Introductory Chemistry: A Foundation - Zumdahl S.S., DeCoste D.J. 2019

Reactions in Aqueous Solutions
Chapter Review

Key Terms

· precipitation (7.2)

· precipitate (7.2)

· precipitation reaction (7.2, 7.6)

· strong electrolyte (7.2)

· soluble solid (7.2)

· insoluble solid (7.2)

· slightly soluble (7.2)

· molecular equation (7.3)

· complete ionic equation (7.3)

· spectator ions (7.3)

· net ionic equation (7.3)

· acid (7.4)

· strong acid (7.4)

· base (7.4)

· strong base (7.4)

· salt (7.4)

· oxidation—reduction reaction (7.5, 7.6)

· precipitation reaction (7.6)

· double-displacement reaction (7.6)

· acid—base reaction (7.6)

· combustion reaction (7.7)

· synthesis (combination) reaction (7.7)

· decomposition reaction (7.7)

For Review

· Four driving forces favor chemical change.

o Formation of a solid

o Formation of water

o Transfer of electrons

o Formation of a gas

· A reaction in which a solid forms is called a precipitation reaction.

· Solubility rules help predict what solid (if any) will form when solutions are mixed.

· Three types of equations are used to describe reactions in solution.

o Molecular (formula) equation, which shows the complete formulas of all reactants and products

o Complete ionic equation in which all strong electrolytes are shown as ions

o Net ionic equation, which includes only those components of the solution that undergo a change

§ Spectator ions (those that remain unchanged) are not shown in the net ionic equation.

· A strong acid is one in which virtually every molecule dissociates (ionizes) in water to an ion and an anion.

· A strong base is a metal hydroxide that is completely soluble in water, giving separate ions and cations.

· The products of the reaction of a strong acid and a strong base are water and a salt.

· Reactions between metals and nonmetals involve a transfer of electrons from the metal to the nonmetal, which is called an oxidation—reduction reaction.

· Reactions can be classified in various ways.

o A synthesis reaction is one in which a compound forms from simpler substances, such as elements.

o A decomposition reaction occurs when a compound is broken down into simpler substances.

o A combustion reaction is an oxidation—reduction reaction that involves .

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.

Consider the mixing of aqueous solutions of lead(II) nitrate and sodium iodide to form a solid.

a. Name the possible products, and determine the formulas of these possible products.

b. What is the precipitate? How do you know?

c. Must the subscript for an ion in a reactant stay the same as the subscript of that ion in a product? Explain your answer.

· 2.

Assume a highly magnified view of a solution of that allows you to “see” the . Draw this magnified view. If you dropped in a piece of magnesium, the magnesium would disappear, and hydrogen gas would be released. Represent this change using symbols for the elements, and write the balanced equation.

· 3.

Why is the formation of a solid evidence of a chemical reaction? Use a molecular-level drawing in your explanation.

· 4.

Sketch molecular-level drawings to differentiate between two soluble compounds: one that is a strong electrolyte, and one that is not an electrolyte.

· 5.

Mixing an aqueous solution of potassium nitrate with an aqueous solution of sodium chloride does not result in a chemical reaction. Why?

· 6.

Why is the formation of water evidence of a chemical reaction? Use a molecular-level drawing in your explanation.

· 7.

Use the Arrhenius definition of acids and bases to write the net ionic equation for the reaction of an acid with a base.

· 8.

Why is the transfer of electrons evidence of a chemical reaction? Use a molecular-level drawing in your explanation.

· 9.

Why is the formation of a gas evidence of a chemical reaction? Use a molecular-level drawing in your explanation.

· 10.

Label each of the following statements as true or false. Explain your answers, and provide an example for each that supports your answer.

a. All nonelectrolytes are insoluble.

b. All insoluble substances are nonelectrolytes.

c. All strong electrolytes are soluble.

d. All soluble substances are strong electrolytes.

· 11.

Look at Fig. 7.2 in the text. It is possible for a weak electrolyte solution to cause the bulb to glow brighter than a strong electrolyte. Explain how this is possible.

· 12.

What is the purpose of spectator ions? If they are not present as part of the reaction, why are they present at all?

· 13.

Which of the following must be an oxidation—reduction reaction? Explain your answer, and include an example oxidation—reduction reaction for all that apply.

a. A metal reacts with a nonmetal.

b. A precipitation reaction.

c. An acid—base reaction.

· 14.

If an element is a reactant or product in a chemical reaction, the reaction must be an oxidation—reduction reaction. Why is this true?

· 15.

Match each name below with the following microscopic pictures of that compound in aqueous solution.

a. barium nitrate

b. sodium chloride

c. potassium carbonate

d. magnesium sulfate

Which picture best represents ? Why aren’t any of the pictures a good representation of ?

· 16.

On the basis of the general solubility rules given in Table 7.1, predict the identity of the precipitate that forms when aqueous solutions of the following substances are mixed. If no precipitate is likely, indicate which rules apply.

· 17.

Write the balanced formula and net ionic equation for the reaction that occurs when the contents of the two beakers are added together. What colors represent the spectator ions in each reaction?

Questions and Problems: 7.1 Predicting Whether a Reaction Will Occur

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

Questions

· 1.

Why is water an important solvent? Although you have not yet studied water in detail, can you think of some properties of water that make it so important?

· 2.

What is a “driving force”? What are some of the driving forces discussed in this section that tend to make reactions likely to occur? Can you think of any other possible driving forces?

Questions and Problems: 7.2 Reactions in Which a Solid Forms

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

Questions

· 3.

A reaction in aqueous solution that results in the formation of a solid is called a reaction.

© Cengage Learning

· 4.

When writing the chemical equation for a reaction, how do you indicate that a given reactant is dissolved in water? How do you indicate that a precipitate has formed as a result of the reaction?

· 5.

Describe briefly what happens when an ionic substance is dissolved in water.

· 6.

When the ionic solute is dissolved in water, what can you say about the number of potassium ions present in the solution compared with the number of phosphate ions in the solution?

· 7.

What is meant by a strong electrolyte? Give two examples of substances that behave in solution as strong electrolytes.

· 8.

How do chemists know that the ions behave independently of one another when an ionic solid is dissolved in water?

· 9.

Suppose you are trying to help your friend understand the general solubility rules for ionic substances in water. Explain in general terms to your friend what the solubility rules mean, and give an example of how the rules could be applied in determining the identity of the precipitate in a reaction between solutions of two ionic compounds.

· 10.

Using the general solubility rules given in Table 7.1, which of the following ions will form a precipitate with ?

a.

b.

c.

d. At least two of the above ions will form a precipitate with .

e. All of the above ions will form a precipitate with .

· 11.

On the basis of the general solubility rules given in Table 7.1, predict which of the following substances are not likely to be soluble in water. Indicate which specific rule(s) led to your conclusion.

a.

b.

c.

d.

e.

f.

g.

h.

· 12.

On the basis of the general solubility rules given in Table 7.1, predict which of the following substances are likely to be appreciably soluble in water. Indicate which specific rule(s) led to your conclusion.

a.

b.

c.

d.

e.

f.

g.

h.

· 13.

On the basis of the general solubility rules given in Table 7.1, for each of the following compounds, explain why the compound would be expected to be appreciably soluble in water. Indicate which of the solubility rules covers each substance’s particular situation.

a. potassium sulfide

b. cobalt(III) nitrate

c. ammonium phosphate

d. cesium sulfate

e. strontium chloride

· 14.

On the basis of the general solubility rules given in Table 7.1, for each of the following compounds, explain why the compound would not be expected to be appreciably soluble in water. Indicate which of the solubility rules covers each substance’s particular situation.

a. iron(III) hydroxide

b. calcium carbonate

c. cobalt(III) phosphate

d. silver chloride

e. barium sulfate

· 15.

On the basis of the general solubility rules given in Table 7.1, predict the identity of the precipitate that forms when aqueous solutions of the following substances are mixed. If no precipitate is likely, indicate which rules apply.

a. copper(II) chloride, , and ammonium sulfide,

b. barium nitrate, , and potassium phosphate,

c. silver acetate, , and calcium chloride,

d. potassium carbonate, , and cobalt(II) chloride,

e. sulfuric acid, , and calcium nitrate,

f. mercurous acetate, , and hydrochloric acid,

· 16.

Lead(II) nitrate is added to four separate beakers that contain the following:

Beaker 1 (sodium chloride)

Beaker 2 (sodium hydroxide)

Beaker 3 (sodium phosphate)

Beaker 4 (sodium sulfate)

After the addition of the lead(II) nitrate solution to each beaker, in which beaker(s) will a precipitate form? Use the general solubility rules given in Table 7.1 to guide you.

Problems

· 17.

On the basis of the general solubility rules given in Table 7.1, write a balanced molecular equation for the precipitation reactions that take place when the following aqueous solutions are mixed. Underline the formula of the precipitate (solid) that forms. If no precipitation reaction is likely for the reactants given, explain why.

a. ammonium chloride, , and sulfuric acid,

b. potassium carbonate, , and tin(IV) chloride,

c. ammonium chloride, , and lead(II) nitrate,

d. copper(II) sulfate, , and potassium hydroxide,

e. sodium phosphate, , and chromium(III) chloride,

f. ammonium sulfide, , and iron(III) chloride,

· 18.

On the basis of the general solubility rules given in Table 7.1, write a balanced molecular equation for the precipitation reactions that take place when the following aqueous solutions are mixed. Underline the formula of the precipitate (solid) that forms. If no precipitation reaction is likely for the solutes given, so indicate.

a. sodium carbonate, , and copper(II) sulfate,

b. hydrochloric acid, , and silver acetate,

c. barium chloride, , and calcium nitrate,

d. ammonium sulfide, , and iron(III) chloride,

e. sulfuric acid, , and lead(II) nitrate,

f. potassium phosphate, , and calcium chloride,

· 19.

Balance each of the following equations that describe precipitation reactions.

a.

b.

c.

· 20.

Balance each of the following equations that describe precipitation reactions.

a.

b.

c.

· 21.

For each of the following precipitation reactions, complete and balance the equation, indicating clearly which product is the precipitate. If no reaction would be expected, so indicate.

a.

b.

c.

· 22.

A solution of zinc nitrate is mixed with a solution of potassium hydroxide. A precipitate forms. Complete and balance the equation for this reaction, including the phases of each reactant and product.

Questions and Problems: 7.3 Describing Reactions in Aqueous Solutions

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

Questions

· 23.

What is a net ionic equation? What species are shown in such an equation, and which species are not shown?

· 24.

Which of the following most accurately describes a spectator ion?

a. An ion that is used up in a chemical reaction; it is limiting.

b. An ion that participates in a chemical reaction but is always present in excess

c. An ion that becomes part of the precipitate in a chemical reaction

d. An ion that does not have a charge but can dissolve in solution and thus does not conduct electricity

e. An ion that is present in solution but does not participate directly in the chemical reaction

Problems

· 25.

Based on the general solubility rules given in Table 7.1, propose five combinations of aqueous ionic reagents that likely would form a precipitate when they are mixed. Write the balanced full molecular equation and the balanced net ionic equation for each of your choices.

· 26.

Write the balanced molecular, complete ionic, and net ionic equations for the reaction between nickel(II) chloride and sodium sulfide.

· 27.

Many chromate salts are insoluble, and most have brilliant colors that have led to their being used as pigments. Write balanced net ionic equations for the reactions of , , , and with chromate ion.

· 28.

The procedures and principles of qualitative analysis are covered in many introductory chemistry laboratory courses. In qualitative analysis, students learn to analyze mixtures of the common positive and negative ions, separating and confirming the presence of the particular ions in the mixture. One of the first steps in such an analysis is to treat the mixture with hydrochloric acid, which precipitates and removes silver ion, lead(II) ion, and mercury(I) ion from the aqueous mixture as the insoluble chloride salts. Write balanced net ionic equations for the precipitation reactions of these three cations with chloride ion.

· 29.

Many plants are poisonous because their stems and leaves contain oxalic acid, , or sodium oxalate, ; when ingested, these substances cause swelling of the respiratory tract and suffocation. A standard analysis for determining the amount of oxalate ion, , in a sample is to precipitate this species as calcium oxalate, which is insoluble in water. Write the net ionic equation for the reaction between sodium oxalate and calcium chloride, , in aqueous solution.

· 30.

Another step in the qualitative analysis of cations (see Exercise 28) involves precipitating some of the metal ions as the insoluble sulfides (followed by subsequent treatment of the mixed sulfide precipitate to separate the individual ions). Write balanced net ionic equations for the reactions of , , , and ions with sulfide ion, .

Questions and Problems: 7.4 Reactions That Form Water: Acids and Bases

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

Questions

· 31.

What is meant by a strong acid? Are the strong acids also strong electrolytes? Explain.

· 32.

What is meant by a strong base? Are the strong bases also strong electrolytes? Explain.

· 33.

The same net ionic process takes place when any strong acid reacts with any strong base. Write the equation for that process.

· 34.

Write the formulas and names of three common strong acids and strong bases.

· 35.

If units were dissolved in a sample of water, the would produce ions and ions.

· 36.

What is a salt? Give two balanced chemical equations showing how a salt is formed when an acid reacts with a base.

Problems

· 37.

Write balanced equations showing how three of the common strong acids ionize to produce hydrogen ion.

· 38.

Along with the three strong acids emphasized in the chapter ( , , and ), hydrobromic acid, , and perchloric acid, , are also strong acids. Write equations for the dissociation of each of these additional strong acids in water.

· 39.

What salt would form when each of the following strong acid/strong base reactions takes place?

a.

b.

c.

d.

· 40.

Complete the following acid—base reactions by indicating the acid and base that must have reacted in each case to produce the indicated salt.

a.

b.

c.

d.

Questions and Problems: 7.5 Reactions of Metals with Nonmetals (Oxidation—Reduction)

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

Questions

· 41.

What is an oxidation—reduction reaction? What is transferred during such a reaction?

· 42.

Give an example of a simple chemical reaction that involves the transfer of electrons from a metallic element to a nonmetallic element.

· 43.

What do we mean when we say that the transfer of electrons can be the “driving force” for a reaction? Give an example of a reaction where this happens.

· 44.

The thermite reaction produces so much energy (heat) that the iron is initially formed as a liquid.

© Cengage Learning

Describe the transfer of electrons for both the aluminum and iron.

· 45.

If atoms of the metal calcium were to react with molecules of the nonmetal fluorine, , how many electrons would each calcium atom lose? How many electrons would each fluorine atom gain? How many calcium atoms would be needed to react with one fluorine molecule? What charges would the resulting calcium and fluoride ions have?

· 46.

If oxygen molecules, , were to react with magnesium atoms, how many electrons would each magnesium atom lose? How many electrons would each oxygen atom gain? How many magnesium atoms would be needed to react with each oxygen molecule? What charges would the resulting magnesium and oxide ions have?

Problems

· 47.

For the reaction , illustrate how electrons are gained and lost during the reaction.

· 48.

For the reaction , show how electrons are gained and lost by the atoms.

· 49.

Balance each of the following oxidation—reduction reactions.

a.

b.

c.

d.

· 50.

Balance each of the following oxidation—reduction chemical reactions.

a.

b.

c.

d.

Questions and Problems: 7.6 Ways to Classify Reactions

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

Questions

directs you to the Chemistry in Focus feature in the chapter

· 51.

a. Give two examples each of a single-displacement reaction and of a double-replacement reaction. How are the two reaction types similar, and how are they different?

b. Give two examples each of a reaction in which formation of water is the driving force and in which formation of a gas is the driving force.

· 52.

The reaction between ammonium perchlorate and aluminum is discussed in the “Chemistry in Focus” segment Oxidation—Reduction Reactions Launch the Space Shuttle. The reaction is labeled as an oxidation—reduction reaction. Explain why this is an oxidation—reduction reaction and defend your answer.

· 53.

Identify each of the following unbalanced reaction equations as belonging to one or more of the following categories: precipitation, acid—base, or oxidation—reduction.

a.

b.

c.

d.

e.

f.

g.

h.

i.

· 54.

Identify each of the following unbalanced reaction equations as belonging to one or more of the following categories: precipitation, acid—base, or oxidation—reduction.

a.

b.

c.

d.

e.

f.

g.

h.

i.

Questions and Problems: 7.7 Other Ways to Classify Reactions

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

Questions

· 55.

How do we define a combustion reaction? In addition to the chemical products, what other products do combustion reactions produce? Give two examples of balanced chemical equations for combustion reactions.

· 56.

Reactions involving the combustion of fuel substances make up a subclass of reactions.

· 57.

What is a synthesis or combination reaction? Give an example. Can such reactions also be classified in other ways? Give an example of a synthesis reaction that is also a combustion reaction. Give an example of a synthesis reaction that is also an oxidation—reduction reaction, but that does not involve combustion.

· 58.

What is a decomposition reaction? Give an example. Can such reactions also be classified in other ways?

Problems

· 59.

Complete and balance each of the following combustion reactions.

a.

b.

c.

· 60.

Complete and balance each of the following combustion reactions.

a.

b.

c.

· 61.

By now, you are familiar with enough chemical compounds to begin to write your own chemical reaction equations. Write two examples of what we mean by a combustion reaction.

· 62.

By now, you are familiar with enough chemical compounds to begin to write your own chemical reaction equations. Write two examples each of what we mean by a synthesis reaction and by a decomposition reaction.

· 63.

Balance each of the following equations that describe synthesis reactions.

a.

b.

c.

· 64.

Balance each of the following equations that describe synthesis reactions.

a.

b.

c.

· 65.

Balance each of the following equations that describe decomposition reactions.

a.

b.

c.

d.

e.

· 66.

Balance each of the following equations that describe oxidation—reduction reactions.

a.

b.

c.

d.

e.

Additional Problems

· 67.

Distinguish between the molecular equation, the complete ionic equation, and the net ionic equation for a reaction in solution. Which type of equation most clearly shows the species that actually react with one another?

· 68.

Which of the following ions form compounds with that are generally soluble in water?

a.

b.

c.

d.

e.

· 69.

Without first writing a full molecular or ionic equation, write the net ionic equations for any precipitation reactions that occur when aqueous solutions of the following compounds are mixed. If no reaction occurs, so indicate.

a. iron(III) nitrate and sodium carbonate

b. mercurous nitrate and sodium chloride

c. sodium nitrate and ruthenium nitrate

d. copper(II) sulfate and sodium sulfide

e. lithium chloride and lead(II) nitrate

f. calcium nitrate and lithium carbonate

g. gold(III) chloride and sodium hydroxide

· 70.

Complete and balance each of the following molecular equations for strong acid/strong base reactions. Underline the formula of the salt produced in each reaction.

a.

b.

c.

d.

· 71.

For the cations listed in the left-hand column, give the formulas of the precipitates that would form with each of the anions in the right-hand column. If no precipitate is expected for a particular combination, so indicate.

· 72.

Which of the following correctly identifies the precipitate formed when aqueous solutions of the following substances are mixed?

· 73.

On the basis of the general solubility rules given in Table 7.1, predict the identity of the precipitate that forms when aqueous solutions of the following substances are mixed. If no precipitate is likely, indicate why (which rules apply).

a. iron(III) chloride and sodium hydroxide

b. nickel(II) nitrate and ammonium sulfide

c. silver nitrate and potassium chloride

d. sodium carbonate and barium nitrate

e. potassium chloride and mercury(I) nitrate

f. barium nitrate and sulfuric acid

· 74.

Below are indicated the formulas of some salts. Such salts could be formed by the reaction of the appropriate strong acid and strong base (with the other product of the reaction being, of course, water). For each salt, write an equation showing the formation of the salt from reaction of the appropriate strong acid and strong base.

a.

b.

c.

d.

· 75.

For each of the following unbalanced molecular equations, write the corresponding balanced net ionic equation for the reaction.

a.

b.

c.

d.

· 76.

Write the balanced molecular, complete ionic, and net ionic equations for the reaction of ammonium sulfide with iron(III) chloride.

· 77.

What strong acid and what strong base would react in aqueous solution to produce the following salts?

a. potassium perchlorate,

b. cesium nitrate,

c. potassium chloride,

d. sodium sulfate,

· 78.

For the reaction , show how electrons are gained and lost by the atoms.

© Cengage Learning

· 79.

For the reaction , show how electrons are gained and lost by the atoms.

· 80.

Balance the equation for each of the following oxidation—reduction chemical reactions.

a.

b.

c.

d.

e.

· 81.

Identify each of the following unbalanced reaction equations as belonging to one or more of the following categories: precipitation, acid—base, or oxidation—reduction.

a.

b.

c.

d.

e.

f.

g.

h.

i.

· 82.

Which of the following statements is/are true regarding solutions?

a. If a solute is dissolved in water, then the resulting solution is considered aqueous.

b. If two solutions are mixed and no chemical reaction occurs, then a net ionic equation cannot be written.

c. If two clear solutions are mixed and then cloudiness results, this indicates that a precipitate formed.

· 83.

Balance each of the following equations that describe synthesis reactions.

a.

b.

c.

d.

e.

· 84.

Balance each of the following equations that describe decomposition reactions.

a.

b.

c.

d.

e.

· 85.

Write a balanced oxidation—reduction equation for the reaction of each of the metals in the left-hand column with each of the nonmetals in the right-hand column.

· 86.

Fill in the following table as if it is a well plate and you are mixing two aqueous compounds at a time to see if a precipitate forms. If a precipitate is expected to form, indicate that by writing the correct formula for the precipitate in the corresponding box in the table. If no precipitate is expected to form, write “No” in the box.

· 87.

Although the metals of Group 2 of the periodic table are not nearly as reactive as those of Group 1, many of the Group 2 metals will combine with common nonmetals, especially at elevated temperatures. Write balanced chemical equations for the reactions of , , , and with , , and .

· 88.

For each of the following metals, how many electrons will the metal atoms lose when the metal reacts with a nonmetal?

a. sodium

b. potassium

c. magnesium

d. barium

e. aluminum

· 89.

For each of the following nonmetals, how many electrons will each atom of the nonmetal gain in reacting with a metal?

a. oxygen

b. fluorine

c. nitrogen

d. chlorine

e. sulfur

· 90.

True or false? When solutions of barium hydroxide and sulfuric acid are mixed, the net ionic equation is: because only the species involved in making the precipitate are included. Whether true or false, include a balanced molecular equation and complete ionic equation for the reaction between barium hydroxide and sulfuric acid to support your answer.

o True

o False

· 91.

Classify the reactions represented by the following unbalanced equations by as many methods as possible. Balance the equations.

a.

b.

c.

d.

e.

· 92.

When a sodium chromate solution and aluminum bromide solution are mixed, a precipitate forms. Complete and balance the equation for this reaction, including the phases of each reactant and product.

· 93.

Corrosion of metals costs us billions of dollars annually, slowly destroying cars, bridges, and buildings. Corrosion of a metal involves the oxidation of the metal by the oxygen in the air, typically in the presence of moisture. Write a balanced equation for the reaction of each of the following metals with : , , , , and .

· 94.

Consider a solution with the following ions present:

All are allowed to react, and there are plenty available of each. List all of the solids that will form using the correct formulas in your explanation.

· 95.

Give a balanced molecular chemical equation to illustrate each of the following types of reactions.

a. a synthesis (combination) reaction

b. a precipitation reaction

c. a double-displacement reaction

d. an acid—base reaction

e. an oxidation—reduction reaction

f. a combustion reaction

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.

· 96.

For the following chemical reactions, determine the precipitate produced when the two reactants listed below are mixed together. Indicate “none” if no precipitate will form.

· 97.

For the following chemical reactions, determine the precipitate produced when the two reactants listed below are mixed together. Indicate “none” if no precipitate will form.

Chapters 6-7. Cumulative Review

Questions

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

· 1.

What kind of visual evidence indicates that a chemical reaction has occurred? Give an example of each type of evidence you have mentioned. Do all reactions produce visual evidence that they have taken place?

· 2.

What, in general terms, does a chemical equation indicate? What are the substances indicated to the left of the arrow called in a chemical equation? To the right of the arrow?

· 3.

What does it mean to “balance” an equation? Why is it so important that equations be balanced? What does it mean to say that atoms must be conserved in a balanced chemical equation? How are the physical states of reactants and products indicated when writing chemical equations?

· 4.

When balancing a chemical equation, why is it not permissible to adjust the subscripts in the formulas of the reactants and products? What would changing the subscripts within a formula do? What do the coefficients in a balanced chemical equation represent? Why is it acceptable to adjust a substance’s coefficient but not permissible to adjust the subscripts within the substance’s formula?

· 5.

What is meant by the driving force for a reaction? Give some examples of driving forces that make reactants tend to form products. Write a balanced chemical equation illustrating each type of driving force you have named.

· 6.

Explain to your friend what chemists mean by a precipitation reaction. What is the driving force in a precipitation reaction? Using the information provided about solubility in these chapters, write balanced molecular and net ionic equations for five examples of precipitation reactions.

· 7.

Define the term strong electrolyte. What types of substances tend to be strong electrolytes? What does a solution of a strong electrolyte contain? Give a way to determine if a substance is a strong electrolyte.

· 8.

Summarize the simple solubility rules for ionic compounds. How do we use these rules in determining the identity of the solid formed in a precipitation reaction? Give examples including balanced complete and net ionic equations.

· 9.

In general terms, what are the spectator ions in a precipitation reaction? Why are the spectator ions not included in writing the net ionic equation for a precipitation reaction? Does this mean that the spectator ions do not have to be present in the solution?

· 10.

Describe some physical and chemical properties of acids and bases. What is meant by a strong acid or base? Are strong acids and bases also strong electrolytes? Give several examples of strong acids and strong bases.

· 11.

What is a salt? How are salts formed by acid—base reactions? Write chemical equations showing the formation of three different salts. What other product is formed when an aqueous acid reacts with an aqueous base? Write the net ionic equation for the formation of this substance.

· 12.

What do we call reactions in which electrons are transferred between atoms or ions? What do we call a loss of electrons by an atom or ion? What is it called when an atom or ion gains electrons? Can we have a process in which electrons are lost by one species without there also being a process in which the electrons are gained by another species? Why? Give three examples of equations in which there is a transfer of electrons between a metallic element and a nonmetallic element. In your examples, identify which species loses electrons and which species gains electrons.

· 13.

What is a combustion reaction? Are combustion reactions a unique type of reaction, or are they a special case of a more general type of reaction? Write an equation that illustrates a combustion reaction.

· 14.

Give an example of a synthesis reaction and of a decomposition reaction. Are synthesis and decomposition reactions always also oxidation—reduction reactions? Explain.

· 15.

List and define all the ways of classifying chemical reactions that have been discussed in the text. Give a balanced chemical equation as an example of each type of reaction, and show clearly how your example fits the definition you have given.

Problems

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

· 16.

The element carbon undergoes many inorganic reactions, as well as being the basis for the field of organic chemistry. Write balanced chemical equations for the reactions of carbon described below.

a. Carbon burns in an excess of oxygen (for example, in the air) to produce carbon dioxide.

b. If the supply of oxygen is limited, carbon will still burn but will produce carbon monoxide rather than carbon dioxide.

c. If molten lithium metal is treated with carbon, lithium carbide, , is produced.

d. Iron(II) oxide reacts with carbon above temperatures of about to produce carbon monoxide gas and molten elemental iron.

e. Carbon reacts with fluorine gas at high temperatures to make carbon tetrafluoride.

· 17.

Balance each of the following chemical equations.

a.

b.

c.

d.

e.

f.

g.

h.

· 18.

The reagent shelf in a general chemistry lab contains aqueous solutions of the following substances: silver nitrate, sodium chloride, acetic acid, nitric acid, sulfuric acid, potassium chromate, barium nitrate, phosphoric acid, hydrochloric acid, lead nitrate, sodium hydroxide, and sodium carbonate. Suggest how you might prepare the following pure substances using these reagents and any normal laboratory equipment. If it is not possible to prepare a substance using these reagents, indicate why.

a.

b.

c.

d.

e.

f.

· 19.

The common strong acids are , , and , whereas and are the common strong bases. Write the neutralization reaction equations for each of these strong acids with each of these strong bases in aqueous solution.

· 20.

Balance each equation. Which equations can be classified as oxidation—reduction reactions?

a.

b.

c.

d.

e.

f.

g.

h.

i.

j.

· 21.

Consider the oxidation—reduction reaction . Why is this chemical equation also classified as a synthesis reaction?

· 22.

Give balanced equations for two examples of each of the following types of reactions.

a. precipitation

b. single-displacement

c. combustion

d. synthesis

e. oxidation—reduction

f. decomposition

g. acid—base neutralization

· 23.

Using the general solubility rules discussed in Chapter 7, give the formulas of five substances that would be expected to be readily soluble in water and five substances that would be expected to not be very soluble in water. For each of the substances you choose, indicate the specific solubility rule you applied to make your prediction.

· 24.

Write the balanced net ionic equation for the reaction that takes place when aqueous solutions of the following solutes are mixed. If no reaction is likely, explain why no reaction would be expected for that combination of solutes.

a. potassium nitrate and sodium chloride

b. calcium nitrate and sulfuric acid

c. ammonium sulfide and lead(II) nitrate

d. sodium carbonate and iron(III) chloride

e. mercurous nitrate and calcium chloride

f. silver acetate and potassium chloride

g. phosphoric acid and calcium nitrate

h. sulfuric acid and nickel(II) sulfate

· 25.

Complete and balance the following equations.

a.

b.

c.

d.

e.

f.

g.

h.