bishop 5 cs - faculty · diprotic acid an acid that can donate two hydrogen ions ... such a...

CHAPTER 5Chemical Reactions

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Objectives

You will be able to do the following.

1. Given a description of a solution of two components, identify the solute and the solvent.2. Write a description of the process for dissolving an ionic compound in water. Your

description should include mention of the nature of the particles in solution and the attractions between the particles in the solution.

3. Given the formula for an ionic compound, predict whether it is soluble in water or not.4. Given formulas for two ionic compounds, (1) predict whether a precipitate will form when

the water solutions of the two are mixed, (2) if there is a reaction, predict the products of the reaction and write their formulas, and (3) if there is a reaction, write the complete, complete ionic, and net ionic equations that describe the reaction.

5. Describe precipitation reactions. Your descriptions should include mention of the nature of the particles in the system before and after the reaction, a description of the cause of the reaction, and a description of the attractions between the particles before and after the reaction.

6. Write a description of the process for the solution of a strong monoprotic acid, sulfuric acid, and a weak monoprotic acid in water.

7. Identify the following as strong monoprotic acids: HCl, HBr, HI, HNO3, and HClO4.8. Identify sulfuric acid, H2SO4, as a strong diprotic acid.9. Given a formula for any acid, identify it as a strong or weak acid.10. Write a description of the process for the solution of ammonia, NH3, in water.11. Given the formula or name of a compound, classify it as a strong electrolyte, weak

electrolyte, or nonelectrolyte. 12. Given a formula for an ionic compound, identify it as representing a weak base, a strong

base, or a substance that is neutral in the Arrhenius acid-base sense. 13. Identify the following as anions that are neutral in the acid-base sense: Cl−, Br−, I−,

NO3−, and ClO4

−.14. Identify the following anions as weak acids: hydrogen sulfate ion, HSO4

−, and dihydrogen phosphate ion, H2PO4

−.15. Identify ammonia as an uncharged, weak base.16. Identify ionic compounds containing hydroxide ions as strong bases.17. Given the pH of a solution, identify the solution as acidic, basic, or neutral. 18. Given the pH of two acidic solutions, identify which solution is more acidic. 19. Given the pH of two basic solutions, identify which solution is more basic.20. Describe how litmus paper can be used in the laboratory to identify whether a solution is

acidic or basic.21. Write a description of the process that takes place at the molecular level when a strong acid

reacts with an aqueous strong base. Your descriptions should include mention of the nature of the particles in the system before and after the reaction, a description of the cause of the reaction, and a description of the attractions between the particles before and after the reaction.

74 Chapter 5 Chemical Reactions

22. Given the formulas for an acid and a base, predict the products that would form from the reaction between them, and write the balanced complete, complete ionic, and net ionic equations that describe the reaction.

23. Identify H2O(l ) and CO2(g) as the products of the reaction of an acid with carbonate, CO3

2−, or hydrogen carbonate, HCO3−.

24. Write an explanation for why a substance can be a Brønsted-Lowry acid in one reaction and a Brønsted-Lowry base in a different reaction. Give an example to illustrate your explanation.

25. Write an explanation for why the Arrhenius definitions for acid and base and not the Brønsted-Lowry definitions are used to describe whether an isolated substance is an acid or base.

26. Write an explanation for why it is useful to have two sets of definitions for acids and bases (the Arrhenius definitions and the Brønsted-Lowry definitions).

27. Given a Brønsted-Lowry neutralization equation, identify the Brønsted-Lowry acid and Brønsted-Lowry base.

28. Given an equation for a chemical reaction, (1) identify whether the equation represents a redox reaction or not, (2) identify the substance that is oxidized and the substance that is reduced, and (3) identify the substance that is the reducing agent and the substance that is the oxidizing agent.

29. Given the mass of substance dissolved in a solution and the total volume of the solution, calculate the molarity of the solution.

30. Convert between moles of solute and volume of solution using molarity as a conversion factor.

31. Convert between amount of one substance and amount of another substance, both involved in a chemical reaction, when the amounts of these substances are described by (1) mass of pure substance, (2) mass of a mixture that contains one of the substances in the reaction with the percentage of that substance described, or (3) volume of a solution that contains one of the substances with the molarity of the solution given.

32. Given the volume of a substance titrated, the molarity of the titrant used in the titration, and the volume of titrant necessary to reach the equivalence point in the titration, calculate the molarity of the substance titrated.

33. Given a volume and molarity for a solution, write a description of how the solution could be made from pure solid and water.

34. Write a description of why when making a solution from solid and water, the solid must be dissolved in water before more water is added to bring the volume to the desired total.

35. Write an explanation for why pure or almost pure acid should always be added to water and not water to pure or almost pure acid.

36. Given the volume and molarity of a dilute solution and the molarity of a more concentrated solution used to make the dilute solution, calculate the volume of the concentrated solution necessary to make the dilute solution.

37. Write a description of the procedure for making a more dilute solution of an acid from a pure or almost pure acid (e.g. 18 M H2SO4).

38. Given three of the following four, calculate the fourth: volume of a more dilute solution, molarity of the more dilute solution, volume of a more concentrated solution of the same substance, and molarity of the more concentrated solution.

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39. Given the volume and molarity of a more dilute solution of a substance and the molarity of a more concentrated solution of the same substance (other than a pure or almost pure acid) used to make the dilute solution, write a description of the procedure for making the more dilute solution from the more concentrated solution.

40. Convert between the definition and the term for the following words or phrases.

Skip Section 5.5 of the text.

Chapter 5 GlossarySolution A mixture whose particles are so evenly distributed that the relative

concentrations of the components are the same throughout. Solutions can also be called homogeneous mixtures.

Aqueous solution A solution in which water is the solvent. Solute The gas in a solution of a gas in a liquid. The solid in a solution of a solid in a

liquid. The minor component in other solutions.Solvent The liquid in a solution of a gas in a liquid. The liquid in a solution of a solid

in a liquid. The major component in other solutions. Hydrated Bound to one or more water molecules. Hydration The binding of one or more water molecules to an ion or molecule.Electrolyte A substance that ionizes or dissociates in water to form an electrically

conducting solution. Strong electrolyte A substance that ionizes or dissociates completely in an aqueous

solution.Nonelectrolyte A substance that ionizes or dissociates incompletely in an aqueous

solution.Ionize To form ions (often as a substance dissolves in water).Dissociate The separation of ions (often as a substance dissolves in water).Precipitation reaction A reaction in which one of the products is insoluble in water

and comes out of solution as a solid. Precipitate A solid that comes out of solution.Precipitation The process of forming a solid in a solution. Crystals Solid particles whose component atoms, ions, or molecules are arranged in

an organized, repeating pattern.Complete ionic equation A chemical equation that describes the actual form for

each substance in solution. For example, ionic compounds that are dissolved in water are described as separate ions.

Spectator ions Ions that play a role in delivering other ions into solution to react but that do not actively participate in the reaction themselves.

Complete equation or molecular equation A chemical equation that includes uncharged formulas for all of the reactants and products. The formulas include the spectator ions, if any.

Net ionic equation A chemical equation for which the spectator ions have been eliminated, leaving only the substances actively involved in the reaction.

Solubility The maximum amount of solute that can be dissolved in a given amount of solvent.

Hydronium ion H3O+


Arrhenius acid According to the Arrhenius theory, any substance that generates hydronium ions, H3O+, when added to water.

Monoprotic acid An acid that donates one hydrogen ion per molecule in a reaction. Polyprotic acid An acid that can donate more than one hydrogen ion per molecule in a

reaction.Diprotic acid An acid that can donate two hydrogen ions per molecule in a reaction.Triprotic acid An acid that can donate three hydrogen ions per molecule in a reaction. Strong acid An acid that donates its H+ ions to water in a reaction that goes completely

to products. Such a compound produces close to one H3O+ ion in solution for each acid molecule dissolved in water.

Completion reaction A reaction that shifts completely to products, that is, a reaction that is not significantly reversible.

Reversible reaction A reaction in which the reactants are constantly forming products and, at the same time, the products are reforming the reactants.

Weak acid A substance that is incompletely ionized in water due to the reversibility of the reaction that forms hydronium ions, H3O+, in water. Weak acids yield significantly less than one H3O+ ion in solution for each acid molecule dissolved in water.

Strong monoprotic acids HCl(aq), HBr(aq), HI(aq), HNO3, HClO4. Strong diprotic acids H2SO4. Arrhenius base A substance that produces hydroxide ions, OH−, when added to water.Strong base A substance that generates at least one hydroxide ion in solution for every

unit of substance added to water. Weak base A substance that produces fewer hydroxide ions in water solution than particles

of the substance added. Miscible Able to be mixed in any proportion, that is, infinitely soluble.Acidic solution A solution with a significant concentration of hydronium ions, H3O+.Basic solution A solution with a significant concentration of hydroxide ions, OH−. Neutralization reaction A chemical reaction between an acid and a base. Brønsted-Lowry acid-base reaction A chemical reaction in which a proton, H+, is

transferred. Brønsted-Lowry Acid A substance that donates protons, H+, in a Brønsted-Lowry

acid-base reaction.Brønsted-Lowry Base A substance that accepts protons, H+, in a Brønsted-Lowry

acid-base reaction. Amphoteric substance A substance that can act as either a Brønsted-Lowry acid and a

Brønsted-Lowry base, depending on the circumstances.Oxidation Any chemical change in which at least one element loses electrons, either

completely or partially. Reduction Any chemical change in which at least one element gains electrons, either

completely or partially. Oxidation-reduction reactions The chemical reactions in which there is a complete or

partial transfer of electrons, resulting in oxidation and reduction. These reactions are also called redox reactions.

Half-reactions Separate oxidation and reduction reaction equations in which electrons are shown as a reactant or product.

Reducing agent A substance that loses electrons, making it possible for another substance to gain electrons and be reduced.

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Oxidizing agent A substance that gains electrons, making it possible for another substance to lose electrons and be oxidized.

Oxidation number A tool for keeping track of the flow of electrons in redox reactions (also called oxidation state).

Figure 5.1 Sodium Chloride Dissolving in Water This shows the mixture immediately after sodium chloride has been added to water. Certain water molecules are highlighted to draw attention to their role in the process.


Figure 5.2 Aqueous Sodium Chloride This image shows a portion of the solution that forms when sodium chloride dissolves in water. Certain water molecules are highlighted to draw attention to their role in the process.

Figure 5.3 Aqueous Calcium Nitrate Notice that there are twice as many −1 nitrate ions as +2 calcium ions.

You can find an animation that shows the process by which sodium chloride dissolves at the following web address:

http://www.mpcfaculty.net/mark_bishop/NaCl_dissolves.htm

http://www.mpcfaculty.net/mark_bishop/NaCl_dissolves.htm

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Figure 5.4 Mixture of Ca(NO3)2(aq) and Na2CO3(aq) at the Instant They Are Combined

Figure 5.5 Product Mixture for the Ca(NO3)2(aq) and Na2CO3(aq) Reaction


You can find an animation that shows this precipitation reaction at the following web address:

http://www.mpcfaculty.net/mark_bishop/precipitation.htm

Category Ions Except with these ions Examples

Soluble Cations

Group 1A metal ions and ammonium, NH4

+All soluble Na2CO3, LiOH, and (NH4)2S are

soluble

Soluble Anions

NO3− and C2H3O2

−, All soluble Bi(NO3)3, and Co(C2H3O2)2 are soluble

Usually Soluble Anions

Cl−, Br−, and I− Soluble except with Ag+ and Pb2+

CuCl2 is water soluble, but AgCl is insoluble.

SO42− Soluble except with Ba2+

and Pb2+FeSO4 is water soluble, but BaSO4 is insoluble.

Usually Insoluble Anions

CO32−, SO3

2−, PO43−,

and OH− Insoluble except with group

1A elements and NH4+

CaCO3, ZnSO3, Ca3(PO4)2, and Mn(OH)2 are insoluble in water,

but (NH4)2CO3, K2SO3, Li3PO4, and CsOH are soluble.

S2− Insoluble except with group 1A and 2A metal ions and

NH4+

CoS is insoluble in water, but MgS is soluble.

EXERCISE 5.1 - Ionic solubility

Predict whether each of the following are soluble or insoluble in water. a. Hg(NO3)2

b. FeCO3

c. SnS

d. K3PO4

e. PbCl2

f. Al(OH)3

Table 5.1 Ionic Solubility

http://www.mpcfaculty.net/mark_bishop/precipitation.htm

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Sample Study Sheet 5.1: Predicting Precipitation Reactions and Writing Precipitation Equations

TIP-OFF You are asked to predict whether a precipitation reaction will take place between two aqueous solutions of ionic compounds, and if the answer is yes, to write the complete equation for the reaction.

GENERAL STEPS STEP 1 Determine the formulas for the possible products using the general double-displacement equation. (Remember to consider ion charges when writing your formulas.)

AB + CD → AD + CB

STEP 2 Predict whether either of the possible products is water insoluble. If either possible product is insoluble, a precipitation reaction takes place, and you may continue with Step 3. If neither is insoluble, write “No reaction”.

STEP 3 Follow these steps to write the complete equation.

• Write the formulas for the reactants separated by a “+”.

• Separate the formulas for the reactants and products with a single arrow.

• Write the formulas for the products separated by a “+”.

• Write the physical state for each formula.

The insoluble product will be followed by (s).

Water-soluble ionic compounds will be followed by (aq).

• Balance the equation.

EXERCISE 5.2 - Precipitation Reactions

Predict whether a precipitate will form when each of the following pairs of water

solutions are mixed. If there is a precipitation reaction, write the complete, complete

ionic, and net ionic equations that describe the reaction.

a. Li2CO3(aq) + Al(NO3)3(aq)

b. KOH(aq) + Fe(NO3)3(aq)

c. NaC2H3O2(aq) + CaS(aq)

d. K2SO4(aq) + Pb(NO3)2(aq)


You can find a description of the procedure for writing complete and net ionic equations at the following web address:

http://www.mpcfaculty.net/mark_bishop/precipitation_equations.htm

Having Trouble?

Are you having trouble with this chapter? People often do. To successfully complete each of the tasks in this chapter, you need to have mastered the skills from previous sections. If you missed anything along the way, you will have trouble. Here is a list of the tasks you need to be able to do in order to work the problems in this chapter. You should go through the list in order and be sure you have mastered each skill before you go on to the next one.

1. Convert between names and symbols for the common elements.

2. Identify whether an element is a metal or a nonmetal.

3. Determine the charges on many of the monatomic ions.

4. Convert between the name and formula for polyatomic ions.

5. Recognize a name or formula as ionic or molecular.

6. Recognize a name or formula as representing a binary acid or an oxyacid.

7. Convert between the name and formula for ionic compounds, binary acids and oxyacids.

8. Balance chemical equations.

9. Predict the products of double displacement reactions.

10. Predict ionic solubility.

11. Predict the states of ionic compounds and H2O.

12. Identify strong, weak, and nonelectrolytes.

13. Write complete ionic and net ionic equations.



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Figure 5.6 Hydrochloric Acid in Water

Figure 5.7 Acetic Acid in Water


You can recognize strong and weak acids by using the following criteria. ♦ You can recognize acids either from their formula or their name.

∗ Acid formulas can be either HX(aq) like HCl(aq), H2S(aq), or HaXbOc, such as HNO2.

∗ Acid names end in acid, such as hydrochloric acid, hydrosulfuric acid and nitrous acid.

♦ The strong monoprotic acids you should know are HCl, HBr, HI, HNO3 and HClO4.

♦ Sulfuric acid, H2SO4, is a strong diprotic acid.

You will recognize a substance as a weak acid if it is not on the list of strong acids.

Figure 5.8 Ammonia in Water

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Table 5.2 Classifying Arrhenius acids

Classification of Arrhenius Acids

Weak or Strong

Characteristics How to Recognize

MonoproticStrong Completely

ionized in waterMemorize HCl(aq), HBr(aq), HI(aq), HNO3, and HClO4

Weak Incompletely ionized in water

General formula of HX(aq) or HXbOc and not on the list of strong acids*

DiproticStrong First H+ ion

lost completely: second incompletely

H2SO4

Weak Both H+ ions lost incompletely

H2S(aq) or General formula of H2XbOc (not H2SO4)

Triprotic Weak All three H+ ions lost incompletely

General formula H3XbOc

*Some acids have formulas that do not fit our general formulas. Acetic acid, HC2H3O2,

is the only example that you need to recognize at this point.

Table 5.3 Classifying Arrhenius bases

Arrhenius Bases Strong or Weak

Characteristics How recognize?

Anions strong Each unit added leads to at least one OH− ion in solution

Soluble Metal hydroxides*

(in ionic compounds)

weak Reversible reaction with water to yield fewer OH− ions than units added

Anions in ionic compounds except with OH−, Cl−, Br−, I−, NO3

−, ClO4

−, HSO4−, and H2PO4

−

Some uncharged molecules

weak Reversible reaction with water to yield fewer OH− ions than molecules added

NH3

* Note: Since a metal hydroxide must be soluble in water to yield a significant

concentration of hydroxide in solution, it is common to consider only water soluble

metal hydroxides like NaOH to be strong bases. Insoluble metal hydroxides like

aluminum hydroxide, Al(OH)3, react like the soluble hydroxides in neutralization

reactions, so for some purposes they can be considered strong bases.


Sample Study Sheet 5.2: Identification of Strong And Weak Acids and Bases

TIP-OFF You can use the following steps to identify a name or chemical formula as representing either (1) an Arrhenius strong acid, (2) an Arrhenius weak acid, (3) an Arrhenius strong base, (4) an Arrhenius weak base, or (5) not acidic or basic in the Arrhenius sense (neutral).

GENERAL STEPS

STEP 1 Identify the substance as an acid, a base, or neither. You can identify acids in the following ways.

• If you are given a name: a. the names of the uncharged acids end in acid.

For example, hydrochloric acid is an acid. b. the names for the only ionic compounds you are expected to

recognize as acidic end in hydrogen sulfate or dihydrogen phosphate. Sodium hydrogen sulfate is acidic.

• If you are given a molecular formula:a. Remember that acids have one of these forms: HX(aq) or HaXbOc.HCl(aq) and H2SO4 are acids. b. Acidic ionic compounds have formulas that include HSO4

− or H2PO4

−. NaH2PO4 is acidic.

You can identify bases in the following ways. • We expect ionic compounds to be basic, except those containing the

following anions. a. Cl−, Br−, I−, NO3

−, ClO4− - neutral

b. HSO4−, H2PO4

− - acidicNaF is basic, NaCl is neutral, and NaHSO4 is acidic.

• Ammonia, NH3, is a base. STEP 2 If you have an acid or base, determine whether it is strong or weak.

♦ We will consider all acids except HCl(aq), HBr(aq), HI(aq), HNO3, HClO4, and H2SO4 to be weak.

HF is a weak acid. ♦ We will consider all bases except metal hydroxides to be weak.

NaF is a weak base.

EXERCISE 5.3 - Acid and Base Classification

Identify HNO2, lithium hydroxide, NaCN, sodium iodide, NaHSO4, nitric acid, CH3OH, hydrofluoric acid, and KC2H3O2 as either (1) an Arrhenius strong acid, (2) an Arrhenius weak acid, (3) an Arrhenius strong base, (4) an Arrhenius weak base, or (5) not acidic or basic in the Arrhenius sense (neutral).

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The following are characteristics of acids.1. Acids have a sour taste.

2. Acids turn litmus from blue to red.

3. Acids react with bases. (When the base includes carbonate or hydrogen

carbonate ions, carbon dioxide gas is released.)

The following are characteristics of bases.1. Bases have a bitter taste.

2. Bases feel slippery on your fingers.

3. Bases turn litmus from red to blue.

4. Bases react with acids.

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Figure 5.9 pH of Common Substances

Table 5.4 Strong, weak and nonelectrolytes

Ions in solution?

Completely ionized?

Types

Strong Electrolytes Yes Yes Strong acids Water soluble ionic compounds

Weak Electrolytes Yes No Uncharged weak acids Ammonia

Nonelectrolytes No Not ionized at all Alcohols Sugars


Table 5.5 Identification of Strong, Weak, and Nonelectrolytes

Category General Type of Substance

How to Recognize Formula

How to Recognize Name Examples

Strong electrolytes

Ionic(aq) MaAb MaHbXcOd or (NH4)aAb (NH4)aHbXcOd

metal (root nonmetal)ide, metal polyatomic ion, ammonium (root nonmetal)ide, or ammonium polyatomic ion

NaCl or Li2HPO4 or NaNO3 or NH4NO3

Strong acids

HCl(aq), HBr(aq), HI(aq), HNO3, HClO4, and H2SO4

hydrochloric acid, hydrobromic acid, hydroiodic acid, nitric acid, perchloric acid, and sulfuric acid

All are listed to the left.

Weak electrolytes Uncharged weak acids

HX(aq), H2S(aq), or HaXbOc

and not on the list of strong acids

hydrosulfuric acid per(root)ic acid(root)ic acid(root)ous acidhypo(root)ous acid and not on the list of strong acids

hydrosulfuric acidperiodic acidchloric acidchlorous acidhypochlorous acid

NH3 NH3 ammonia

Nonelectrolytes Alcohol CaHbOH name ends in -anol or alcohol

methanol or methyl alcohol

Sugar C6H12O6 or C12H22O11

name ends in -ose glucose or sucrose

M=symbol of metal, A=symbol of nonmetal, X=some element other than H or O, and a, b, c & d indicate subscripts

EXERCISE 5.4 - Strong, Weak, and Nonelectrolytes

Identify Al(NO3)3, acetic acid, NH3, ammonium acetate, HCl, glucose, CH3OH, barium chloride, K2Cr2O7, nitric acid, HBrO2, 2–propanol (or isopropyl alcohol) as either a (1) strong electrolyte, (2) weak electrolyte, or (3) nonelectrolyte, and in parentheses write the type of compound each name or formula represents. Each dissolves in water.

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Figure 5.11 Water Solution of Nitric Acid and Sodium Hydroxide before Reaction

Figure 5.10 Aqueous Nitric Acid


Figure 5.12 After Reaction of Nitric Acid and Sodium Hydroxide

Table 5.6 Prediction of whether neutralization reactions are completion reactions or reversible reactions

Strong Acid Weak Acid

Strong Base Completion Completion

Weak Base Completion Reversible or Completion*

*When you need to know whether a neutralization reaction is a completion reaction or reversible, you will be told.

Sample Study Sheet 5.3:

Steps for Writing Neutralization Equations

TIP-OFF You will be given formulas or names of two substances and asked to predict whether a neutralization reaction will take place between them. If it does, you will be asked to write the complete, complete ionic, and net ionic equations for the reaction. The following steps allow you to do this.GENERAL STEPS

STEP 1 Ask, “Do you have an acid and a base?” If yes, go to Step 2. If no, say “No reaction”. See Previous Study Sheets.

Any one of the following conditions will lead you to predict “no reaction”. 1) One of the substances is neutral in the Arrhenius acid-base sense.2) Both substances given are acids. 3) Both substances given are bases.

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STEP 2 Describe a completion reaction with a single arrow, (→). Describe a reversible reaction with a double arrow, ( ). Determine whether the reaction is reversible or goes to completion by applying the following guidelines.

a. When one or both of the acid and base are strong, the reaction is a completion reaction. See Previous Study Sheets.

b. If you have a weak acid and a weak base, you can assume that the reaction is reversible.

STEP 3 Write the formulas and states of the products.a. You must first decide on the correct form of the equation.

1) Most neutralization reactions that you will see in this chapter will be double displacement. If the base is ammonia, NH3, the reaction will have the second form.

Double Displacement AB + CD → AD + CB

NH3(aq) + HX(aq) → NH4X(aq)

2) For double displacement reactions, follow these steps.

a. Identify the A, B, C and D.

(1) For acids the positive component is H+.

(2) Split ionic compounds into cation and anion.

b. Write the AD and CB formulas. Be careful to balance charges.

c. Remember that carbonic acid, H2CO3(aq), decomposes to form CO2(g) and H2O(l ).

d. We will assume that all of the H+ ions react for polyprotic acids, e.g. H2SO3 loses two H+ ions, and H3PO4 loses three H+ ions.

b. Write the states of reactants and products.

1. The states are usually given for the original reactants.

2. For the ionic products:

1). Most ionic products of neutralization reactions are water soluble and described with (aq).

2). If they are insoluble, describe them with (s).

3. Describe acids as aqueous, (aq), unless stated otherwise.

4. Ammonia is aqueous, (aq), when in water and gaseous, (g), when pure.

5. Carbon dioxide is insoluble in water and described as a gas, (g).

6. Water is a liquid, (l ).

STEP 4 Balance the complete equation.


STEP 5 Write the complete ionic and net ionic equations.

a. To get the complete ionic equation, describe each reactant and product as ions

or with a complete (not ionized) formula.

1. Strong acids and ionic(aq) are strong electrolytes and completely ionized.

a. For this task, describe strong acids, HA(aq), as H+(aq) and A−(aq).

(The A−(aq) represents an anion)

b. Even though describing strong acids, HA(aq), as H3O+(aq) and

A−(aq) is a better description of the nature of the ions in solution, it

is probably more common to use H+(aq) and A−(aq) for the strong

acids.

2. Everything else is described with a complete (not ionized) formula.

b. To get the net ionic equation, eliminate spectator ions. Remember that

these are the ions that are in an identical form on both sides of the equation.

Rewrite what is left and balance.

c. You can check that you have written the net ionic equation correctly by

making sure the following are true.

1. Strong monoprotic acids are described as H+(aq) in a net ionic equation.

(Even though H3O+(aq) provides a better description of the ions in

solution, it is probably more common to use H+(aq).)

2. The strong diprotic acid sulfuric acid, H2SO4, is described as H+ and

HSO4−.

3. Weak, uncharged acids, like acetic acid, HC2H3O2(aq), are described

with complete formulas.

a. A very common mistake is to describe the weak acids like the strong

acids, as H+ in the net ionic equation.

b. Ionic formulas that contain the anions that are weak acids are

described as ions. For instance, NaHSO4(aq) is described as

Na+(aq) and HSO4− (aq).

4. Any ions in strong electrolytes on both sides of the equation are

spectator ions and should be eliminated in a net ionic equation.

5. Formulas with (g), (l ), and (s) are described with complete formulas.

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EXERCISE 5.5 - Writing Neutralization Equations

For each of the following pairs of possible reactants, predict whether a neutralization reaction will take place between them. If there is no reaction, write, “No Reaction”. If there is a reaction, write complete, complete ionic, and net ionic equations for the reaction. (The reactions between weak acids and weak bases given here are reversible reactions. If an acid has more than one acidic hydrogen, assume that there is enough base to remove all of them. Assume that there is enough acid to add as many protons to the base as possible)

a. HCl(aq) + NaOH(aq)

b. HF(aq) + LiOH(aq)

c. NH3(aq) + HNO3(aq)

d. NH3(aq) + HClO(aq)

e. HC2H3O2(aq) + LiF(aq)

f. Na2CO3(aq) + HBr(aq)

g. HCl(aq) + HNO2(aq)

h. H3PO3(aq) + LiOH(aq)

i. Fe(OH)3(s) + HNO3(aq)

j. NaI(aq) + HCl(aq)


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Figure 5.13 Brønsted-Lowry Conjugate Acid-Base Pairs

EXERCISE 5.6 - Brønsted-Lowry Acids and Bases

For each of the following equations, identify the Brønsted-Lowry acid and base.

a. HNO2(aq) + NaBrO(aq) → HBrO(aq) + NaNO2(aq)

b. H2PO4−(aq) + HNO2(aq) H3PO4(aq) + NO2

−(aq)

c. H2PO4−(aq) + 2OH−(aq) → PO4

3−(aq) + 2H2O(l)

d. H2SO3(aq) + 2NaOH(aq) → Na2SO3(aq) + 2H2O(l)

Table 5.7 Key definitions for Redox Reactions

Term Definition

oxidation-reduction reaction

An electron transfer reaction

oxidation Loss of electrons

reduction Gain of electrons

reducing agent The substance that loses electrons (is oxidized) and makes it possible for something else to gain electrons (be reduced)

oxidizing agent The substance that gains electrons (is reduced) and makes it possible for something else to lose electrons (be oxidized)

half-reaction One-half of a redox reaction; just the oxidation or just the reduction

95


Assignment of Oxidation Numbers

TIP-OFF You are asked to determine the oxidation number of an atom, or you need to assign oxidation numbers to atoms to determine whether a reaction is a redox reaction, and if it is, to identify which element is oxidized, which is reduced, what the oxidizing agent is, and what the reducing agent is.

GENERAL STEPS Use the following guidelines to assign oxidation numbers to as many atoms as you can. (The following table provides a summary of these guidelines with examples.)

• The oxidation number for each atom in a pure element is zero. • The oxidation number of a monatomic ion is equal to its charge. • When fluorine atoms are combined with atoms of other elements, their

oxidation number is −1. • When oxygen atoms are combined with atoms of other elements, their

oxidation number is −2, except in peroxides, like hydrogen peroxide, H2O2, where their oxidation number is −1.

• The oxidation number for each hydrogen atom in a molecular compound or a polyatomic ion is +1.

If a compound’s formula contains one element for which you cannot assign an oxidation number using the guidelines listed above, calculate the oxidation number according to the following rules.

• The sum of the oxidation numbers for the atoms in an uncharged formula is equal to zero.

• The sum of the oxidation numbers for the atoms in a polyatomic ion is equal to the overall charge on the ion.

Table 5.8 Oxidation Numbers for Some Elements

Oxidation Number

Examples Exceptions

Pure, uncharged element

0 Each atom is 0 in Zn, H2, P4, and Cl2

None

Monatomic ions charge on ion

Cd in CdCl2 is +2. Cl in CdCl2 is −1.

H in LiH is −1.

None

Fluorine in the combined form

−1 F in AlF3 is −1. F in CF4 is −1.

None

Oxygen in the combined form

−2 O in ZnO is −2. O in H2O is −2.

O is −1 in peroxides, such as H2O2 and O2

2−.

Covalently bonded hydrogen

+1 H in H2O is +1. None


Equations for redox reactions can be difficult to balance, but your ability to determine oxidation numbers can help. You can find a description of the process for balancing redox equations at the following web address.

www.mpcfaculty.net/mark_bishop/redox_balancing.htm

EXERCISE 5.7 - Oxidation Numbers

Determine the oxidation number for the atoms in P4, PF3, PH3, P2O3, H3PO4, N2,

N3−, K3N, Co3O2, NaH, Na, HSO3−, Cu(NO3)2, K2O2, and Fe2(SO4)3.

Table 5.9 Questions Answered by the Determination of Oxidation Numbers

Question How to answer the question

Is the reaction redox? If any atoms change their oxidation number, the reaction is redox.

What’s oxidized? The element that increases its oxidation number is oxidized.

What’s reduced? The element that decreases its oxidation number is reduced.

What’s the reducing agent? The substance that contains the element that is oxidized is the reducing agent.

What’s the oxidizing agent? The substance that contains the element that is reduced is the oxidizing agent.

www.mpcfaculty.net/mark_bishop/redox_balancing.htm

97

EXERCISE 5.8 - Redox Reactions

Identify whether the following equations describe redox reactions or not. For each of the redox reactions, identify what is oxidized, what is reduced, what the reducing agent is, and what the oxidizing agent is.

a. Ca(s) + F2(g) → CaF2(s)

∆b. CaCO3(s) → CaO(s) + CO2(g)

c. 2Al(s) + 3H2O(g) → Al2O3(s) + 3H2(g)

d. Cr2O72−(aq) + 6Cl−(aq) + 14H+(aq) → 2Cr3+(aq) + 3Cl2(g) + 7H2O(l )

Figure 5.14 Single-Displacement Reaction Between Copper(II) Sulfate and Solid Zinc

98 Chapter 5 Relating Quantities of Reactants and Products


General Equation Stoichiometry

TIP-OFF - You are asked to convert from amount of one substance in a chemical reaction to amount of another substance in the reaction.

GENERAL PROCEDURE - Use the unit analysis format to make the following conversions.

1. If you are not given it, write and balance the chemical equation for the reaction.

2. Start your unit analysis in the usual way, setting the desired units of substance 2 equal to the given units of substance 1.

3. Convert from the units that you are given for substance 1 to moles of substance 1.

• For pure solids and liquids, this means converting grams to moles using the molar mass of the substance. (It might be necessary to insert one or more additional conversion factors to convert from the given unit to grams.)

• Molarity can be used to convert from volume of solution to moles of solute. (It might be necessary to insert one or more additional conversion factors to convert from the given unit to liters or milliliters.)

4. Convert from moles of substance 1 to moles of substance 2 using the coefficients from the balanced equation.

5. Convert from moles of substance 2 to the desired units for substance 2. • For pure solids and liquids, this means converting moles to mass using

the molar mass of substance 2. • Molarity can be used to convert from moles of solute to volume of

solution. 6. Calculate your answer and report it with the correct significant figures (in

scientific notation, if necessary) and with the correct unit.

EXERCISE 5.9 - Stoichiometry and Molarity

How many milliliters of 6.00 M HNO3 are necessary to neutralize the carbonate in 75.0 mL of 0.250 M Na2CO3?

EXERCISE 5.10 - Molarity

What is the molarity of a AgClO4 solution made by dissolving 29.993 g of silver perchlorate in water and diluting with water to 50.00 mL total?

99

Figure 5.15 General Equation Stoichiometry

Visit the following website to see a description of a procedure called titration that can be used to determine the molarities of acidic and basic solutions:

www.mpcfaculty.net/mark_bishop/titration.htm

EXERCISE 5.11 - General Stoichiometry

What is the maximum number of grams of silver chloride that will precipitate from a solution made by mixing 25.00 mL of 0.050 M MgCl2 with an excess a AgNO3 solution?

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www.mpcfaculty.net/mark_bishop/titration.htm



Titration Problems

TIP-OFF - For the problems that you will see in this text, you will be given the volume of a solution of an acid or base (the titrant – solution 1) necessary to react completely with a given volume of solution being titrated (solution 2). You will also be given the molarity of the titrant (solution 1). You will be asked to calculate the molarity of solution 2. GENERAL PROCEDURE - You can use the following steps. Use the unit analysis process, with the following general format.

∗ The first conversion factor is only necessary if you are not given liters of 2. (Because you are usually given milliliters, this conversion factor often converts from milliliters to liters.)

∗ The second conversion factor is only necessary if you are not given either milliliters or liters of solution 1. (You are usually given milliliters, so if you use the form of the molarity conversion factor that includes “103 mL 1 soln”, this conversion factor is not necessary.)

∗ The coefficients in the final conversion factor come from the balanced equation for the reaction.

∗ Complete the calculation in the usual way.

(given) (volume unit) 1

(given) (volume unit) 2

--- (voloume unit) 2

--- L

--- L (or mL)

--- (volume unit) 1

(molarity) mol 1

1 L (or 103 mL) 1 soln

(coef. 2) mol 2

(coef. 1) mol 1

? mol 2

1 L 2 soln=

EXERCISE 5.12 - Titration Problem

When 34.2 mL of a 1.02 M NaOH solution is added from a burette to 25.00 mL of a phosphoric acid solution that contains phenolphthalein, the solution changes from colorless to red.

a. What is the titrant for this process?

b. What is the molarity of the phosphoric acid?


Making Solutions from Pure, Solid Substance

101

TIP-OFF - You are asked to describe how you would make an aqueous solution from a pure solid, and you are given the volume and the molarity of solution desired. GENERAL PROCEDURE – Follow these steps.

• Calculate the mass in grams of solid to be weighed.

The first conversion factor is only necessary if you are not given liters or milliliters of X solution.

• Answer the question with a sentence like the following.

Dissolve (calculated) grams of X in a minimum amount of water, and dilute with water to the desired total volume.

? g X = (given) (volume unit) X soln--- L (or --- mL)

--- (volume unit)

(molarity) mol X

1 L (or 103 mL) X soln

(molar mass) g X

1 mol X

EXERCISE 5.13 - Making Solution from Solid

An experiment calls for a total of 1.50 L of 0.200 M KMnO4 for a class of chemistry students. How would this solution be made from pure, solid potassium permanganate and water?



Dilution Problems

TIP-OFF - There are several tip-offs for this type of problem. Perhaps the most general is the mention of two concentrations of the same substance. Other tip-offs include mention of dilution or the specific reference to making a more dilute solution from a more concentrated solution.

GENERAL PROCEDURE - There are two possible approaches. UNIT ANALYSIS - You can use the general unit analysis thought-process, taking care to clearly identify the units of concentrated solution, dilute solution, and pure solute. The molarities given in the problem are used as conversion factors. The following is one general form of this approach.

WITH DILUTION EQUATION - You can also use the following steps.

• Assign the variables MC, VC, MD, and VD to the values you are given and the value that you want.

• Write the dilution equation.

MCVC = MDVD

• Solve the equation for the unknown variable.

• Plug in the values with their units for the other variables.

• Cancel your units.

• If your units do not cancel to yield the desired unit, make the necessary unit conversions so the units do cancel correctly.

• Complete the calculation, round your answer to the correct significant figures, and report the correct unit.

? mL conc soln = (given) mL dil soln soln(dil molarity) mol solute

103 mL dil X soln

103 mL conc X

(conc molarity) mol solute

EXERCISE 5.14 - Dilution Problems

What is the molarity of a solution made by diluting 5.00 mL of a 14.8 M NH3 solution to 75.0 mL?

103

Sample Study Sheet 5.9: Making Solutions

Tip-off - You are asked to describe how to make a solution of a specific molarity.

General Procedure - The first step is to recognize whether you are starting with pure solid or with a more concentrated solution.

• If you are starting with pure solid, use the procedure described on the study sheet called Descriptions of Making Solutions From Solid and Water.

• If you are starting with a more concentrated solution, follow these steps. (In this text, you will be given the volume and molarity desired for the more dilute solution, and you will be given the molarity of the more concentrated starting solution.)

∗ Calculate the volume of the more concentrated solution that must be measured to start the process. This can be done with unit analysis or with the shortcut for dilution problems. (See the sample study sheet for Dilution Problems.)

MCVC = MDVD VC =MDVD

MC

∗ If the more concentrated solution is a pure or almostpure acid, such as 18 M H2SO4 or 17 M HC2H3O2, describe the process for making the solution in the following way:

♦ Add the (calculated) volume of the concentrated acid solution to enough water to dilute the acid significantly.

♦ Add water until the volume reaches the desired total.∗ If the more concentrated solution is not a pure or almostpure acid,

describe the process for making the solution in the following way: ♦ Add the (calculated) volume of the more concentrated solution to a

volume-measuring instrument. ♦ Add water until the volume reaches the desired total.

? mL conc soln = (given) mL dil soln soln(dil molarity) mol solute

103 mL dil X soln

103 mL conc X

(conc molarity) mol solute

EXERCISE 5.15 - Making Solution from Concentrated Acid

How would you make 250.0 milliliters of 2.00 M acetic acid from concentrated acetic acid (called glacial acetic acid) that is 17.4 M HC2H3O2?


EXERCISE 5.16 - Making Solution from More Concentrated Solution

How would you make 50.0 milliliters of 0.250 M hydrochloric acid from 2.0 M HCl?

You can get more information about making solutions of a specific molarity at the

following Web address:

http://www.mpcfaculty.net/mark_bishop/making_solutions.htm

http://www.mpcfaculty.net/mark_bishop/making_solutions.htm

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