed a very complex sense.

The following descriptive terms may help you as you complete this investigation:

animal (musk, sweaty) fruity (various fruits)

balsamic (heavy, sweet odors like cocoa, vanilla, green (odor of fresh cut grass)

cinnamon) medicinal (disinfectants, Lysol, ben-gay)

camphoraceous (camphor) metallic (brass, steel)

citrus (sweet, odor of citrus fruits) minty (peppermint-like odor)

earthy (humus-like, humid earth) mossy (forests and seaweed)

fatty (smelling of animal fat) spicy (various spices)

floral (various flowers) woody (cedarwood, sandlewood)

101

Use the following chart during this investigation:

Vial Molecular Formu-la and Name

Functional Group(s)

Structural Formula Predicted Smell

Actual Smell

A menthol

OH

CH

CH CH

2

CH

C

H2

C

H2

CH3

OH

CH

CH3

CH3

B

NH2

CH3

C

H2

CH

2

C

H2

CH

2

CH CH

3

NH2

C

O

O

CH3

CO

C

H2

CH

2

C

H2

CH

2

CH3

O

D

O

O

CH

3

CO

C

H2

CH3

O

E methyl salicylate

O

O

OH

CH

C

CCH

CH

CH

CO

CH3

O

OH

F

O

OH

CH3

C

H2

CH

2

C

O

OH

G

OH

O

CH3

CH

2

C

H2

CH

2

C

H2

CH

2

COH

O

102

H methyl benzoate

O

O

CH

CH

CCH

CH

CH

CO

CH3

O

I benzaldehyde

CH

O

CH

CH

CCH

CH

CH

CH

O

J cinnamaldehyde

CH

O

CH

CH

CCH

CH

CH

CH

CH

CH

O

K borneol

OH

CH

CH2

CH

C

C

CH2

CH

2

CH3

CH3

CH3OH

L camphor

O

C

CH2

CH

C

C

CH2

CH

2

CH3

CH3

CH3O

M vanillin

CH

O

OH

O

CH3

CH

O

C

CH

CH

CH

CCOH

O

CH3

103

Analysis Questions

1. When you smelled each sample, were you smelling the entire molecule (as shown by the structur-

al formula), or were you instead smelling individual atoms that broke away from the original

molecules? Explain your answer in terms of intermolecular and intramolecular forces.

2. Do molecules with the same molecular formula always smell the same? Give at least one specific

example.

3. Identify at least one set of isomers (two molecules) from the odorants A – M. Write the name,

molecular formula, and functional groups present in each isomer.

4. Describe the difference in smell between the two isomers you selected in number three.

5. Identify two compounds that had similar smells but not necessarily the same chemical formula or

functional groups. Describe similarities between the two molecules (i.e. shape, molecular mass,

etc.)

6. How close were your predictions to the actual smells? Were there any that you were surprised

by?

7. Key concept: What accounts for the different smells of two molecules with identical molecular

formulas?

104

105

Intro to Organic Reactions Name__________________________________

1. Describe the key characteristics of a combustion reaction and give the two possible products of an incomplete combustion.

2. Besides combustion, what is the only other type of reaction that an alkane can undergo? What condi-tions are necessary for this type of reaction to occur?

3. Define the terms:

a. Free radical

b. Homolytic fission

c. Mechanism

d. Nucleophile

e. Monomer

f. Polymer

g. Reflux

h. Catalyst

4. Name and describe each of the three steps that are necessary in the mechanism of a substitution re-action in order for an alkane to be turned in to a halogenoalkane.

106

5. Write the substitution reaction that would occur between propane and bromine.

6. Identify the necessary reactants for a nucleophilic substitution reaction and identify the class of

compound that would be formed by a hydroxide nucleophile.

7. Each type of nucleophilic substitution involves two steps in its mechanism: nucleophilic attack and

the halogen leaving. List the order of these two steps for SN1 and SN2:

SN1 Steps— 1. _______________________________ 2. _______________________________

SN2 Steps— 1. _______________________________ 2. _______________________________

8. Explain what each symbol in the term SN1 stands for.

9. For each of the following classes of compounds, identify whether it would undergo SN1 or SN2.

Primary halogenoalkane Secondary halogenoalkane Tertiary halogenoalkane

10. Why does an SN1 mechanism occur faster than an SN2 mechanism for most molecules?

11. Complete the blanks below for the following addition reactions. Be sure to add catalysts over the ar-

row when necessary.

a. Ethene + Hydrogen __________________________

b. Ethene + ____________________ Ethanol

c. Ethene + Hydrochloric acid __________________________

d. Ethene + ____________________ Dichloroethane

e. LOTS of ethane __________________________

12. For each of the following classes of compounds, identify the product(s) (if any) of its oxidation.

Primary alcohol Secondary alcohol Tertiary alcohol

107

PPASTAST IB TIB TESTEST QQUESTIONSUESTIONS -- TTOPICOPIC 1010

1. Which of the structures below is an aldehyde?

A.

B.

C.

D.

2. What product results from the reaction of

CH2==CH2 with Br2?

A. CHBrCHBr

B. CH2CHBr

C. CH3CH2Br

D. CH2BrCH2Br

3. What is the final product formed when

CH3CH2OH is refluxed with acidified potassi-

um dichromate(VI)?

A. CH3CHO

B. CH2==CH2

C. CH3COOH

D. HCOOCH3

Name ________________________________

4. Which of the substances below is least soluble

in water?

A. CH2OHCHOHCH2OH

B.

C.

D.

5. Which substance(s) could be formed during

the incomplete combustion of a hydrocarbon?

I. Carbon

II. Hydrogen

III. Carbon monoxide

A. I only

B. I and II only

C. I and III only

D. II and III only

6. Which formulas represent butane or its iso-

mer?

I. CH3(CH2)2CH3

II. CH3CH(CH3)CH3

III. (CH3)3CH

A. I and II only

B. I and III only

C. II and III only

D. I, II and III

CH CH CH

O

3 2

CH CCH

O

3 3

CH CH COH

O

3 2

CH COCH

O

3 3

CH CCH

O

3 3

CH CH COH

O

3 2

CH COCH

O

3 3

108

7. Which compound can exist as optical isomers?

A. CH3CHBrCH3

B. CH2BrCHBrCH3

C. CH2BrCHBrCH2Br

D. CHBr2CHBrCHBr2

8. Which statement about the reactions of

halogenoalkanes with aqueous sodium hy-

droxide is correct?

A. Primary halogenoalkanes react mainly

by an SN1 mechanism.

B. Chloroalkanes react faster than iodoal-

kanes.

C. Tertiary halogenoalkanes react faster

than primary halogenoalkanes.

D. The rate of an SN1 reaction depends on

the concentration of aqueous sodium

hydroxide.

9. Which type of compound must contain a mini-

mum of three carbon atoms?

A. An aldehyde

B. A carboxylic acid

C. An ester

D. A ketone

10. What is the IUPAC name for CH3CH2CH

(CH3)2?

A. 1,1-dimethylpropane

B. 2-methylbutane

C. isopentane

D. ethyldimethylmethane

11. Which product is formed by the reaction be-

tween CH2CH2 and HBr?

A. CH3CH2Br

B. CH2CHBr

C. BrCHCHBr

D. CH3CHBr2

12. Which reaction(s) involve(s) the formation of

a positive ion?

I. CH3CH2CH2Br + OH–

II. (CH3)3CBr + OH–

A. I only

B. II only

C. Both I and II

D. Neither I nor II

13. Consider the following compounds.

I. CH3CH2CH(OH)CH3

II. CH3CH(CH3)CH2OH

III. (CH3)3COH

The compounds are treated separately with

acidified potassium dichromate(VI) solution.

Which will produce a colour change from or-

ange to green?

A. I and II only

B. I and III only

C. II and III only

D. I, II and III

109

14. Which compound reacts most rapidly by a SN1

mechanism?

A. (CH3)3CCl

B. CH3CH2CH2CH2Br

C. (CH3)3CBr

D. CH3CH2CH2CH2Cl

15. Which compound is a member of the same

homologous series as 1-chloropropane?

A. 1-chloropropene

B. 1-chlorobutane

C. 1-bromopropane

D. 1,1-dichloropropane

16. What is the organic product of the reaction

between ethanol and ethanoic acid?

A. CH3CHO

B. CH3COOCH3

C. CH3CH2COOCH3

D. CH3COOCH2CH3

17. Which statement is correct about the reaction

between methane and chlorine?

A. It involves heterolytic fission and Cl–

ions.

B. It involves heterolytic fission and Cl

radicals.

C. It involves homolytic fission and Cl–

ions.

D. It involves homolytic fission and Cl

radicals.

18. Which compound is converted to butanal by

acidified potassium dichromate(VI) solution?

A. butan-1-ol

B. butan-2-ol

C. butanone

D. butanoic acid

19. Which formula represents a tertiary alcohol?

20. Which reaction type is typical for halogenoal-

kanes?

A. nucleophilic substitution

B. electrophilic substitution

C. electrophilic addition

D. nucleophilic addition

21. Which substance is not readily oxidized by

acidified potassium dichromate(VI) solution?

A. propan-1-ol

B. propan-2-ol

C. propanal

D. propanone

CHCH 3 CH 2 CH3

CH2OH

A. B. CHCH3 CH2

CH3

CH2 OH

C.

CCH3 CH2

CH3

CH3

OH

D.

CHCH3 CH CH3

CH 3

OH

110

22. Which are characteristics typical of a free radi-

cal?

I. It has a lone pair of electrons.

II. It can be formed by the homolytic

fission of a covalent bond.

III. It is uncharged.

A. I and II only

B. I and III only

C. II and III only

D. I, II and III

23. Which pair of compounds can be used to pre-

pare CH3COOCH3?

A. Ethanol and methanoic acid

B. Methanol and ethanoic acid

C. Ethanol and ethanoic acid

D. Methanol and methanoic acid

24. What is the reaction type when (CH3)3CBr re-

acts with aqueous sodium hydroxide to form

(CH3)3COH and NaBr?

A. Addition

B. Elimination

C. SN1

D. SN2

25. Which species is a free radical?

A. •CH3

B. +CH3

C. –CH3

D. :CH3

26. Nylon is a condensation polymer made up of

hexanedioic acid and 1,6-diaminohexane.

Which type of linkage is present in nylon?

A. Amide

B. Ester

C. Amine

D. Carboxyl

27. How many chiral carbon atoms are present in a

molecule of glucose?

A. 1

B. 2

C. 3

D. 4

28. Which amino acid can exist as optical isomers?

A.

B.

C.

D.

111

29. The following is a three-dimensional representation of an organic molecule.

Which statement is correct?

A. The correct IUPAC name of the molecule is 2-methylpentane.

B. All the bond angles will be approximately 90°.

C. One isomer of this molecule is pentane.

D. The boiling point of this compound would be higher than that of pentane.

30. What is the product of the following reaction?

CH3CH2CH2CN + H2

A. CH3CH2CH2NH2

B. CH3CH2CH2CH3

C. CH3CH2CH2CH2CH3

D. CH3CH2CH2CH2NH2

31. (i) A compound D has the molecular formula C2H4O2 and is obtained from a reaction between methano-

ic acid and methanol. Write an equation for this reaction and state the name of D.

…………………………………………………………………………………………..

…………………………………………………………………………………………..

………………………………………………………………………………………….. (3)

(ii) A second compound, E, has the same molecular formula as D and has acidic properties.

State the name of compound E.

………………………………………………………………………………………….. (1)

Ni

112

32. The first synthetic thread was made from a polyester. A section of the polyester is drawn below:

–––CH2COO–––CH2COO–––CH2COO–––

(i) Give the structural formula of the monomer (containing two functional groups) that could be used to

make this polyester and state the names of the two functional groups.

…………………………………………………………………………………………..

…………………………………………………………………………………………..

…………………………………………………………………………………………..

…………………………………………………………………………………………..

………………………………………………………………………………………….. (3)

(ii) State, giving a reason, whether this polyester is made by a condensation reaction or an addition reac-

tion.

…………………………………………………………………………………………..

………………………………………………………………………………………….. (2)

33. Hexanedioic acid and 1,6-diaminohexane react together to form a synthetic polymer. There are many natu-

ral polymers, some of the most familiar being proteins formed from 2-amino acids.

(i) Give the structural formula of each monomer in the synthetic polymer.

……………………………………………………………………………………………

……………………………………………………………………………………………

……………………………………………………………………………………………

…………………………………………………………………………………………… (2)

(ii) State the type of polymerization reaction that occurs between these two monomers and identify the

structural feature needed in the monomers.

…………………………………………………………………………………………… (2)

(iii) Draw the structure of and state the type of linkage formed in this polymer, and identify the other

product of this polymerization reaction.

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……………………………………………………………………………………………

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113

(iv) The structures of some 2-amino acids are shown in Table 20 of the Data Booklet. Using alanine as an

example, explain what is meant by the term optical activity, identify the structural feature that needs

to be present and illustrate your answer with suitable diagrams of both isomers.

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……………………………………………………………………………………………

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(v) Identify a 2-amino acid from Table 20 which does not show optical activity.

…………………………………………………………………………………………… (1)

(vi) Polyesters are formed in a similar polymerization reaction to proteins. Their monomers are esters.

State one use of esters and identify the two compounds that react together to form the ester ethyl

methanoate.

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……………………………………………………………………………………………

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34. The compound, 2-bromobutane, CH3CHBrCH2CH3, can react with sodium hydroxide to form

compounds F, G and H.

Compound F, C4H10O, exists as a pair of optical isomers. Compounds G and H, C4H8, are structural iso-

mers, and compound H exists as a pair of geometrical isomers.

(i) Draw the structures of the two optical isomers of F.

……………………………………………………………………………………………

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…………………………………………………………………………………………… (2)

(ii) Outline the use of a polarimeter in distinguishing between the optical isomers.

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…………………………………………………………………………………………… (2)

114

(iii) Draw diagrams to show the shapes of the two geometrical isomers of H.

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35. (a) There are geometrical isomers of the cyclic compound C4H6Cl2. Draw the structural formula of two

isomers and explain why these two isomers exist.

cis-isomer trans-isomer

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.................................................................................................................................... (3)

(b) (i) Draw the structural formulas of two isomers of but-2-ene-1,4-dioic acid.

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……………………………………………………………………………………………

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…………………………………………………………………………………………… (2)

(ii) State and explain which isomer will have a lower melting point.

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……………………………………………………………………………………………

……………………………………………………………………………………………

…………………………………………………………………………………………… (2)

115

(c) Consider the following compounds:

1-chloropentane, 2-chloropentane, 3-chloropentane

(i) Identify the compound which exhibits optical isomerism and draw the structures of the two

isomers.

......................................................................................................................... (3)

(ii) Describe how these two isomers can be distinguished experimentally.

.........................................................................................................................

......................................................................................................................... (1)

36. (i) Ethanoic acid reacts with ethanol in the presence of concentrated sulfuric acid and heat. Identify the

type of reaction that takes place. Write an equation for the reaction, name the organic product formed

and draw its structure.

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………………………………………………………………………………………….. (4)

(ii) State and explain the role of sulfuric acid in this reaction.

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………………………………………………………………………………………….. (2)

(iii) State one major commercial use of the organic product from this type of reaction.

…………………………………………………………………………………………..

………………………………………………………………………………………….. (1)

116

37. Bromoethane reacts with ammonia as follows.

CH3CH2Br + NH3 → CH3CH2NH3+ + Br–

CH3CH2NH3+ + NH3 → CH3CH2NH2 + NH4+

The mechanism for this reaction is described as SN2.

(a) State the meaning of each of the symbols in SN2.

.....................................................................................................................................

..................................................................................................................................... (2)

(b) State the name of the organic product of the reaction, CH3CH2NH2.

..................................................................................................................................... (1)

117

Food Chemistry F.1 – Food Groups

1. Distinguish between a food and a nutrient

2. Describe the chemical composition of lipids (fats and oils), carbohydrates and proteins

F.2 – Fats and Oils 1. Describe the difference in structure between saturated and unsaturated (mono- and poly-unsaturated)

fatty acids

2. Predict the degree of crystallization (solidification) and melting point of fats and oils from their struc-

tures, and explain the relevance of this property in the home and in industry

3. Deduce the stability of fats and oils from their structure

4. Describe the process of hydrogenation of unsaturated fats

5. Discuss the advantages and disadvantages of hydrogenating fats and oils

F.3 – Shelf Life 1. Explain the meaning of the term shelf life

2. Discuss the factors that affect the shelf life and quality of food

3. Describe the rancidity of fats

4. Compare the process of hydrolytic and oxidative rancidity in lipids

5. Describe ways to minimize the rate of rancidity and prolong the shelf life of food

6. Describe the traditional methods used by different cultures to extend the shelf life of foods

7. Define the term antioxidant

8. List examples of common naturally occurring antioxidants and their sources

9. Compare the structural features of the major synthetic antioxidants in food

10. Discuss the advantages and disadvantages associated with natural and synthetic antioxidants

11. List some antioxidants found in the traditional foods of different cultures that may have health benefits

F.4 – Color 1. Distinguish between a dye and a pigment

2. Explain the occurrence of color in naturally occurring pigments

3. Describe the range of colors and sources of the naturally occurring pigments anthocyanins, carotenoids,

chlorophyll and heme

4. Describe the factors that affect the color stability of anthocyanins, carotenoids, chlorophyll and heme

5. Discuss the safety issues associated with the use of synthetic colorants in food

6. Compare the two processes of non-enzymatic browning (Millard reaction) and carmelization that cause

the browning of food

F.5 – Genetically Modified Foods 1. Define a genetically modified (GM) food

2. Discuss the benefits and concerns of using GM foods

118

F.6 – Texture 1. Describe a dispersed system in food

2. Distinguish between the following types of dispersed systems: suspensions, emulsions and foams in food

3. Describe the action of emulsifiers

F.7 – Oxidative Rancidity (auto-oxidation) (AHL) 1. Describe the steps in the free-radical chain mechanism during oxidative rancidity

F.8 – Antioxidants (AHL) 1. Explain the differences between the three main types of antioxidants

F.9 – Stereochemistry in Food (AHL) 1. Explain the three different conventions used for naming the different enantiomeric forms

2. Distinguish between the properties of the different enantiomeric forms of stereoisomers found in food

F.10 – Chemical Structure and Color (AHL) 1. Compare the similarities and differences in the structure of natural pigments: anthocyanins, carotenoids,

chlorophyll and heme

2. Explain why anthocyanins, carotenoids, chlorophyll and heme form colored compounds while many other

organic molecules are colorless

3. Deduce whether anthocyanins and carotenoids are water- or fat-soluble from their structures

119

HL 2 Food Chemistry Homework Name ________________________

1. Distinguish between a food and a nutrient .

................................................................................................................................................

................................................................................................................................................

................................................................................................................................................

................................................................................................................................................ (2)

2. Oils can be hydrogenated. One possible problem is that partial hydrogenation may occur which produces an

oil containing trans fatty acids. Explain the structural difference between a cis fatty acid and a trans fatty

acid and state one disadvantage of ingesting oils containing trans fatty acids.

Difference:

................................................................................................................................................

................................................................................................................................................

Disadvantage:

................................................................................................................................................

................................................................................................................................................ (2)

3. The preservation of food is important around the world.

(a) Explain the meaning of the term shelf life.

......................................................................................................................................

......................................................................................................................................

......................................................................................................................................

...................................................................................................................................... (2)

(b) Discuss two factors that can affect the shelf life of food.

......................................................................................................................................

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......................................................................................................................................

......................................................................................................................................

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120

4. Distinguish between the following types of dispersed systems.

Suspension ............................................................................................................................

...............................................................................................................................................

Emulsion ...............................................................................................................................

...............................................................................................................................................

Foam .....................................................................................................................................

............................................................................................................................................... (3)

5. State three characteristic features of all monosaccharide molecules.

................................................................................................................................................

................................................................................................................................................

................................................................................................................................................

................................................................................................................................................

................................................................................................................................................

................................................................................................................................................ (3)

6. Describe the step by step mechanism for the free radical reaction that causes an oily fish such as mackerel

to become rancid. Include the name of each step and an equation for each step in your answer.

Step 1:

................................................................................................................................................

................................................................................................................................................

................................................................................................................................................

Step 2:

................................................................................................................................................

................................................................................................................................................

................................................................................................................................................

Step 3:

................................................................................................................................................

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121

7. Genetic engineering is an important technique used to alter the properties of foods.

(a) Define the term genetically modified (GM) food.

......................................................................................................................................

...................................................................................................................................... (1)

(b) Discuss one benefit and one concern of using genetically modified (GM) crops in food.

Benefit:

......................................................................................................................................

......................................................................................................................................

Concern:

......................................................................................................................................

...................................................................................................................................... (2)

8. The structure below shows –(l)-carvone.

–(l)-carvone has another optical isomer.

(a) State its name and, by means of a diagram, predict its structure in the space above.

...................................................................................................................................... (2)

(b) Describe the structural feature of the carvone molecule that allows it to exist as optical isomers.

......................................................................................................................................

...................................................................................................................................... (1)

(c) State the effect that the difference in the structures of the two optical isomers has on the flavour of the

compounds.

......................................................................................................................................

...................................................................................................................................... (2)

122

(d) Explain what is meant by the –(l) notation, and how this is different to the (R) notation.

......................................................................................................................................

......................................................................................................................................

......................................................................................................................................

...................................................................................................................................... (2)

9. Antioxidants can be used to prolong the shelf life of food.

(a) Define the terms:

(i) shelf life

...........................................................................................................................

........................................................................................................................... (1)

(ii) antioxidant.

...........................................................................................................................

........................................................................................................................... (1)

(b) Using Table 22 of the Data Booklet, and the structure of propyl gallate (PG) below, compare the

structural features of the three common antioxidants 3-BHA, BHT and PG.

......................................................................................................................................

......................................................................................................................................

......................................................................................................................................

......................................................................................................................................

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...................................................................................................................................... (4)

123

(c) State one example of a common naturally occurring antioxidant and state one possible long-term

health benefit of consuming food in which it is present.

Antioxidant:

......................................................................................................................................

Long-term health benefit:

......................................................................................................................................

...................................................................................................................................... (2)

10. Fats and vegetable oils are triesters of glycerol with fatty acids. Many of these acids contain 18 carbon at-

oms. The table shows the relative percentages of various C18 fatty acid chains in four common fats and oils.

(a) Deduce which fat or oil from the table could best be described as:

saturated

......................................................................................................................................

mono-unsaturated

......................................................................................................................................

poly-unsaturated.

...................................................................................................................................... (2)

(b) (i) Explain the meaning of the term shelf life and suggest which fat or oil from the table would

have the shortest shelf life.

...........................................................................................................................

...........................................................................................................................

........................................................................................................................... (2)

(ii) Describe two ways in which shelf life could be increased.

...........................................................................................................................

...........................................................................................................................

...........................................................................................................................

........................................................................................................................... (2)

Fat/Oil C17H35COO–

/ %

C17H33COO–

/ %

C17H31COO–

/ %

C17H29COO–

/ %

Tallow 52 44 3 1

Linseed Oil 5 32 18 45

Olive Oil 2 83 15 0

Peanut Oil 7 47 46 0

124

11. (a) Explain why pigments such as anthocyanins are colored.

......................................................................................................................................

......................................................................................................................................

......................................................................................................................................

...................................................................................................................................... (2)

(b) The wavelength of visible light lies between 400 and 750 nm. The absorption spectrum of a particular

anthocyanin is shown below.

(i) Explain what effect, if any, the absorption at 375 nm will have on the color of the anthocyanin.

...........................................................................................................................

........................................................................................................................... (1)

(ii) Explain what effect, if any, the absorption at 530 nm will have on the color of the anthocyanin.

...........................................................................................................................

........................................................................................................................... (1)

(c) List two factors which could alter the precise color of a particular anthocyanin.

......................................................................................................................................

......................................................................................................................................

......................................................................................................................................

......................................................................................................................................

...................................................................................................................................... (2)