﻿ Chapter Review - Chemical Quantities

# Chemical QuantitiesChapter Review

Key Terms

· mole ratio (9.2)

· stoichiometry (9.3)

· limiting reactant (9.4)

· limiting reagent (9.4)

· theoretical yield (9.6)

· percent yield (9.6)

For Review

· A balanced chemical equation gives relative numbers (or moles) of reactant and product molecules that participate in a chemical reaction.

· Stoichiometric calculations involve using a balanced chemical equation to determine the amounts of reactants needed or products formed in a reaction.

· To convert between moles of reactants and moles of products, we use mole ratios derived from the balanced chemical equation.

· To calculate masses from the moles of reactants needed or products formed, we can use the molar masses of substances for finding the masses (g) needed or formed.

· Often, reactants in a chemical reaction are not present in stoichiometric quantities (i.e., they do not “run out” at the same time).

o In this case, we must determine which reactant runs out first and thus limits the amount of products that can form—this is called the limiting reactant.

· The actual yield (amount produced) of a reaction is usually less than the maximum expected (theoretical yield).

· The actual yield is often expressed as a percentage of the theoretical yield:

Active Learning Questions

These questions are designed to be considered by groups of students in class. Often these questions work well for introducing a particular topic in class.

· 1.

Relate Active Learning Question 2 from Chapter 2 to the concepts of chemical stoichiometry.

· 2.

You are making cookies and are missing a key ingredient—eggs. You have plenty of the other ingredients, except that you have only cups of butter and no eggs. You note that the recipe calls for cups of butter and eggs (plus the other ingredients) to make dozen cookies. You telephone a friend and have him bring you some eggs.

a. How many eggs do you need?

b. If you use all the butter (and get enough eggs), how many cookies can you make?

Unfortunately, your friend hangs up before you tell him how many eggs you need. When he arrives, he has a surprise for you—to save time he has broken the eggs in a bowl for you. You ask him how many he brought, and he replies, “All of them, but I spilled some on the way over.” You weigh the eggs and find that they weigh g. Assuming that an average egg weighs g:

c. How much butter is needed to react with all the eggs?

d. How many cookies can you make?

e. Which will you have left over, eggs or butter?

f. How much is left over?

g. Relate this question to the concepts of chemical stoichi ometry.

· 3.

Nitrogen and hydrogen react to form ammonia . Consider the mixture of and in a closed container as illustrated below: Assuming the reaction goes to completion, draw a representation of the product mixture. Explain how you arrived at this representation.

· 4.

Which of the following equations best represents the reaction for Question 3?

a.

b.

c.

d.

e. For choices you did not pick, explain what you feel is wrong with them, and justify the choice you did pick.

· 5.

You know that chemical reacts with chemical . You react g with g . What information do you need to know to determine the amount of product that will be produced? Explain.

· 6.

If g of hydrogen gas is reacted with g of oxygen gas according to the equation

we should not expect to form g of water. Why not? What mass of water can be produced with a complete reaction?

· 7.

The limiting reactant in a reaction:

a. has the lowest coefficient in a balanced equation.

b. is the reactant for which you have the fewest number of moles.

c. has the lowest ratio: moles available/coefficient in the balanced equation.

d. has the lowest ratio: coefficient in the balanced equation/moles available.

e. None of the above.

For choices you did not pick, explain what you feel is wrong with them, and justify the choice you did pick.

· 8.

Given the equation , if moles of is reacted with moles of , which of the following is true?

a. The limiting reactant is the one with the higher molar mass.

b. is the limiting reactant because you need moles of and have moles.

c. is the limiting reactant because you have fewer moles of than moles of .

d. is the limiting reactant because three molecules react with every one molecule.

e. Neither reactant is limiting.

For choices you did not pick, explain what you feel is wrong with them, and justify the choice you did pick.

· 9.

What happens to the weight of an iron bar when it rusts?

a. There is no change because mass is always conserved.

b. The weight increases.

c. The weight increases, but if the rust is scraped off, the bar has the original weight.

d. The weight decreases.

Justify your choice and, for choices you did not pick, explain what is wrong with them. Explain what it means for something to rust.

· 10.

Consider the equation . If you mix mole of and mole of , how many moles of can be produced?

· 11.

What is meant by the term mole ratio? Give an example of a mole ratio, and explain how it is used in solving a stoichiometry problem.

· 12.

Which would produce a greater number of moles of product: a given amount of hydrogen gas reacting with an excess of oxygen gas to produce water, or the same amount of hydrogen gas reacting with an excess of nitrogen gas to make ammonia? Support your answer.

· 13.

Consider a reaction represented by the following balanced equation

You find that it requires equal masses of and so that there are no reactants left over. Which of the following is true? Justify your choice.

a. The molar mass of must be greater than the molar mass of .

b. The molar mass of must be less than the molar mass of .

c. The molar mass of must be the same as the molar mass of .

· 14.

Consider a chemical equation with two reactants forming one product. If you know the mass of each reactant, what else do you need to know to determine the mass of the product? Why isn’t the mass necessarily the sum of the mass of the reactants? Provide a real example of such a reaction, and support your answer mathematically.

· 15.

Consider the balanced chemical equation

When equal masses of and are reacted, which is limiting, or ? Justify your choice.

a. If the molar mass of is greater than the molar mass of , then must be limiting.

b. If the molar mass of is less than the molar mass of , then must be limiting.

c. If the molar mass of is greater than the molar mass of , then must be limiting.

d. If the molar mass of is less than the molar mass of , then must be limiting.

· 16.

Which of the following reaction mixtures would produce the greatest amount of product, assuming all went to completion? Justify your choice.

Each involves the reaction symbolized by the equation

a. moles of and moles of .

b. moles of and moles of .

c. moles of and mole of .

d. moles of and mole of .

e. Each would produce the same amount of product.

· 17.

Baking powder is a mixture of cream of tartar and baking soda . When it is placed in an oven at typical baking temperatures (as part of a cake, for example), it undergoes the following reaction ( makes the cake rise):

You decide to make a cake one day, and the recipe calls for baking powder. Unfortunately, you have no baking powder. You do have cream of tartar and baking soda, so you use stoichiometry to figure out how much of each to mix.

Of the following choices, which is the best way to make baking powder? The amounts given in the choices are in teaspoons (that is, you will use a teaspoon to measure the baking soda and cream of tartar). Justify your choice.

Assume a teaspoon of cream of tartar has the same mass as a teaspoon of baking soda.

a. Add equal amounts of baking soda and cream of tartar.

b. Add a bit more than twice as much cream of tartar as baking soda.

c. Add a bit more than twice as much baking soda as cream of tartar.

d. Add more cream of tartar than baking soda, but not quite twice as much.

e. Add more baking soda than cream of tartar, but not quite twice as much.

· 18.

You have seven closed containers each with equal masses of chlorine gas . You add g of sodium to the first sample, g of sodium to the second sample, and so on (adding g of sodium to the seventh sample). Sodium and chloride react to form sodium chloride according to the equation

After each reaction is complete, you collect and measure the amount of sodium chloride formed. A graph of your results is shown below. a. Explain the shape of the graph.

b. Calculate the mass of formed when g of sodium is used.

c. Calculate the mass of in each container.

d. Calculate the mass of formed when g of sodium is used.

e. Identify the leftover reactant and determine its mass for parts b and d above.

· 19.

You have a chemical in a sealed glass container filled with air. The setup is sitting on a balance as shown below. The chemical is ignited by means of a magnifying glass focusing sunlight on the reactant. After the chemical has completely burned, which of the following is true? Explain your answer. a. The balance will read less than g.

b. The balance will read g.

c. The balance will read greater than g.

d. Cannot be determined without knowing the identity of the chemical.

· 20.

Consider an iron bar on a balance as shown. As the iron bar rusts, which of the following is true? Explain your answer.

a. The balance will read less than g.

b. The balance will read g.

c. The balance will read greater than g.

d. The balance will read greater than g, but if the bar is removed, the rust scraped off, and the bar replaced, the balance will read g.

· 21.

Consider the reaction between and represented below. What is the balanced equation for this reaction, and what is the limiting reactant?

Questions and Problems: 9.1 Information Given by Chemical Equations

Questions and Problems with answers below also have full solutions in the Student Solutions Guide.

Questions

· 1.

What do the coefficients of a balanced chemical equation tell us about the proportions in which atoms and molecules react on an individual (microscopic) basis?

· 2.

The vigorous reaction between aluminum and iodine gives the balanced equation: What do the coefficients in this balanced chemical equation tell us about the proportions in which these substances react on a macroscopic (mole) basis?

· 3.

Although mass is a property of matter we can conveniently measure in the laboratory, the coefficients of a balanced chemical equation are not directly interpreted on the basis of mass. Explain why.

· 4.

Which of the following statements is true for the reaction of nitrogen gas with hydrogen gas to produce ammonia ? Choose the best answer.

a. Subscripts can be changed to balance this equation, just as they can be changed to balance the charges when writing the formula for an ionic compound.

b. The nitrogen and hydrogen will not react until you have added the correct mole ratios.

c. The mole ratio of nitrogen to hydrogen in the balanced equation is .

d. Ammonia will not form unless mole of nitrogen and moles of hydrogen have been added.

e. The balanced equation allows you to predict how much ammonia you will make based on the amount of nitrogen and hydrogen present.

Problems

· 5.

For each of the following reactions, give the balanced equation for the reaction and state the meaning of the equation in terms of the numbers of individual molecules and in terms of moles of molecules.

a.

b.

c.

d.

· 6.

For each of the following reactions, give the balanced chemical equation for the reaction and state the meaning of the equation in terms of individual molecules and in terms of moles of molecules.

a.

b.

c.

d.

Questions and Problems: 9.2 Mole—Mole Relationships

Questions and Problems with answers below also have full solutions in the Student Solutions Guide.

Questions

· 7.

Consider the reaction represented by the chemical equation

Because the coefficients of the balanced chemical equation are all equal to , we know that exactly g of will react with exactly g of . True or false? Explain.

· 8.

For the balanced chemical equation for the combination reaction of hydrogen gas and oxygen gas

explain why we know that g of reacting with g of will not result in the production of g of .

· 9.

Consider the balanced chemical equation

What mole ratio would you use to calculate how many moles of oxygen gas would be needed to react completely with a given number of moles of aluminum metal? What mole ratio would you use to calculate the number of moles of product that would be expected if a given number of moles of aluminum metal reacts completely?

· 10.

Write the balanced chemical equation for the complete combustion of heptene, . In combustion, heptene reacts with oxygen to produce carbon dioxide and water. What is the mole ratio that would enable you to calculate the number of moles of oxygen needed to react exactly with a given number of moles of heptene? What mole ratios would you use to calculate how many moles of each product form from a given number of moles of heptene?

Problems

· 11.

For each of the following balanced chemical equations, calculate how many moles of product(s) would be produced if mole of the first reactant were to react completely.

a.

b.

c.

d.

· 12.

For each of the following unbalanced chemical equations, calculate how many moles of each product would be produced by the complete conversion of mole of the reactant indicated in boldface. State clearly the mole ratio used for the conversion.

a.

b.

c.

d.

· 13.

For each of the following balanced chemical equations, calculate how many grams of the product(s) would be produced by complete reaction of mole of the first reactant.

a.

b.

c.

d.

· 14.

For each of the following balanced chemical equations, calculate how many moles and how many grams of each product would be produced by the complete conversion of mole of the reactant indicated in boldface. State clearly the mole ratio used for each conversion.

a.

b.

c.

d.

· 15.

For each of the following unbalanced equations, indicate how many moles of the second reactant would be required to react exactly with mole of the first reactant. State clearly the mole ratio used for the conversion.

a.

b.

c.

d.

· 16.

For each of the following unbalanced equations, indicate how many moles of the first product are produced if mole of the second product forms. State clearly the mole ratio used for each conversion.

a.

b.

c.

d.

Questions and Problems: 9.3 Mass Calculations

Questions and Problems with answers below also have full solutions in the Student Solutions Guide.

Questions

· 17.

What quantity serves as the conversion factor between the mass of a sample and how many moles the sample contains?

· 18.

What does it mean to say that the balanced chemical equation for a reaction describes the stoichiometry of the reaction?

Problems directs you to the Chemistry in Focus feature in the chapter

· 19.

Using the average atomic masses given inside the front cover of this book, calculate how many moles of each substance the following masses represent.

a. g of silicon,

b. mg of gold(III) chloride,

c. kg of sulfur,

d. g of iron(III) chloride,

e. g of magnesium oxide,

· 20.

Using the average atomic masses given inside the front cover of this book, calculate the number of moles of each substance contained in the following masses.

a. g of silver

b. mg of ammonium sulfide

c. of uranium

d. kg of sulfur dioxide

e. g of iron(III) nitrate

· 21.

Using the average atomic masses given inside the front cover of this book, calculate the mass in grams of each of the following samples.

a. moles of germanium,

c. mole of ammonia,

d. moles of hexane,

e. moles of iodine monochloride,

· 22.

Using the average atomic masses given inside the front cover of this book, calculate the mass in grams of each of the following samples.

a. mole of potassium nitride

b. mmol of neon

c. mole of manganese(II) oxide

d. moles of silicon dioxide

e. mole of iron(III) phosphate

· 23.

For each of the following unbalanced equations, calculate how many moles of the second reactant would be required to react completely with moles of the first reactant.

a.

b.

c.

d.

· 24.

For each of the following unbalanced equations, calculate how many moles of the second reactant would be required to react completely with grams of the first reactant.

a.

b.

c.

d.

· 25.

For each of the following unbalanced equations, calculate how many grams of each product would be produced by complete reaction of g of the reactant indicated in boldface. Indicate clearly the mole ratio used for the conversion.

a.

b.

c.

d.

· 26.

Boron nitride reacts with iodine monofluoride in trichlorofluoromethane at to produce pure nitrogen triiodide and by-product .

a. What mass of iodine monofluoride must be used to produce g of nitrogen triiodide?

b. When g of nitrogen triiodide is produced, what is the maximum mass of by-product created?

· 27.

“Smelling salts,” which are used to revive someone who has fainted, typically contain ammonium carbonate, . Ammonium carbonate decomposes readily to form ammonia, carbon dioxide, and water. The strong odor of the ammonia usually restores consciousness in the person who has fainted. The unbalanced equation is

Calculate the mass of ammonia gas that is produced if g of ammonium carbonate decomposes completely.

· 28.

Calcium carbide, , can be produced in an electric furnace by strongly heating calcium oxide (lime) with carbon. The unbalanced equation is

Calcium carbide is useful because it reacts readily with water to form the flammable gas acetylene, , which is used extensively in the welding industry. The unbalanced equation is

What mass of acetylene gas, , would be produced by complete reaction of g of calcium carbide?

· 29.

When elemental carbon is burned in the open atmosphere, with plenty of oxygen gas present, the product is carbon dioxide.

However, when the amount of oxygen present during the burning of the carbon is restricted, carbon monoxide is more likely to result.

What mass of each product is expected when a -g sample of pure carbon is burned under each of these conditions?

· 30.

If baking soda (sodium hydrogen carbonate) is heated strongly, the following reaction occurs:

Calculate the mass of sodium carbonate that will remain if a -g sample of sodium hydrogen carbonate is heated.

· 31.

Although we usually think of substances as “burning” only in oxygen gas, the process of rapid oxidation to produce a flame may also take place in other strongly oxidizing gases. For example, when iron is heated and placed in pure chlorine gas, the iron “burns” according to the following (unbalanced) reaction:

How many milligrams of iron(III) chloride result when mg of iron is reacted with an excess of chlorine gas?

· 32.

When yeast is added to a solution of glucose or fructose, the sugars are said to undergo fermentation, and ethyl alcohol is produced.

This is the reaction by which wines are produced from grape juice. Calculate the mass of ethyl alcohol, , produced when g of glucose, , undergoes this reaction.

· 33.

Sulfurous acid is unstable in aqueous solution and gradually decomposes to water and sulfur dioxide gas (which explains the choking odor associated with sulfurous acid solutions).

If g of sulfurous acid undergoes this reaction, what mass of sulfur dioxide is released?

· 34.

Small quantities of ammonia gas can be generated in the laboratory by heating an ammonium salt with a strong base. For example, ammonium chloride reacts with sodium hydroxide according to the following balanced equation:

What mass of ammonia gas is produced if g of ammonium chloride reacts completely?

· 35.

Elemental phosphorus burns in oxygen with an intensely hot flame, producing a brilliant light and clouds of the oxide product. These properties of the combustion of phosphorus have led to its being used in bombs and incendiary devices for warfare.

If g of phosphorus is burned, what mass of oxygen does it combine with?

· 36.

Although we tend to make less use of mercury these days because of the environmental problems created by its improper disposal, mercury is still an important metal because of its unusual property of existing as a liquid at room temperature. One process by which mercury is produced industrially is through the heating of its common ore cinnabar (mercuric sulfide, ) with lime (calcium oxide, ).

What mass of mercury would be produced by complete reaction of kg of ?

· 37.

Ammonium nitrate has been used as a high explosive because it is unstable and decomposes into several gaseous substances. The rapid expansion of the gaseous substances produces the explosive force.

Calculate the mass of each product gas if g of ammonium nitrate reacts.

· 38.

If common sugars are heated too strongly, they char as they decompose into carbon and water vapor. For example, if sucrose (table sugar) is heated, the reaction is

What mass of carbon is produced if g of sucrose decomposes completely?

· 39.

Thionyl chloride, , is used as a very powerful drying agent in many synthetic chemistry experiments in which the presence of even small amounts of water would be detrimental. The unbalanced chemical equation is

Calculate the mass of water consumed by complete reaction of g of .

· 40. In the “Chemistry in Focus” segment Cars of the Future, the claim is made that the combustion of gasoline for some cars causes about lb of to be produced for each mile traveled.

Estimate the gas mileage of a car that produces about lb of per mile traveled. Assume gasoline has a density of g/mL and is octane . While this last part is not true, it is close enough for an estimation. The reaction can be represented by the following unbalanced chemical equation:

Questions and Problems: 9.5 Calculations Involving a Limiting Reactant

Questions and Problems with answers below also have full solutions in the Student Solutions Guide.

Questions

· 41.

Which of the following statements is(are) true?

a. A balanced equation relates the numbers of molecules of reactants and products (or numbers of moles of reactants and products).

b. To convert between moles of reactants and moles of products, we use mole ratios derived from the balanced equation.

c. Often reactants are not mixed in stoichiometric quantities (they do not “run out” at the same time). In that case, we must use the limiting reactant to calculate the amounts of products formed.

d. When a chemical reaction occurs, it must follow the law of conservation of mass.

· 42.

Explain how one determines which reactant in a process is the limiting reactant. Does this depend only on the masses of the reactant present? Give an example of how to determine the limiting reactant by using a Before—Change—After (BCA) table with a balanced chemical equation and reactant starting amounts.

· 43.

Consider the equation: . If g of reacts with g of , how is the limiting reactant determined? Choose the best answer and explain.

a. Choose the reactant with the smallest coefficient in the balanced chemical equation. So in this case, the limiting reactant is .

b. Choose the reactant with the smallest mass given. So in this case, the limiting reactant is .

c. The mass of each reactant must be converted to moles and then compared to the ratios in the balanced chemical equation. So in this case, the limiting reactant cannot be determined without the molar masses of and .

d. The mass of each reactant must be converted to moles first. The reactant with the fewest moles present is the limiting reactant. So in this case, the limiting reactant cannot be determined without the molar masses of and .

e. The mass of each reactant must be divided by their coefficients in the balanced chemical equation, and the smallest number present is the limiting reactant. So in this case, there is no limiting reactant because and are used up perfectly.

· 44.

Balance the following chemical equation, and then answer the question below.

Which reactant is the limiting reactant? Choose the best answer.

a. Both and are equally limiting because they react in a mole ratio.

b. is the limiting reactant because only moles are available compared with moles of .

c. is the limiting reactant because it is present in excess.

d. Neither nor is a limiting reactant because total moles are present on the reactant side compared with total moles on the product side.

e. The limiting reactant cannot be determined because the starting amounts are not given.

Problems

· 45.

For each of the following unbalanced reactions, suppose exactly moles of each reactant are taken. Determine which reactant is limiting, and also determine what mass of the excess reagent will remain after the limiting reactant is consumed. For each reaction, solve the problem three ways:

i. Set up and use Before—Change—After (BCA) tables.

ii. Compare the moles of reactants to see which runs out first.

iii. Consider the amounts of products that can be formed by completely consuming each reactant.

d.

e.

· 46.

For each of the following unbalanced chemical equations, suppose that exactly g of each reactant is taken. Determine which reactant is limiting, and calculate what mass of each product is expected (assuming that the limiting reactant is completely consumed).

.

a.

b.

c.

· 47.

For each of the following unbalanced chemical equations, suppose g of each reactant is taken. Show by calculation which reactant is the limiting reagent. Calculate the mass of each product that is expected.

.

a.

b.

c.

· 48.

For each of the following unbalanced chemical equations, suppose that exactly g of each reactant is taken. Determine which reactant is limiting, and calculate what mass of the product in boldface is expected (assuming that the limiting reactant is completely consumed).

.

a.

b.

c.

· 49.

For each of the following unbalanced chemical equations, suppose g of each reactant is taken. Show by calculation which reactant is limiting. Calculate the mass of each product that is expected.

.

a.

b.

c.

· 50.

For each of the following unbalanced chemical equations, suppose that exactly g of each reactant are taken. Using Before—Change—After (BCA) tables, determine which reactant is limiting, and calculate what mass of each product is expected. (Assume that the limiting reactant is completely consumed.)

.

a.

b.

c.

· 51.

Lead(II) carbonate, also called “white lead,” was formerly used as a pigment in white paints. However, because of its toxicity, lead can no longer be used in paints intended for residential homes. Lead(II) carbonate is prepared industrially by reaction of aqueous lead(II) acetate with carbon dioxide gas. The unbalanced equation is

Suppose an aqueous solution containing g of lead(II) acetate is treated with g of carbon dioxide. Calculate the theoretical yield of lead carbonate.

· 52.

Copper(II) sulfate has been used extensively as a fungicide (kills fungus) and herbicide (kills plants). Copper(II) sulfate can be prepared in the laboratory by reaction of copper(II) oxide with sulfuric acid. The unbalanced equation is

If g of copper(II) oxide is treated with g of pure sulfuric acid, which reactant would limit the quantity of copper(II) sulfate that could be produced?

· 53.

Lead(II) oxide from an ore can be reduced to elemental lead by heating in a furnace with carbon.

Calculate the expected yield of lead if kg of lead oxide is heated with kg of carbon.

· 54.

If steel wool (iron) is heated until it glows and is placed in a bottle containing pure oxygen, the iron reacts spectacularly to produce iron(III) oxide.

If g of iron is heated and placed in a bottle containing mole of oxygen gas, what mass of iron(III) oxide is produced?

· 55.

A common method for determining how much chloride ion is present in a sample is to precipitate the chloride from an aqueous solution of the sample with silver nitrate solution and then to weigh the silver chloride that results. The balanced net ionic reaction is

Suppose a -g sample of pure sodium chloride is dissolved in water and is then treated with a solution containing g of silver nitrate. Will this quantity of silver nitrate be capable of precipitating all the chloride ion from the sodium chloride sample?

· 56.

Although many sulfate salts are soluble in water, calcium sulfate is not ( Table 7.1). Therefore, a solution of calcium chloride will react with sodium sulfate solution to produce a precipitate of calcium sulfate. The balanced equation is

If a solution containing g of calcium chloride is combined with a solution containing of sodium sulfate, which is the limiting reactant? Which reactant is present in excess?

· 57.

Hydrogen peroxide is used as a cleaning agent in the treatment of cuts and abrasions for several reasons. It is an oxidizing agent that can directly kill many microorganisms; it decomposes upon contact with blood, releasing elemental oxygen gas (which inhibits the growth of anaerobic microorganisms); and it foams upon contact with blood, which provides a cleansing action. In the laboratory, small quantities of hydrogen peroxide can be prepared by the action of an acid on an alkaline earth metal peroxide, such as barium peroxide.

What amount of hydrogen peroxide should result when g of barium peroxide is treated with mL of hydrochloric acid solution containing g of per mL?

· 58.

Silicon carbide, , is one of the hardest materials known. Surpassed in hardness only by diamond, it is sometimes known commercially as carborundum. Silicon carbide is used primarily as an abrasive for sandpaper and is manufactured by heating common sand (silicon dioxide, ) with carbon in a furnace.

What mass of silicon carbide should result when kg of pure sand is heated with an excess of carbon?

Questions and Problems: 9.6 Percent Yield

Questions and Problems with answers below also have full solutions in the Student Solutions Guide.

Questions

· 59.

Your text talks about several sorts of “yield” when experiments are performed in the laboratory. Students often confuse these terms. Define, compare, and contrast what are meant by theoretical yield, actual yield, and percent yield.

· 60.

The text explains that one reason why the actual yield for a reaction may be less than the theoretical yield is side reactions. Suggest some other reasons why the percent yield for a reaction might not be .

· 61.

According to his prelaboratory theoretical yield calculations, a student’s experiment should have produced g of magnesium oxide. When he weighed his product after reaction, only g of magnesium oxide was present. What is the student’s percent yield?

· 62.

An air bag is deployed by utilizing the following reaction (the nitrogen gas produced inflates the air bag):

If g of is decomposed, what theoretical mass of sodium should be produced? If only g of sodium is actually collected, what is the percent yield?

Problems

· 63.

The compound sodium thiosulfate pentahydrate, , is important commercially to the photography business as “hypo,” because it has the ability to dissolve unreacted silver salts from photographic film during development. Sodium thiosulfate pentahydrate can be produced by boiling elemental sulfur in an aqueous solution of sodium sulfite.

What is the theoretical yield of sodium thiosulfate pentahydrate when g of sulfur is boiled with g of sodium sulfite? Sodium thiosulfate pentahydrate is very soluble in water. What is the percent yield of the synthesis if a student doing this experiment is able to isolate (collect) only g of the product?

· 64.

Alkali metal hydroxides are sometimes used to “scrub” excess carbon dioxide from the air in closed spaces (such as submarines and spacecraft). For example, lithium hydroxide reacts with carbon dioxide according to the unbalanced chemical equation

Suppose a lithium hydroxide canister contains g of . What mass of will the canister be able to absorb? If it is found that after hours of use the canister has absorbed g of carbon dioxide, what percentage of its capacity has been reached?

· 65.

Although they were formerly called the inert gases, at least the heavier elements of Group 8 do form relatively stable compounds. For example, xenon combines directly with elemental fluorine at elevated temperatures in the presence of a nickel catalyst.

What is the theoretical mass of xenon tetrafluoride that should form when g of xenon is reacted with g of ? What is the percent yield if only g of is actually isolated?

· 66.

Solid copper can be produced by passing gaseous ammonia over solid copper(II) oxide at high temperatures. Ken O’Donoghue

The other products of the reaction are nitrogen gas and water vapor. The balanced equation for this reaction is:

What is the theoretical yield of solid copper that should form when g of is reacted with g of ? If only g of copper is actually collected, what is the percent yield?

· 67.

Natural waters often contain relatively high levels of calcium ion, , and hydrogen carbonate ion (bicarbonate), , from the leaching of minerals into the water. When such water is used commercially or in the home, heating of the water leads to the formation of solid calcium carbonate, , which forms a deposit (“scale”) on the interior of boilers, pipes, and other plumbing fixtures.

If a sample of well water contains mg of per milliliter, what mass of scale would mL of this water be capable of depositing?

· 68.

One process for the commercial production of baking soda (sodium hydrogen carbonate) involves the following reaction, in which the carbon dioxide is used in its solid form (“dry ice”) both to serve as a source of reactant and to cool the reaction system to a temperature low enough for the sodium hydrogen carbonate to precipitate:

Because they are relatively cheap, sodium chloride and water are typically present in excess. What is the expected yield of when one performs such a synthesis using g of ammonia and g of dry ice, with an excess of and water?

· 69.

A favorite demonstration among chemistry instructors, to show that the properties of a compound differ from those of its constituent elements, involves iron filings and powdered sulfur. If the instructor takes samples of iron and sulfur and just mixes them together, the two elements can be separated from one another with a magnet (iron is attracted to a magnet, sulfur is not). If the instructor then combines and heats the mixture of iron and sulfur, a reaction takes place and the elements combine to form iron(II) sulfide (which is not attracted by a magnet).

Suppose g of iron filings is combined with g of sulfur. What is the theoretical yield of iron(II) sulfide?

· 70.

When the sugar glucose, , is burned in air, carbon dioxide and water vapor are produced. Write the balanced chemical equation for this process, and calculate the theoretical yield of carbon dioxide when g of glucose is burned completely.

· 71.

When elemental copper is strongly heated with sulfur, a mixture of and is produced, with predominating.

What is the theoretical yield of when g of is heated with g of ? (Assume only is produced in the reaction.) What is the percent yield of if only g of can be isolated from the mixture?

· 72.

Barium chloride solutions are used in chemical analysis for the quantitative precipitation of sulfate ion from solution.

Suppose a solution is known to contain on the order of mg of sulfate ion. What mass of barium chloride should be added to guarantee precipitation of all the sulfate ion?

· 73.

The traditional method of analysis for the amount of chloride ion present in a sample is to dissolve the sample in water and then slowly to add a solution of silver nitrate. Silver chloride is very insoluble in water, and by adding a slight excess of silver nitrate, it is possible to effectively remove all chloride ion from the sample.

Suppose a -g sample is known to contain chloride ion by mass. What mass of silver nitrate must be used to completely precipitate the chloride ion from the sample? What mass of silver chloride will be obtained?

· 74.

For each of the following reactions, give the balanced equation for the reaction and state the meaning of the equation in terms of numbers of individual molecules and in terms of moles of molecules.

a.

b.

c.

d.

· 75.

True or false? For the reaction represented by the balanced chemical equation

for mole of , mol of will be needed.

· 76.

Consider the balanced equation

What mole ratio enables you to calculate the number of moles of oxygen needed to react exactly with a given number of moles of ? What mole ratios enable you to calculate how many moles of each product form from a given number of moles of ?

· 77.

For each of the following balanced reactions, calculate how many moles of each product would be produced by complete conversion of mole of the reactant indicated in boldface. Indicate clearly the mole ratio used for the conversion.

a.

b.

c.

d.

· 78.

For each of the following balanced equations, indicate how many moles of the product could be produced by complete reaction of g of the reactant indicated in boldface. Indicate clearly the mole ratio used for the conversion.

a.

b.

c.

d.

· 79.

Using the average atomic masses given inside the front cover of the text, calculate how many moles of each substance the following masses represent.

a. g of copper(II) sulfate

b. g of barium nitrate

c. mg of water

d. g of tungsten

e. lb of sulfur

f. g of ethyl alcohol,

g. g of carbon

· 80.

Using the average atomic masses given inside the front cover of the text, calculate the mass in grams of each of the following samples.

a. moles of nitric acid

b. mole of mercury

c. mole of potassium chromate

d. moles of aluminum chloride

e. moles of sulfur hexafluoride

f. moles of ammonia

g. mole of sodium peroxide

· 81.

For each of the following incomplete and unbalanced equations, indicate how many moles of the second reactant would be required to react completely with mole of the first reactant.

a.

b.

c.

d.

· 82.

One step in the commercial production of sulfuric acid, , involves the conversion of sulfur dioxide, , into sulfur trioxide, .

If kg of reacts completely, what mass of should result?

· 83.

Many metals occur naturally as sulfide compounds; examples include and . Air pollution often accompanies the processing of these ores, because toxic sulfur dioxide is released as the ore is converted from the sulfide to the oxide by roasting (smelting). For example, consider the unbalanced equation for the roasting reaction for zinc:

How many kilograms of sulfur dioxide are produced when kg of is roasted in excess oxygen by this process?

· 84.

If sodium peroxide is added to water, elemental oxygen gas is generated:

Suppose g of sodium peroxide is added to a large excess of water. What mass of oxygen gas will be produced?

· 85.

When elemental copper is placed in a solution of silver nitrate, the following oxidation—reduction reaction takes place, forming elemental silver:

What mass of copper is required to remove all the silver from a silver nitrate solution containing mg of silver nitrate?

· 86.

When small quantities of elemental hydrogen gas are needed for laboratory work, the hydrogen is often generated by chemical reaction of a metal with acid. For example, zinc reacts with hydrochloric acid, releasing gaseous elemental hydrogen:

What mass of hydrogen gas is produced when g of zinc is reacted with excess aqueous hydrochloric acid?

· 87.

The gaseous hydrocarbon acetylene, , is used in welders’ torches because of the large amount of heat released when acetylene burns with oxygen.

How many grams of oxygen gas are needed for the complete combustion of g of acetylene?

· 88.

For each of the following unbalanced chemical equations, suppose exactly g of each reactant is taken. Determine which reactant is limiting, and calculate what mass of each product is expected, assuming that the limiting reactant is completely consumed.

a.

b.

c.

d.

· 89.

For each of the following unbalanced chemical equations, suppose g of each reactant is taken. Show by calculation which reactant is limiting. Calculate the theoretical yield in grams of the product in boldface.

a.

b.

c.

d.

· 90.

Hydrazine, , emits a large quantity of energy when it reacts with oxygen, which has led to hydrazine’s use as a fuel for rockets:

How many moles of each of the gaseous products are produced when g of pure hydrazine is ignited in the presence of g of pure oxygen? How many grams of each product are produced?

· 91.

Consider the following reaction:

a. If a container were to have only molecules of and molecules of , how many total molecules (reactant and product) would be present in the container after the above reaction goes to completion?

b. Using “microscopic” pictures, draw the total molecules present inside the container after the reaction occurs.

c. What mass of is present in the container after the reaction occurs? (Report your final answer to significant figures.)

· 92.

Before going to lab, a student read in his lab manual that the percent yield for a difficult reaction to be studied was likely to be only of the theoretical yield. The student’s prelab stoichiometric calculations predict that the theoretical yield should be g. What is the student’s actual yield likely to be?

ChemWork Problems

These multiconcept problems (and additional ones) are found interactively online with the same type of assistance a student would get from an instructor.

· 93.

Consider the following unbalanced chemical equation for the combustion of pentane :

If a -gram sample of pentane is burned in excess oxygen, what mass of water can be produced, assuming yield?

· 94.

A -g sample of impure sodium nitrate (contains sodium nitrate plus inert ingredients) was heated, converting all the sodium nitrate to g of sodium nitrite and oxygen gas. Determine the percent of sodium nitrate in the original sample.

· 95.

Consider the following unbalanced chemical equation.

If g of lithium hydroxide reacts with excess carbon dioxide, what mass of lithium carbonate will be produced?

· 96.

Over the years, the thermite reaction has been used for welding railroad rails, in incendiary bombs, and to ignite solid fuel rocket motors. The reaction is

a. What mass of iron(III) oxide must be used to produce g of iron?

b. What mass of aluminum must be used to produce g of iron?

c. What is the maximum mass of aluminum oxide that could be produced along with g of iron?

· 97.

Consider the following unbalanced chemical equation:

Determine the maximum number of moles of produced from moles of and moles of .

· 98.

Ammonia gas reacts with sodium metal to form sodium amide and hydrogen gas. The unbalanced chemical equation for this reaction is as follows:

Assuming that you start with g of ammonia gas and g of sodium metal and assuming that the reaction goes to completion, determine the mass (in grams) of each product.

· 99.

Sulfur dioxide gas reacts with sodium hydroxide to form sodium sulfite and water. The unbalanced chemical equation for this reaction is as follows:

Assuming you react g of sulfur dioxide with g of sodium hydroxide and assuming that the reaction goes to completion, calculate the mass of each product formed.

· 100.

The production capacity for acrylonitrile in the United States is over billion pounds per year. Acrylonitrile, the building block for polyacrylonitrile fibers and a variety of plastics, is produced from gaseous propylene, ammonia, and oxygen:

a. Assuming yield, determine the mass of acrylonitrile which can be produced from the mixture below:

 Mass Reactant g propylene g ammonia g oxygen

b. What mass of water is formed from your mixture?

c. Calculate the mass (in grams) of each reactant after the reaction is complete.

Chapters 8-9. Cumulative Review

Questions

Questions with answers below also have full solutions in the Student Solutions Guide.

· 1.

What does the average atomic mass of an element represent? What unit is used for average atomic mass? Express the atomic mass unit in grams. Why is the average atomic mass for an element typically not a whole number?

· 2.

Perhaps the most important concept in introductory chemistry concerns what a mole of a substance represents. The mole concept will come up again and again in later chapters in this book. What does one mole of a substance represent on a microscopic, atomic basis? What does one mole of a substance represent on a macroscopic, mass basis? Why have chemists defined the mole in this manner?

· 3.

How do we know that g of oxygen contains the same number of atoms as does g of carbon, and that g of sodium contains the same number of atoms as each of these? How do we know that g of contains the same number of carbon atoms as does g of carbon, but three times as many oxygen atoms as in g of oxygen, and twice as many sodium atoms as in g of sodium?

· 4.

Define molar mass. Using as an example, calculate the molar mass from the atomic masses of the elements.

· 5.

What is meant by the percent composition by mass for a compound? Describe in general terms how this information is obtained by experiment for new compounds. How can this information be calculated for known compounds?

· 6.

Define, compare, and contrast what are meant by the empirical and molecular formulas for a substance. What does each of these formulas tell us about a compound? What information must be known for a compound before the molecular formula can be determined? Why is the molecular formula an integer multiple of the empirical formula?

· 7.

When chemistry teachers prepare an exam question on determining the empirical formula of a compound, they usually take a known compound and calculate the percent composition of the compound from the formula. They then give students this percent composition data and have the students calculate the original formula. Using a compound of your choice, first use the molecular formula of the compound to calculate the percent composition of the compound. Then use this percent composition data to calculate the empirical formula of the compound.

· 8.

Rather than giving students straight percent composition data for determining the empirical formula of a compound (see Question 7), sometimes chemistry teachers will try to emphasize the experimental nature of formula determination by converting the percent composition data into actual experimental masses. For example, the compound contains carbon by mass. Rather than giving students the data in this form, a teacher might say instead, “When g of a compound was analyzed, it was found to contain g of carbon, with the remainder consisting of hydrogen.” Using the compound you chose for Question 7, and the percent composition data you calculated, reword your data as suggested in this problem in terms of actual “experimental” masses. Then from these masses, calculate the empirical formula of your compound.

· 9.

Balanced chemical equations give us information in terms of individual molecules reacting in the proportions indicated by the coefficients, and also in terms of macroscopic amounts (that is, moles). Write a balanced chemical equation of your choice, and interpret in words the meaning of the equation on the molecular and macroscopic levels.

· 10.

Consider the unbalanced equation for the combustion of propane:

First, balance the equation. Then, for a given amount of propane, write the mole ratios that would enable you to calculate the number of moles of each product as well as the number of moles of that would be involved in a complete reaction. Finally, show how these mole ratios would be applied if mole of propane is combusted.

· 11.

In the practice of chemistry one of the most important calculations concerns the masses of products expected when particular masses of reactants are used in an experiment. For example, chemists judge the practicality and efficiency of a reaction by seeing how close the amount of product actually obtained is to the expected amount. Using a balanced chemical equation and an amount of starting material of your choice, summarize and illustrate the various steps needed in such a calculation for the expected amount of product.

· 12.

What is meant by a limiting reactant in a particular reaction? In what way is the reaction “limited”? What does it mean to say that one or more of the reactants are present in excess? What happens to a reaction when the limiting reactant is used up?

· 13.

For a balanced chemical equation of your choice, and using g of each of the reactants in your equation, illustrate and explain how you would determine which reactant is the limiting reactant. Indicate clearly in your discussion how the choice of limiting reactant follows from your calculations.

· 14.

What do we mean by the theoretical yield for a reaction? What is meant by the actual yield? Why might the actual yield for an experiment be less than the theoretical yield? Can the actual yield be more than the theoretical yield?

Problems

Problems with answers below also have full solutions in the Student Solutions Guide.

· 15.

Consider -g samples of each of the following elements or compounds. Calculate the number of moles of the element or compound present in each sample.

a.

b.

c.

d.

e.

f.

g.

h.

· 16.

Calculate the percent by mass of the element whose symbol occurs first in the following compounds’ formulas.

a.

b.

c.

d.

e.

f.

g.

h.

· 17.

A compound was analyzed and was found to have the following percent composition by mass: sodium, ; carbon, ; oxygen, . Determine the empirical formula of the compound.

· 18.

For each of the following balanced equations, calculate how many grams of each product would form if g of the reactant listed first in the equation reacts completely (there is an excess of the second reactant).

a.

b.

c.

d.

· 19.

Consider the reaction as represented by the following unbalanced chemical equation:

You react g of hydrogen gas with g of oxygen gas. Determine the amount of reactant in excess (in grams) after the reaction is complete.

· 20.

Solid calcium carbide reacts with liquid water to produce acetylene gas and aqueous calcium hydroxide.

a. Write the balanced equation for the reaction that is occurring, including all phases.

b. If a -g sample of calcium carbide is initially reacted with g of water, which reactant is limiting?

c. Prove that mass is conserved for the reactant amounts used in part b.

· 21.

A traditional analysis for samples containing calcium ion was to precipitate the calcium ion with sodium oxalate solution and then to collect and weigh either the calcium oxalate itself or the calcium oxide produced by heating the oxalate precipitate:

Suppose a sample contained g of calcium ion. What theoretical yield of calcium oxalate would be expected? If only g of calcium oxalate is collected, what percentage of the theoretical yield does that represent?

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