19.1 Acid-Base Theories 19t1lara.weebly.com/uploads/1/6/3/2/1632178/ch19pdf.pdf · Lowry theory?...

Acids, Bases, and Salts 587

19.1

FOCUSObjectives19.1.1 Define the properties of acids

and bases.19.1.2 Compare and contrast acids

and bases as defined by the theories of Arrhenius, Brønsted-Lowry, and Lewis.

Guide for Reading

Build VocabularyWord Parts Explain that the prefixes mono, di, and tri have Greek origins, meaning one, two, and three, respec-tively. Have students think of other words that use these prefixes. (monop-oly, monolog; dipole; triangle, triarchy)

Reading StrategyPredict Have students make a two-col-umn chart with the headings Acid and Base. Have them list substances they think contain either acids or bases. Ask students how they identified each sub-stance as an acid or base.

INSTRUCT

Have students study the photograph and read the text that opens the sec-tion. Ask, Why do visitors to Bracken Cave in Texas need to wear protec-tive goggles and respirators? (Stu-dents will likely suggest that ammonia has a strong, pungent odor and that high levels of ammonia are dangerous.)

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Section Resources

Connecting to Your World

Section 19.1 Acid-Base Theories 587

19.1 Acid-Base Theories

Bracken Cave, near San Anto-nio, Texas, is home to twenty to forty million bats, which is probably the largest colony of mammals in the world. Visitors to the cave must wear

protective goggles and respirators to protect themselves from the dangerous levels of

ammonia in the cave. Ammonia is a by-product of the bats’ urine. In this section, you will learn why ammonia is considered

a base.

Guide for Reading

Key Concepts• What are the properties of acids

and bases?• How did Arrhenius define an

acid and a base?• What distinguishes an acid

from a base in the Brønsted-Lowry theory?

• How did Lewis define an acid and a base?

Vocabularymonoprotic acidsdiprotic acidstriprotic acidsconjugate acidconjugate baseconjugate acid-base pairhydronium ion (H3O�)amphotericLewis acidLewis base

Reading StrategyBuilding Vocabulary As youread the section, write a definition and the formula of an example of each type of substance listed in the key terms list.

Properties of Acids and BasesAcids and bases play a central role in much of the chemistry that affectsyour daily life. Your body needs acids and bases to function properly. Vine-gar, carbonated drinks, and foods such as citrus fruits contain acids. Theelectrolyte in a car battery is an acid. Most manufacturing processes useacids or bases. Bases are present in many commercial products, includingantacids and household cleaning agents. Figure 19.1 shows some of themany products that contain acids and bases.

Acids Acids have several distinctive properties with which you are proba-bly familiar. Acidic compounds give foods a tart or sour taste. For example,vinegar imparts a tart taste to salad dressing. Vinegar contains ethanoicacid, sometimes called acetic acid. Lemons, which taste sour enough tomake your mouth pucker, contain citric acid.

Aqueous solutions of acids are electrolytes. Recall from Chapter 15 thatelectrolytes conduct electricity. Some acid solutions are strong electrolytes,and others are weak electrolytes. Acids cause certain chemical dyes, calledindicators, to change color. Many metals, such as zinc and magnesium,react with aqueous solutions of acids to produce hydrogen gas. Acids reactwith compounds containing hydroxide ions to form water and a salt.

Acids taste sour, will change the color of an acid-base indicator, and canbe strong or weak electrolytes in aqueous solution.

Figure 19.1 Many items contain acids or bases, or produce acids and bases when dissolved in water.

Citrus fruit contain citric acid. Tea contains tannic acid. Antacids use bases to neutralize excess stomach acid.

The base calcium hydroxide is a component of mortar.

abc

d

b c da

Print• Guided Reading and Study Workbook,

Section 19.1• Core Teaching Resources, Section 19.1

Review• Transparencies, T213–T214

Technology• Interactive Textbook with ChemASAP,

Animation 25, Problem-Solving 19.1, Assessment 19.1

• Go Online, Section 19.1

588 Chapter 19

Section 19.1 (continued)

Properties of Acids and Bases

TEACHER DemoTEACHER Demo

Reactive AcidsPurpose Demonstrate that many met-als and acids react to produce hydro-gen gas.

Materials 5–10 mL of 1M HCl, test tube, small pieces of zinc, match

Safety Wear safety goggles, gloves, and a lab apron while performing the demonstration.

Procedure Add 5–10 mL 1M HCl to a test tube and add a few small pieces of zinc metal.

Expected Outcome Gas bubbles are liberated from the solution. Explain that many metals react with aqueous acids to produce hydrogen gas. Remind students that hydrogen is a highly flammable gas. Demonstrate the liberation of the gas by placing a lit match at the mouth of the test tube to ignite the hydrogen.

Arrhenius Acids and BasesUse VisualsTable 19.1 Note that the acids in the table are listed in order of decreasing tendency to yield hydrogen ions. Ask, What element do all the acids in this table have in common? (hydrogen) Do diprotic acids always produce more hydrogen ions in solution than monoprotic acids? (They may, but there is no direct relationship between the number of hydrogen atoms in a for-mula and the number of hydrogen ions produced.)

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Acid Names and Uses Acids can often be purchased at hardware and home supply stores, but they are often labeled with “old names.” For example, muri-atic acid, a term no longer used in chemistry,

is hydrochloric acid. Among other uses, this acid is purchased to clean lime deposits from tile and porcelain and to clean mortar stains from brick and tile.

Facts and Figures

588 Chapter 19

Bases Bases have properties with which you are less familiar. Bases have abitter taste, but tasting most bases is hazardous. Soap is a familiar materialthat has the properties of a base. If you have tasted soap, you know that ithas a bitter taste. The slippery feel of soap is another property of bases.

Like acids, aqueous solutions of bases are electrolytes and will cause anindicator to change color. Water and a salt are formed when a base thatcontains hydroxide ions reacts with an acid. With a few exceptions, none ofthe foods you eat are bases. Milk of magnesia (a suspension of magnesiumhydroxide in water) is a base used to treat the problem of excess stomachacid. Bases taste bitter, feel slippery, will change the color of an acid-baseindicator, and can be strong or weak electrolytes in aqueous solution.

Checkpoint Give an example of a base.

Arrhenius Acids and BasesAlthough chemists had recognized the properties of acids and bases formany years, they were not able to propose a theory to explain this behavior.Then, in 1887, the Swedish chemist Svante Arrhenius (1859–1927) pro-posed a revolutionary way of defining and thinking about acids and bases.

Arrhenius said that acids are hydrogen-containing compounds that ion-ize to yield hydrogen ions (H�) in aqueous solution. He also said that bases arecompounds that ionize to yield hydroxide ions (OH�) in aqueous solution.

Arrhenius Acids Table 19.1 lists some common acids. Acids that containone ionizable hydrogen, such as nitric acid (HNO3), are called monoprotic

acids. Acids that contain two ionizable hydrogens, such as sulfuric acid(H2SO4), are called diprotic acids. Acids that contain three ionizable hydro-

gens, such as phosphoric acid (H3PO4), are called triprotic acids. Not all

compounds that contain hydrogen are acids, however. Also, not all thehydrogens in an acid may be released as hydrogen ions. Only the hydrogensin very polar bonds are ionizable. In such bonds, hydrogen is joined to avery electronegative element. When a compound that contains such bondsdissolves in water, it releases hydrogen ions because the hydrogen ions arestabilized by solvation. An example is the hydrogen chloride molecule,shown in Figure 19.2. Hydrogen chloride is a polar covalent molecule. Itionizes to form an aqueous solution of hydronium ions and chloride ions.

Hydrogen Hydrogen Chloridechloride ion ion

(hydrochloric acid)

Hd+

¬Cld-

1g 2¡H2O

H+1aq 2 + Cl-1aq 2

Figure 19.2 Hydrochloric acid is actually an aqueous solution of hydrogen chloride. The hydrogen ion forms hydronium ions, making this compound an acid.

Table 19.1

Some Common Acids

� —→� �

�

HClHydrogen chloride

H2OWater

H3O�

Hydronium ionCl�

Chloride ion

Name Formula

Hydrochloric acid HCl

Nitric acid HNO3

Sulfuric acid H2SO4

Phosphoric acid H3PO4

Ethanoic acid CH3COOH

Carbonic acid H2CO3


Use VisualsTable 19.2 Have students examine the table. Ask them to examine the metal cations in each of the bases, noting that the bases containing Group 1A elements appear to be highly soluble in water while those containing Group 2A elements are not. Ask, Based on your generalizations, is LiOH likely to be soluble in water? (LiOH is proba-bly soluble in water.)

Explain that hydroxides involving Group 2A elements increase in solubility going down the periodic table column. Ask, Do you think Be(OH)2 is soluble in water? (insoluble) Sr(OH)2? (quite solu-ble)

DiscussAcid–base concepts have been a part of chemistry for more than 300 years. Over time, theories have been devel-oped to explain the observed behavior of acids and bases. Arrhenius’s theory addresses compounds whose formulas contain obvious clues to their acidic or basic nature.

Download a worksheet on Acids and Bases in the Home for students to complete, and find additional teacher support from NSTA SciLinks.

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Answers to...

Checkpoint

Ammonia, soap, and milk of magnesia are examples of bases.


Table 19.2

Some Common Bases

For: Links on Acids and Bases in the home

Visit: www.SciLinks.orgWeb Code: cdn-1191

In contrast, the four hydrogens in methane (CH4) are attached by weaklypolar C—H bonds. Methane has no ionizable hydrogens and is not an acid.Ethanoic (acetic) acid (CH3COOH), used in the manufacture of plastics,pharmaceuticals, and photographic chemicals, is different. Although thismolecule contains four hydrogens, ethanoic acid is a monoprotic acid. Thestructural formula shows why.

The three hydrogens attached to the carbon are in weakly polar bonds. Theydo not ionize. Only the hydrogen bonded to the highly electronegative oxygencan be ionized. As you gain more experience looking at written formulas foracids, you will be able to recognize which hydrogen atoms can be ionized.

Arrhenius Bases Table 19.2 lists some common bases. The base withwhich you are perhaps most familiar is sodium hydroxide (NaOH). Sodiumhydroxide is commonly known as lye. Sodium hydroxide is an ionic solid. Itdissociates into sodium ions and hydroxide ions in aqueous solution.

Because of its extremely caustic nature, sodium hydroxide is a major com-ponent of consumer products used to clean clogged drains.

Potassium hydroxide (KOH) is another ionic solid that dissociates toform potassium ions and hydroxide ions in aqueous solution.

Sodium and potassium are Group 1A elements. The elements inGroup 1A, the alkali metals, react with water to produce solutions that arebasic. Sodium metal reacts violently with water to form sodium hydroxideand hydrogen gas. The following equation illustrates the reaction ofsodium metal with water.

Ethanoic acid(CH3COOH)

H¬

Hƒ

Cƒ

H

¬

O‘C¬O¬H

Sodium Sodium Hydroxidehydroxide ion ion

NaOH 1 s 2¡H2O

Na+ 1aq 2 + OH- 1aq 2

KOH(s) K�(aq) + OH�(aq)Potassium Potassium Hydroxidehydroxide ion ion

¡H2O

Sodium Water Sodium Hydrogenmetal hydroxide

2Na 1s 2 + 2H2O 1 l 2¡ 2NaOH 1aq 2 + H2 1g 2

Name Formula Solubility in water

Potassium hydroxide KOH High

Sodium hydroxide NaOH High

Calcium hydroxide Ca(OH)2 Very low

Magnesium hydroxide Mg(OH)2 Very low

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Brønsted-Lowry Acids and BasesDiscussBrønsted arrived at his theory of acids and bases through his work in kinetics and thermodynamics. Brønsted noticed that a great many compounds that would not be classified as acids and bases by the Arrhenius definition behaved like acids and bases in his experiments. He found that these com-pounds were capable of hydrogen-ion transfer reactions. This led him to define acids as hydrogen-ion donors and bases as hydrogen-ion acceptors in chemical reactions. Some students may think the conjugate base of an acid in a given reaction appears on the same side of the equation as the acid. Explain that an acid does not react with its conjugate base but instead produces it or is produced by it in a reaction.

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Sodium hydroxide and potassium hydroxide are very soluble in water.Concentrated solutions of these compounds can be readily prepared. Suchsolutions, like other basic solutions, would have a bitter taste and slipperyfeel. However, they are extremely caustic to the skin and can cause deep,painful, slow-healing wounds if not immediately washed off.

Calcium hydroxide (Ca(OH)2) and magnesium hydroxide (Mg(OH)2)are compounds of Group 2A metals. These compounds are not very solublein water. Their solutions are always very dilute, even when saturated. Theconcentration of hydroxide ions in such solutions is correspondingly low. Asaturated solution of calcium hydroxide contains only 0.165 g Ca(OH)2 per100 g of water. Magnesium hydroxide is much less soluble than calciumhydroxide. A saturated solution contains only 0.0009 g Mg(OH)2 per 100 gof water. Suspensions of magnesium hydroxide in water contain low con-centrations of hydroxide ion. People take these suspensions internallyas milk of magnesia, shown in Figure 19.3, which is an antacid and a mildlaxative.

Brønsted-Lowry Acids and BasesThe Arrhenius definition of acids and bases is not a very comprehensiveone. It defines acids and bases rather narrowly and does not include certainsubstances that have acidic or basic properties. For example, aqueous solu-tions of sodium carbonate (Na2CO3(aq)) and ammonia (NH3(aq)) are basic.Neither of these compounds is a hydroxide, however, and neither would beclassified as a base under the Arrhenius definition. In 1923, the Danishchemist Johannes Brønsted (1879–1947) and the English chemist ThomasLowry (1874–1936) independently proposed a new definition. TheBrønsted-Lowry theory defines an acid as a hydrogen-ion donor, and a base asa hydrogen-ion acceptor. All the acids and bases included in the Arrheniustheory are also acids and bases according to the Brønsted-Lowry theory.Some compounds not included in the Arrhenius theory are classified asbases in the Brønsted-Lowry theory.

Why Ammonia is a Base The behavior of ammonia as a base can beunderstood by using the Brønsted-Lowry theory. Ammonia gas is very solu-ble in water. When ammonia dissolves in water, it acts as a base because itaccepts a hydrogen ion from water.

In this reaction, ammonia is the hydrogen-ion acceptor and thereforeis a Brønsted-Lowry base. Water, the hydrogen-ion donor, is a Brønsted-Lowry acid. Hydrogen ions are transferred from water to ammonia, as isshown in Figure 19.5. This causes the hydroxide-ion concentration to begreater than it is in pure water. As a result, solutions of ammonia are basic.

Checkpoint Why is ammonia considered to be a Brønsted-Lowry base?

Ammonia Water Ammonium Hydroxide ion(hydrogen-ion (hydrogen-ion ion (makes the

acceptor, Brønsted- donor, Brønsted- solution basic)Lowry base) Lowry acid)

NH4+ 1aq 2 + OH- 1aq 2mH2O 1 l 2�NH3 1aq 2

Figure 19.4 Sodium carbonate decahydrate (Na2CO3·10H2O) is used as a laundry aid. Its common name is washing soda.

Figure 19.3 Milk of magnesia is a base used as an antacid. Bases are usually hazardous when taken internally, but the low solubility of milk of magnesia makes it safe to use.


Use VisualsTable 19.3 Have students study the Brønsted-Lowry acid-base conjugate pairs. Ask, How are hydrogen ions related to Brønsted-Lowry acids? (Acids lose hydrogen ions.) How are hydrogen ions related to Brønsted-Lowry bases? (Bases gain hydrogen ions.)

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Answers to...Figure 19.5 It is not a hydroxide.

Checkpoint

Ammonia is considered to be a Brønsted-Lowry base because it is a hydrogen-ion acceptor.


Conjugate Acids and Bases Because all gases become less soluble inwater as temperature increases, increasing the temperature of an aqueoussolution of ammonia releases ammonia gas. As ammonia gas leaves thesolution, the equilibrium in the equation shifts to the left. The ammoniumion (NH4

�) reacts with OH� to form NH3 and H2O. When the reaction goesfrom right to left, NH4

� gives up a hydrogen ion; it acts as a Brønsted-Lowryacid. The hydroxide ion accepts an H�; it acts as a Brønsted-Lowry base.Overall, then, this equilibrium has two acids and two bases.

When ammonia dissolves and then reacts with water, NH4� is the con-

jugate acid of the base NH3. A conjugate acid is the particle formed when abase gains a hydrogen ion. Similarly, OH� is the conjugate base of the acidwater. A conjugate base is the particle that remains when an acid hasdonated a hydrogen ion. Conjugate acids and bases are always paired witha base or an acid, respectively. A conjugate acid-base pair consists of twosubstances related by the loss or gain of a single hydrogen ion. The ammo-nia molecule and ammonium ion are a conjugate acid-base pair. The watermolecule and hydroxide ion are also a conjugate acid-base pair.

The Brønsted-Lowry theory also applies to acids. Consider the dissoci-ation of hydrogen chloride in water.

In this reaction, hydrogen chloride is the hydrogen-ion donor. Thus it is aBrønsted-Lowry acid. Water is the hydrogen-ion acceptor and thereforewater is a Brønsted-Lowry base. A water molecule that gains a hydrogen ionbecomes a positively charged hydronium ion (H3O�). The chloride ion is theconjugate base of the acid HCl. The hydronium ion is the conjugate acid ofthe base water.

Base Acid Conjugate Conjugateacid base

NH3 1aq 2 + H2O 1l 2m NH4+ 1aq 2 + OH- 1aq 2

Base Acid Conjugate Conjugateacid base

NH3 1aq 2 + H2O 1 l 2m NH4+ 1aq 2 + OH- 1aq 2

Acid Base Conjugate Conjugateacid base

HCl 1g 2 + H2O 1l 2m H3O+ 1aq 2 + Cl- 1aq 2

NH3

AmmoniaH2O

WaterNH4

�

Ammonium ionOH�

Hydroxide ion

� —→

� �

�

Figure 19.5 Ammonia dissolves in water to form ammonium ions and hydroxide ions. In this reaction, the water molecule donates a hydrogen ion to the ammonia molecule. Explaining Why is ammonia not classified as an Arrhenius base?

Table 19.3

Several Conjugate Acid-Base Pairs

Acid Base

HCl Cl�

H2SO4 HSO4�

H3O� H2O

HSO4� SO4

2�

CH3COOH CH3COO�

H2CO3 HCO3�

HCO3� CO3

2�

NH4� NH3

H2O OH�

592 Chapter 19


Word OriginsAn amphibious airplane is designed to take off from and land on either land or water. The prefix, amph-, is also used in amphibian which is an animal that has aquatic larvae, with gills that grow into air-breathing adults with lungs.

Lewis Acids and BasesDiscussPoint out that Lewis was the first scien-tist to discuss the significance of elec-tron pairs in bonding (Lewis electron-dot diagrams). His theory of acids and bases was an extension of his concept of electron pairs. A Lewis acid accepts a pair of electrons to form a covalent bond and a Lewis base donates a pair of electrons to form a covalent bond.

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Answers to...Figure 19.6 hydronium ion; hydrogen sulfate ion

Checkpoint A Lewis acid is an electron-pair acceptor; a Lewis base is an electron-pair donor.

592 Chapter 19

Table 19.4

Acid-Base Definitions

Figure 19.6 shows the reaction that occurs when sulfuric acid dissolvesin water. The solution formed consists of hydronium ions and hydrogensulfate ions.

Sometimes water accepts a hydrogen ion. At other times, it donates ahydrogen ion. A substance that can act as both an acid and a base is said tobe amphoteric. Water is amphoteric. In the reaction with HCl, wateraccepts a proton and is therefore a base.

Lewis Acids and BasesA third theory of acids and bases was proposed by Gilbert Lewis (1875–1946).

Lewis proposed that an acid accepts a pair of electrons during a reac-tion, while a base donates a pair of electrons. This concept is more generalthan either the Arrhenius theory or the Brønsted-Lowry theory. A Lewis acid is asubstance that can accept a pair of electrons to form a covalent bond. A Lewisbase is a substance that can donate a pair of electrons to form a covalent bond.A hydrogen ion (Brønsted-Lowry acid) can accept a pair of electrons in forminga bond. A hydrogen ion, therefore, is also a Lewis acid. A Brønsted-Lowry base,or a substance that accepts a hydrogen ion, must have a pair of electrons avail-able and, therefore, is also a Lewis base. Consider the reaction of H� and OH�.

In this reaction, a hydroxide ion is a Lewis base. It is also a Brønsted-Lowrybase. The hydrogen ion is both a Lewis acid and a Brønsted-Lowryacid. The Lewis definition also includes some compounds not classified asBrønsted-Lowry acids or bases.

Ammonia dissolved in water is another example of a Lewis acid andLewis base. In this reaction, the hydrogen ion from the dissociation ofwater is an electron-pair acceptor and is a Lewis acid. Ammonia is an elec-tron-pair donor and is a Lewis base.

Lewisacid

Lewisbase

Figure 19.6 When sulfuric acid dissolves in water, it forms hydronium ions and hydrogen sulfate ions. Identify Whichion is the conjugate acid and which is the conjugate base?

H2SO4

Sulfuric acidH2O

WaterH3O

�

Hydronium ionHSO4

�

Hydrogen sulfate ion

� —→�

�

�

Word OriginsAmphoteric comes from the Greek word ampho-teros, meaning “partly one and partly the other.” An amphoteric substance can act as both an acid and a base. What makes an amphibious airplane dif-ferent from other air-planes?

withChemASAP

Animation 25 Compare the three important definitions of acids and bases.

Checkpoint What is a Lewis acid and a Lewis base?

Type Acid Base

Arrhenius H� producer OH� producer

Brønsted-Lowry H� donor H� acceptor

Lewis electron-pair acceptor electron-pair donor


CONCEPTUAL PROBLEM 19.1

Answers1. a. H+ is the Lewis acid; H2O is the

Lewis base.b. AlCl3 is the Lewis acid; Cl- is the Lewis base.

2. A Lewis base; it has a nonbonding pair of electrons that it can donate.

Practice Problems PlusIdentify each reactant as an acid or base.

a.

(KOH is the base; HBr is the acid)b. HCl + H2O → H3O+ + Cl–

(HCl is the acid; H2O is the base)

ASSESSEvaluate UnderstandingHave students define Arrhenius acids and bases and give an example of each. (Arrhenius acids are hydrogen-containing compounds that ionize to yield hydrogen ions in aqueous solution. HCl is an example. Arrhenius bases are compounds that ionize to yield hydrox-ide ions in aqueous solution. NaOH is an example.)

Drain cleaners contain caustic NaOH. Reaction with aluminum generates heat, which softens greases and oils, and also generates hydrogen, which agitates the mixture.

with ChemASAP

If your class subscribes to the Interactive Textbook, use it to review key concepts in Section 19.1.

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KOH + HBr KBr + H2O

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Section 19.1 Assessment3. Both are electrolytes and change the color

of an acid-base indicator. Acids have a sour taste; bases taste bitter.

4. An acid gives hydrogen ions. A base gives hydroxide ions.

5. Acids are hydrogen-ion donors and bases are hydrogen-ion acceptors.

6. A Lewis acid is an electron-pair acceptor; a Lewis base is an electron-pair donor.

7. a. two substances related by the loss or gain of a single hydrogen ionb. HNO3 + H2O → H3O+ + NO3

–; HNO3 is

the hydrogen-ion donor; its conjugate base is NO3

–. H2O is the hydrogen-ion acceptor; its conjugate acid is H3O+.

CO3

2– is the hydrogen-ion acceptor; its con-jugate acid is HCO3

–. H2O is the hydrogen-ion donor; its conjugate base is OH–.

8. a. diprotic; two ionizable hydrogensb. triprotic; three ionizable hydrogensc. monoprotic; one ionizable hydrogend. diprotic; two ionizable hydrogens

CO32– + H2O HCO3

– + OH–


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Identifying Lewis Acids and BasesAmmonia is widely used in fertilizers, plastics, and explosives. Identify the Lewis acid and the Lewis base in this reaction involving ammonia.

H

H

H

N F

F

F

B

H

H

H N B

F

F

F�..

Analyze Identify the relevant concepts.

The Lewis acid–Lewis base definitions, which are to be used in solving the problem, are based on the acceptance and donation of a pair of electrons.

Solve Apply concepts to this situation.

Ammonia is donating a pair of electrons. Boron trifluoride is accepting a pair of elec-trons. Lewis bases donate electrons, so ammonia is acting as a Lewis base. Lewis acids accept a pair of electrons, so boron trifluoride is acting as a Lewis acid.

1. Identify the Lewis acid and Lewis base in each reaction.

a.

b. AlCl3 � Cl�¡ AlCl4�

H

H H

H O.O. ..

3�

2. Would you predict PCl3 to be a Lewis acid or a Lewis base in typical reactions? Explain your prediction.

19.1 Section Assessment

3. Key Concept What are the properties of acids and bases?

4. Key Concept How did Arrhenius define an acid and a base?

5. Key Concept How are acids and bases defined by the Brønsted-Lowry theory?

6. Key Concept What is the Lewis theory of acids and bases?

7. a. What is a conjugate acid-base pair? b. Write equations for the ionization of HNO3 in

water and the reaction of CO32� with water. For

each equation, identify the hydrogen-ion donor and hydrogen-ion acceptor. Then label the conju-gate acid-base pairs in each equation.

8. Identify the following acids as monoprotic, di-protic, or triprotic. Explain your reasoning.

a. H2CO3 b. H3PO4 c. HCl d. H2SO4

Write a Report Household drain cleaners contain pellets of sodium hydroxide (NaOH) and small metal particles. Use the library or Internet to find out how drain cleaners work. Include the identity of the metal particles in your written report.

Assessment 19.1 Test yourself on the concepts in Section 19.1.

Practice Problems

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Problem-Solving 19.1 Solve Problem 1 with the help of an interactive guided tutorial.

594 Chapter 19



Review• Transparencies, T215–T222• Laboratory Manual, Lab 40• Small-Scale Chemistry Laboratory Manual,

Lab 30


Problem-Solving 19.10, 19.12, 19.14, 19.15, Assessment 19.2

• GoOnline, Section 19.2

19.2

FOCUSObjectives19.2.1 Describe how [H+] and [OH] are

related in an aqueous solution.19.2.2 Classify a solution as neutral,

acidic, or basic given the hydrogen-ion or hydroxide-ion concentration.

19.2.3 Convert hydrogen-ion concen-trations into pH values and hydroxide-ion concentrations into pOH values.

19.2.4 Describe the purpose of an acid-base indicator.

Guide for Reading

Build VocabularyGraphic Organizers Students can compare and contrast the properties of a neutral solution, an acid solution, and a basic solution using a table.

Reading StrategyIdentify Main Idea/Details Have stu-dents write down the titles of the red heads in the section. Have them follow each title with a sentence that expresses the section’s main idea.

INSTRUCT

Have students study the photograph and read the text that opens the sec-tion. Ask, What scale is commonly used to express the acidity of a solu-tion? (pH scale)

Hydrogen Ions from WaterUse VisualsFigure 19.7 Ask students to examine the figure. Illustrate the reaction on the board: Show how two water molecules can react to yield a hydronium ion and a hydroxide ion. Ask, Which element is donating a pair of electons? (oxygen) Which element is accepting a pair of electrons? (hydrogen)

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19.2 Hydrogen Ions and Acidity

A patient is brought to a hospital unconscious and with a fruity odor on his breath. The doctor suspects the patient has fallen into a diabetic coma. To confirm her diagnosis, she orders several tests, including one of the acidity of the patient’s blood. The results from this test will be expressed in units of pH, not molar concentration. In this section, you will learn how the pH scale is used to indicate the acidity of a solution and why the pH scale is used.

Guide for Reading

Key Concepts• How are [H�] and [OH�] related

in an aqueous solution?• How is the hydrogen-ion

concentration used to classify a solution as neutral, acidic, or basic?

• What is the most important characteristic of an acid-base indicator?

Vocabularyself-ionization

neutral solution

ion-product constant for water (Kw)

acidic solution

basic solution

alkaline solutions

pH

Reading StrategyRelating Text and VisualsAs you read about pH, look carefully at the diagram in Figure 19.10. Make sure that you can explain why the differences in [H�]and [OH�] exist in acidic, neutral, and basic solutions.

Hydrogen Ions from WaterAs you already know, water molecules are highly polar and are in con-tinuous motion, even at room temperature. Occasionally, the collisionsbetween water molecules are energetic enough to transfer a hydrogen ionfrom one water molecule to another. A water molecule that loses a hydro-gen ion becomes a negatively charged hydroxide ion (OH�). A water mole-cule that gains a hydrogen ion becomes a positively charged hydroniumion (H3O�).

The reaction in which water molecules produce ions is called the self-ionization of water. This reaction can be written as a simple dissociation.

In water or aqueous solution, hydrogen ions (H�) are always joined towater molecules as hydronium ions (H3O�). Hydrogen ions in aqueoussolution have several names. Some chemists call them protons. Othersprefer to call them hydrogen ions or hydronium ions. In this textbook,either H� or H3O� is used to represent hydrogen ions in aqueous solution.Figure 19.7 shows how two water molecules react to form one hydroniumion and one hydroxide ion.

Hydrogen ion Hydroxide ion

H2O1l 2m H+1aq 2 + OH-1aq 2

Figure 19.7 The self-ionization of water. A proton (hydrogen ion) transfers from one water molecule to another water molecule. The result is one hydronium ion (H3O�) and one hydroxide ion (OH�).

HH

�

�

�

�

O

Hydroniumion

Hydroxideion

Watermolecule

Watermolecule

H OH HHOH

HO

H2O H2O H3O OH

�

�

�

�

� �

�

�

—→

—→

—→


Ion Product Constant for Water

CLASS ActivityCLASS

Using a pH MeterPurpose Students observe how a pH meter measures the PH of solutions.

Materials several test tubes; 0.1M HCl; 0.1M NaOH; household products such as lemon juice, vinegar, shampoo; and liquid detergent; pH meter

Procedure Set up test tubes with acids and bases of varying strength. In addition to HCl and NaOH, include household products such as lemon juice, vinegar, shampoo, and liquid detergent. Use a pH meter to measure the pH of each solution. Explain that the reading on the pH meter depends upon the concentration of hydrogen ions in the solutions.

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Answers to...Figure 19.8 Hydrochloric acid will increase [H+] and decrease [OH–]; sodium hydroxide will increase [OH–] and decrease [H+].

Checkpoint A basic solu-tion is one in which [H+] is less than [OH–]

Differentiated InstructionEnglish LearnersStudents who do not have a good back-ground in scientific notation may have trou-ble understanding the relationship between hydrogen- and hydroxide-ion concentrations. Draw diagrams to help students visualize the shift in equilibrium that occurs when one or the other kind of ion is added or removed.

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Section 19.2 Hydrogen Ions and Acidity 595

Figure 19.8 Acids and bases have many uses in the home and in industry. Unrefined hydrochloric acid, commonly called muriatic acid, is used to clean stone buildings and swimming pools. Sodium hydroxide, or lye, is commonly used as a drain cleaner.Predicting How will each chemical affect the hydrogen-ion and hydroxide-ion concen-tration of an aqueous solution?

a

b

a

b

The self-ionization of water occurs to a very small extent. In pure waterat 25°C, the equilibrium concentration of hydrogen ions ([H�]) and theequilibrium concentration of hydroxide ions ([OH�]) are each only1 � 10�7M. This means that the concentrations of H� and OH� are equal inpure water. Any aqueous solution in which [H�] and [OH�] are equal isdescribed as a neutral solution.

Ion Product Constant for WaterIn any aqueous solution, when [H�] increases, [OH�] decreases. When [H�]decreases, [OH�] increases. Le Châtelier’s principle, which you learnedabout in Chapter 18, applies here. If additional ions (either hydrogen ionsor hydroxide ions) are added to a solution, the equilibrium shifts. The con-centration of the other type of ion decreases. More water molecules areformed in the process.

H�(aq) � OH�(aq)∆ H2O(l)

For aqueous solutions, the product of the hydrogen-ion concentra-tion and the hydroxide-ion concentration equals 1.0 � 10�14.

[H�] � [OH�] � 1.0 � 10�14

This equation is true for all dilute aqueous solutions at 25°C. As you will see,the concentrations of H� and OH� may change when substances are addedto water. However, the product of [H�] and [OH�] is always 1 � 10�14.

The product of the concentrations of the hydrogen ions and hydroxideions in water is called the ion-product constant for water (Kw).

Kw � [H�] � [OH�] � 1.0 � 10�14

Not all solutions are neutral. When some substances dissolve in water,they release hydrogen ions. For example, when hydrogen chloride dis-solves in water, it forms hydrochloric acid.

In such a solution, the hydrogen-ion concentration is greater than thehydroxide-ion concentration. The hydroxide ions are present from the self-ionization of water. An acidic solution is one in which [H�] is greater than[OH�]. The [H�] of an acidic solution is greater than 1 � 10�7M.

When sodium hydroxide dissolves in water, it forms hydroxide ions insolution.

In such a solution, the hydrogen-ion concentration is less than thehydroxide-ion concentration. Remember, the hydrogen ions are presentfrom the self-ionization of water. A basic solution is one in which [H�] is lessthan [OH�]. The [H�] of a basic solution is less than 1 � 10�7M. Basic solu-tions are also known as alkaline solutions. Some uses for acids and basesare shown in Figure 19.8.

Checkpoint What is a basic solution?

HCl1g 2 ¡H2OH+1aq 2 + Cl-1aq 2

NaOH1 s 2 ¡H2ONa+1aq 2 + OH-1aq 2

596 Chapter 19


Sample Problem 19.1

Answers9. a. basic

b. basicc. acidicd. neutral

10. 1.0 × 10-11M; basic

Practice Problems PlusThe [OH-] of a carbonated soft drink is 1.0 × 10-11M. What is the [H+] of this solution? Is the solution acidic or basic? (1.0 × 10-3M; acidic)

Math HandbookFor a math refresher and practice, direct students to scientific notation, page R56.

The pH ConceptDiscussPure water self-ionizes to form hydro-gen and hydroxide ions, yet it does not conduct electric current well. Ask, Why does this phenomenon occur? (Pure water is a poor conductor because the concentrations of the ions are low.)

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Ions in WaterExplain to the students that all aqueous sys-tems contain both hydrogen and hydroxide ions due to self-ionization of water. At 25˚C, about 1 molecule of water out of 550 000 000 will dissociate. This amounts to 1 g of hydro-gen ions and 17 g of hydroxide ions in 10 000 000 L of water. No matter how small the con-centrations, both ions are always present in a water solution.

The relationship between the concentrations of H+ and OH- in aqueous solutions at con-stant temperature is similar to the relationship between the pressure and volume of a gas at constant temperature. In both cases, the rela-tionship is inverse. The product of the two quantities is a constant—as one quantity increases, the other decreases.

Facts and Figures

596 Chapter 19

The pH ConceptExpressing hydrogen-ion concentration in molarity is cumbersome. Amore widely used system for expressing [H�] is the pH scale, proposed in1909 by the Danish scientist Søren Sørensen (1868–1939). On the pH scale,which ranges from 0 to 14, neutral solutions have a pH of 7. A pH of 0 isstrongly acidic. A solution with a pH of 14 is strongly basic.

Calculating pH The pH of a solution is the negative logarithm of thehydrogen-ion concentration. The pH may be represented mathematicallyusing the following equation.

pH � �log[H�]

Practice ProblemsPractice Problems

SAMPLE PROBLEM 19.1

Finding the [OH�] of a SolutionColas are slightly acidic. If the [H�] in a solution is 1.0 � 10�5M, is thesolution acidic, basic, or neutral? What is the [OH�] of this solution?

Analyze List the knowns and the unknowns.

Knowns• [H�] � 1.0 � 10�5M• Ion-product constant for water:

Kw � [H�] � [OH�] � 1 � 10�14

Unknowns• solution � acidic, basic, or neutral?• [OH�] � ?M

Calculate Solve for the unknowns.

[H�] is 1.0 � 10�5M. Because this is greater than 1.0 � 10�7M, thesolution is acidic.

By definition Kw � [H�] � [OH�]. Therefore, .

Substituting the known numerical values, compute [OH�] as follows.

Evaluate Do the results make sense?

If [H�] is greater than 1.0 � 10�7M, the [OH�] must be less than1.0 � 10�7M. At 1 � 10�9M, [OH�] is less than 1 � 10�7M.

3OH- 4 �Kw

3H+ 4

3OH- 4 � 1.0 � 10-14

1.0 � 10-5 � 1.0 � 10-9M

9. Classify each solution as acidic, basic, or neutral.a. [H�] � 6.0 � 10�10Mb. [OH�] � 3.0 � 10�2Mc. [H�] � 2.0 � 10�7Md. [OH�] � 1.0 � 10�7M

10. If the hydroxide-ion concen-tration of an aqueous solution is 1 � 10�3M, what is the [H�]in the solution? Is the solution acidic, basic, or neutral?

Math Handbook

For help with scientific notation, go to page R56.

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Download a worksheet on pH for students to complete, and find addi-tional teacher support from NSTA SciLinks.

Use VisualsFigure 19.9 Have students study the figure and the text that follows. Explain that the pH scale shows the relation-ship between pH and the hydrogen-ion concentration. Write the expression for the ion-product constant for water on the board, and remind students about the relationship between the concentrations of hydronium ion and hydroxide ion in an aqueous solution. Point out that the product of these concentrations in aqueous solutions is always 1 × 10-14 at 25˚C.

Arbitrarily assign concentration values to each of the hydronium-ion bars in Figure 19.9 for the acidic and basic solutions (for the neutral solution, [H+] is 1 × 10-7M); then have students calcu-late the pH of those solutions. Ask, What is [OH-] for these solutions? (Have students refer to Table 19.5 to find examples of aqueous systems with these [H+] and [OH-].)

Interpreting Graphsa. [H3O+] = 1 × 10-7Mb. [H3O+] = [OH-]c. basic: [OH-] > [H3O+]

acidic: [H3O+] > [OH-]

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Gifted and TalentedHave students identify occupations in which people need to measure pH. Interested stu-dents may select a local person in an occupa-tion from the list and ask to interview that person about his or her use of pH measure-ment. Have students write up their inter-views as if they were preparing them for a newspaper article.

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Differentiated Instruction


[H3O�] and [OH�] in Acidic, Neutral, and Basic Solutions

In a neutral solution, the [H�] � 1 � 10�7M. The pH of a neutral solution is 7.

You can calculate the pH of a solution using the log function key on acalculator. (You can review finding the logarithm of a number on a calcula-tor in the Math Handbook on page R78.)

Figure 19.9 shows how the hydrogen-ion concentration of a solutionis used to classify the solution as neutral acidic, or basic. A solution inwhich [H��] is greater than 1 �� 10��7M has a pH less than 7.0 and is acidic.The pH of pure water or a neutral aqueous solution is 7.0. A solution witha pH greater than 7 is basic and has a [H��] of less than 1 �� 10��7M.

� 7.0

� -10.0 + 1-7.02 2

� -1log 1 + log 10-72

pH � -log 11 � 10-72

• Acidic solution: pH < 7.0 [H�] greater than 1 � 10�7M• Neutral solution: pH � 7.0 [H�] equals 1 � 10�7M• Basic solution: pH > 7.0 [H�] less than 1 � 10�7M

The pH values of several common aqueous solutions are listed inTable 19.5. The table also summarizes the relationship among [H�], [OH�],and pH. You may notice that pH can sometimes be read from the value of[H�]. If [H�] is written in scientific notation and has a coefficient of 1, thenthe pH of the solution equals the exponent, with the sign changed fromminus to plus. For example, a solution with [H�] � 1 � 10�2M has a pH of2.0 and a solution with [H�] � 1 � 10�13M has a pH of 13.0.

If the pH is an integer number, it is also possible to directly write thevalue of [H�]. A solution with a pH of 9.0 has a [H�] � 1 � 10�9M. A pH of 4indicates a [H�] of 1 � 10�4M.

For: Links on pHVisit: www.SciLinks.orgWeb Code: cdn–1192

Figure 19.9 The hydrogen-ion concentrate of a solution is used to classify the solution as acidic, neutral, or basic.

Acidic Solution Neutral Solution Basic Solution

[H3O�]

[H3O�]

[H3O�][OH�]

[OH�]

[OH�]

Greater than1 � 10�7M

1 � 10�7M

Less than1 � 10�7M

Co

ncen

trati

on

INTERPRETING GRAPHS

a. Identify What is [H3O�] in a neutral solution?b. Describe How does [H3O�]compare with [OH�] in an acidic solution?c. Compare and Contrast Interms of ion concentrations, how are basic solutions different from acidic solutions?

598 Chapter 19


DiscussIntroduce pH as a simpler way to express hydrogen-ion concentration. Point out that it is easier to say the pH of a solution is 3.00 than to say that the hydrogen ion concentration is equal to 1.0 × 10-3 moles per liter. For students who have no concept of logarithms, explain that pH is found by taking the negative of the power (exponent) of 10 that expresses the hydrogen ion con-centration. Show students this exam-ple: [H] = 0.00010 = 1.0 × 10-4M; pH = –log10(1.0 × 10-4) = 4.00. Most pH val-ues are positive numbers, but negative values are also possible. Assuming 100% ionization, for example, the pH of a 10M HCl solution is –1.00.

DiscussHave students compare the pH and pOH of a solution by posing a series of questions such as the following: In an acidic solution with a pH of 3.25, are there any OH- ions? (Yes, there are always some OH- ions in an aqueous solution.) How do you know this? (The product of the hydrogen ion concentra-tion and the hydroxide ion concentration must always equal Kw.) What is the pOH of this solution? (pH + pOH = 14; pOH = 14 - 3.25 = 10.75)

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598 Chapter 19

Calculating pOH The pOH of a solution equals the negative logarithmof the hydroxide-ion concentration.

pOH � �log[OH�]

A neutral solution has a pOH of 7. A solution with a pOH less than 7 is basic.A solution with a pOH greater than 7 is acidic. A simple relationship betweenpH and pOH allows you to find either one when the other is known.

pH and Significant Figures For pH calculations, you should expressthe hydrogen-ion concentration in scientific notation. For example, ahydrogen-ion concentration of 0.0010M, rewritten as 1.0 � 10�3M in scien-tific notation, has two significant figures. The pH of this solution is 3.00,with the two numbers to the right of the decimal point representing the twosignificant figures in the concentration. A solution with a pH of 3.00 isacidic, as shown in Figure 19.10.

pOH � 14 - pH

pH � 14 - pOH

pH + pOH � 14

Table 19.5

Relationship among [H�], [OH�], and pH

Figure 19.10 The pH scale shows the relationship between pH and the hydrogen-ion concentration. Interpreting Diagrams What happens to [H��] as pH increases?

0 1 2 3 4 5 6 7 8 9 10 11 12 1413

100

or 110�1 10�2 10�3 10�4 10�5 10�6 10�7 10�8 10�9 10�10 10�11 10�12 10�1410�13

Neu

tral

pH

[H�]

Increasing acidity Increasing basicity

[H�] (mol/L) [OH�] (mol/L) Aqueous system

1 � 100 1 � 10�14

1 � 10�1 1 � 10�13

1 � 10�2 1 � 10�12

1 � 10�3 1 � 10�11

1 � 10�4 1 � 10�10

1 � 10�5 1 � 10�9

1 � 10�6 1 � 10�8

1 � 10�7 1 � 10�7

1 � 10�8 1 � 10�6

1 � 10�9 1 � 10�5

1 � 10�10 1 � 10�4

1 � 10�11 1 � 10�3

1 � 10�12 1 � 10�2

1 � 10�13 1 � 10�1

1 � 10�14 1 � 100

Incr

easi

ng

aci

dit

y

Neutral

Incr

easi

ng

bas

icit

y

0.0

1.0

2.0

3.0

4.0

5.0

6.0

7.0

8.0

9.0

10.0

11.0

12.0

13.0

14.0

pH

1M HCI

0.1M HCIGastric juiceLemon juice

Tomato juiceBlack coffee

MilkPure water

Blood

Sodium bicarbonate,sea water

Milk of magnesia

Household ammoniaWashing soda

0.1M NaOH

1M NaOH


Sample Problem 19.2

Answers11. a. pH = –log[H+] = –log (1 × 10-4)

= 4.0b. pH = –log[H+] = –log (1.5 × 10-3) = 2.82

12. a. pH = –log[H+] = –log (1.0 × 10-12) = 12.00b. pH = –log[H+] = –log (4.5 × 10-2) = 1.35

Practice Problems PlusWhat are the pH values of the follow-ing three solutions, based on their hydrogen ion concentrations?a. [H+] = 1.0 × 10-11M (11.00)b. [H+] = 1.0 × 10-8M (8.00)c. [H+] = 0.00001M (5.0)

LogarithmsThe CHEMath feature will help stu-dents who are struggling with log-arithms. In Section 19.2, focus on solutions of strong acids and bases. Save calculating dissocia-tion constants for Section 19.3.

Math HandbookFor a math refresher and practice, direct students to logarithms, page R78.

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Answers to...Figure 19.10 [H+] decreases


Math Handbook


Most pH values are not whole numbers. For example, milk of magnesiahas a pH of 10.5. Using the definition of pH, this means that [H�] mustequal 1 � 10�10.5M. Thus [H�] must be less than 1 � 10�10M (pH 10.0), butgreater than 1 � 10�11M (pH 11.0). If [H�] is written in scientific notationbut its coefficient is not 1, then you use a calculator with a log function keyto calculate pH.

SAMPLE PROBLEM 19.2

Calculating pH from [H�]What is the pH of a solution with a hydrogen-ion concentration of4.2 � 10�10M?

Analyze List the knowns and the unknown.

Knowns• [H�] � 4.2 � 10�10M• pH � �log[H�]

Unknown• pH � ?

Calculate Solve for the unknown.

Using a calculator, log (4.2 � 10�10) � �9.37675pH � �(�9.38) � 9.38.


The calculated pH is between 9 ([H�] � 1 � 10�9M) and 10 ([H�] � 1 � 10�10M). The pH is rounded to two decimal places because the hydrogen-ion concentration has two significant figures.

� 10.00

� -1-10.002

� -10.00 + 1-10.002 2

� -1log 1.0 + log 10-102

� -log 11.0 � 10-102

pH � -log 3H+ 4

11. Find the pH of each solution.a. [H�] � 1 � 10�4Mb. [H�] � 0.0015M

12. What are the pH values of the following solutions, based on their hydrogen-ion concentrations?a. [H�] � 1.0 � 10�12Mb. [H�] � 0.045M

You can calculate the hydrogen-ion concentration of a solution if youknow the pH. For example, if the solution has a pH of 3.00, then[H�] � 1 � 10�3M. When the pH is not a whole number, you will need a cal-culator with a yx function key to calculate the hydrogen-ion concentration.For example, if the pH is 3.70, the hydrogen-ion concentration is greaterthan 1 � 10�4M (pH 4.0) and less than 1 � 10�3M (pH 3.0). To get an accu-rate value, use a calculator as shown in the following example.

LogarithmsThe common logarithm (log) of a number (N) is the expo-nent (x) to which the base 10 must be raised to yield the number. If N � 10x, thenlog N � x.

The use of logarithms allows a large range of values to be conveniently expressed as small, nonexponential numbers. For example, log 103 � 3 and log 106 � 6. Although there is a range of three orders of magnitude (10 � 10 � 10 or 1000) between the numbers 103

and 106, the range of the log values is only 3.

The concentration of hydrogen ions in most aque-ous solutions, although small, can vary over many orders of magnitude. Scientists use pH, a logarithmic scale which ranges between 1 and 14, to more conveniently express the hydrogen ion concentra-tion of a solution.

For help calculating logarithms, go to page R78.

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600 Chapter 19


Sample Problem 19.3

Answers13. a. –log[H+] = 5.00; log[H+] =

–5.00; [H+] = 1.0 × 10-5Mb. –log[H+] = 12.83; log[H+] = –12.83; [H+] = 1.5 × 10-13M

14. a. –log[H+] = 4.00; log[H+] = –4.00; [H+] = 1.0 × 10-4Mb. –log[H+] = 11.55; log[H+] = –11.55; [H+] = 2.8 × 10-12M

Practice Problems PlusWhat are the hydroxide-ion concen-trations for solutions with the fol-lowing pH values?

a. 4.00 (1.0 × 10-10M)b. 8.00 (1.0 × 10-6M)c. 12.00 (1.0 × 10-2M)

A soft drink has a pH of 3.80. What is the hydrogen-ion concentration in the drink? (1.6 × 10-4M)

Measuring pH


pH IndicatorsPurpose Students observe how pH indicators react to the acidity of their environment.

Materials 0.1% solutions of thymol blue, methyl red, bromthymol blue, and phenolphthalein; test tubes; aqueous buffers spanning pH 4 to pH 10 (Either create a range of buffers spanning pH 4 to pH 10 or higher, or purchase standard buffers with different pH values. )

Procedure Explain that many natural and synthetic pigments are weak acids that change color with varying pH. Trans-fer a small volume of each buffer to a test tube. Label the tubes with their respec-tive pH values. Add 5 drops of indicator to each test tube. Set up a separate rack of tubes for each indicator and ask stu-dents to infer the pH range for which each indicator is most suitable. Students should note the colors and pH values of the solutions in tubes. In addition, add 5 drops of the indicators to two or three tubes containing a solution with an “unknown” pH and have students esti-mate its pH value using the reference

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tubes. Ask, Why would an investigator prefer to use a pH meter to measure pH? (Individual indi-cators are only responsive to pH changes in narrow ranges. Many different indicators are needed to span the entire pH range. A pH meter allows an

investigator to collect quantitative values on a con-tinuous basis throughout the entire pH range. The precision and accuracy of pH meters are superior to standard indicators. pH meters can be calibrated and used reliably at different temperatures.)

600 Chapter 19

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Math Handbook


SAMPLE PROBLEM 19.3

Using pH to Find [H�]The pH of an unknown solution is 6.35. What is its hydrogen-ionconcentration?


Knowns• pH � 6.35• pH � �log[H�]

Unknown• [H�] � ?M


First, rearrange the equation for the definition of pH to solve for theunknown.

�log[H�] � pH

Next, substitute the value of pH.

�log[H�] � 6.35

Change the signs on both sides of the equation.

log[H�] � �6.35

Finally, determine the number that has a log of �6.35, the antilog of�6.35, using the 10x key on the calculator. The antilog of �6.35 is4.5 � 10�7. Therefore, [H�] � 4.5 � 10�7M.


The hydrogen ion concentration must be between 1 � 10�6M (pH � 6)and 1 � 10�7M (pH � 7). The answer is rounded to two significantfigures because the pH was measured to two decimal places.

13. Calculate [H�] for each solution.a. pH � 5.00b. pH � 12.83

14. What are the hydrogen-ion concentrations for solutions with the following pH values?a. 4.00b. 11.55


Measuring pHPeople need to be able to measure the pH of the solutions they use. Frommaintaining the correct acid-base balance in a swimming pool, to creatingsoil conditions ideal for plant growth, to making medical diagnoses, pHmeasurements have valuable applications. For preliminary pH measure-ments and for small-volume samples, an indicator such as the one shownin Figure 19.11 is often used. For precise and continuous measurements, apH meter is preferred.

For help calculating logarithms, go to page R78.

Figure 19.11 Acid-base indicators respond to pH changes over a specific range. Phenolphthalein changes from colorless to pink at pH 7–9.


Sample Problem 19.4

Answers15. a. [H+] = Kw/[OH-] = (1.0 × 10-14)/

(4.3 × 10-5) = 2.3 × 10-10M; pH = –log(2.3 × 10-10) = 9.63b. [H+] = Kw/[OH-] = (1.0 × 10-14)/(4.5 × 10-11) = 2.2 × 10-4M; pH = –log(2.2 × 10-4) = 3.65

16. a. pH = –log(5.0 × 10-5) = 4.30b. pH = –log(8.3 × 10-10) = 9.08

Practice Problems PlusCalculate the pH for each solutiona. [OH-] = 2.5 × 10-11M (3.40)b. [OH-] = 6.5 × 10-8M (6.81)c. [OH-] = 8.7 × 10-5M (9.94)

Math HandbookFor a math refresher and practice, direct students to logarithms, page R78.


Observing pH ChangePurpose Students observe a color change resulting from a change in pH.

Materials Lemon juice, tea, Alka-Seltzer tablet, glass of 10% grape juice, household ammonia

Procedure Without explanation, add lemon juice to tea, or add an Alka-Selt-zer tablet to a glass of 10% unsweet-ened grape juice, or a few drops of household ammonia to another glass of juice. Have the students try to explain the changes they observe.

Expected Outcome Color changes occur when lemon juice, an Alka-Selt-zer tablet, or ammonia are added. The pigments in tea and grape juice are weak acids that change color with varying pH.

RelatePoint out to students that the pH scale can be compared to the Richter scale, which measures the strength of earth-quakes. On each scale, a change of one unit represents a tenfold change in the value being measured. On the Richter scale, for example, a tremor measuring 4.0 is ten times stronger than one measuring 3.0. On the pH scale, the hydrogen-ion concentration of a solu-tion with a pH of 3.0 is ten times

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greater than that of a solution with a pH of 4.0. Remind students that the lower the pH value, the more acidic the solution.


Practice Problems

Practice Problems

SAMPLE PROBLEM 19.4

Calculating pH from [OH�]What is the pH of a solution if [OH�] � 4.0 � 10�11M?


Knowns• [OH�] � 4.0 � 10�11M• Kw � [OH�] � [H�] � 1 � 10�14

• pH � �log[H�]

Unknown• pH � ?


To calculate pH, first calculate [H�] by using the definition of Kw.

With the value of [H�] determined, use the definition of pH to solve for the pH.

A calculator indicates that log 2.5 � 10�4M � �3.60205, therefore


A solution in which [OH�] is less than 1 � 10�7M would be acidicbecause [H�] would be greater than 1 � 10�7M. The pH is expressed to2 decimal places because [OH�] is expressed to 2 significant figures.

� 2.5 � 10-4M

3H+ 4 �Kw

3OH- 4� 1.0 � 10-14

4.0 � 10-11 � 0.25 � 10-3M

Kw � 3OH- 4 � 3H+ 4

pH � -log 3H+ 4 � -log 12.5 � 10-42

� 3.60

pH � -1-3.602

15. Calculate the pH of each solution.a. [OH�] � 4.3 � 10�5Mb. [OH�] � 4.5 � 10�11M

16. Calculate the pH of eachsolution.a. [H�] � 5.0 � 10�5Mb. [H�] � 8.3 � 10�10M

withChemASAP


Acid-Base Indicators An indicator (HIn) is an acid or a base that undergoes dis-sociation in a known pH range. An indicator is a valuable tool for measuringpH because its acid form and base form have different colors in solution. Thefollowing generalized equation represents the dissociation of an indicator (HIn).

The acid form dominates the dissociation equilibrium at low pH (high[H�]), and the base form dominates the equilibrium at high pH (high [OH�]).

Acid form Base form

HIn 1aq 2mOH-

H+H+1aq 2 + In-1aq 2

If you know the [OH�] of a solution, you can find its pH. The ion-productfor water defines the relationship between [H�] and [OH�]. Therefore, youcan determine [H�] by dividing Kw by the known [OH�].

602 Chapter 19


Interpreting Graphsa. thymol blueb. The indicators change color over a

very limited range of pH. This allows for a more accurate reading of pH change.

c. bromphenol blue or bromcresol green

Enrichment QuestionWhich indicator(s) would best show a change in pH from 4.6 to 4.9? (Using both bromcresol green and methyl red would give the best results.)

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602 Chapter 19

For each indicator, the change from dominating acid form to dominatingbase form occurs in a narrow range of approximately two pH units. Withinthis range, the color of the solution is a mixture of the colors of the acid andthe base forms. Knowing the pH range over which this color change occurscan give you a rough estimate of the pH of a solution. At all pH values belowthis range, you would see only the color of the acid form. At all pH valuesabove this range, you would see only the color of the base form. You couldget a more precise estimate of the pH of the solution by repeating the exper-iment with indicators that have different pH ranges for their color changes.Many different indicators are needed to span the entire pH spectrum.Figure 19.12 shows the pH ranges of some commonly used indicators.

Indicators have certain characteristics that limit their usefulness. Thelisted pH values of indicators are usually given for 25°C. At other tempera-tures, an indicator may change color at a different pH. If the solution beingtested is not colorless, the color of the indicator may be distorted. Dissolvedsalts in a solution may also affect the indicator’s dissociation. Using indica-tor strips can help overcome these problems. An indicator strip is a piece ofpaper or plastic impregnated with an indicator. The paper is dipped into anunknown solution and compared with a color chart to measure the pH.Some indicator paper is impregnated with multiple indicators. The colorsthat result, which cover a wide pH range, are shown in Figure 19.13.

Thymol blue

Bromphenol blue

Bromcresol green

Methyl red

Alizarin

Bromthymol blue

Phenol red

Phenolphthalein

Alizarin yellow R

1 20 3 4 5 6 7 8 9 10 11 12 13 14

pH

Color Ranges of Acid-Base Indicators

INTERPRETING GRAPHS

a. Identify Which indicator changes color in a solution with a pH of 2?b. Compare and ContrastWhat do you notice about the range over which each indicator changes color?c. Apply Concepts Whichindicator would you choose to show that a solution has changed from pH 3 to pH 5?

Figure 19.12 Indicators change color at a different pH.

Figure 19.13 You can find acidic and basic substances in your home. Universal indicator solution has been added to solutions of known pH in the range from 1 to 12 to produce a set of reference colors.

Universal indicator has been added to samples of vinegar, soda water, and ammonia solution. Interpreting Photographs Use the refer-ence colors to assign pH values to vinegar, soda water, and ammonia solution.

a

b

a b


CLASS ActivityCLASS

Comparing pH Indicators and pH MetersPurpose Students compare pH indi-cators and pH meters.

Materials chemical supply catalogs

Procedure Allow students to compare the costs and capabilities of pH meters versus other quantitative acid-base indi-cators. Some pH meters are hand held devices, while others are larger, bench instruments, which can be connected to computer hardware and are intended for more detailed analyses. Students should compare the resolution and pre-cision of each of the instruments and pH indicators. Ask, What are some advantages and uses of hand held devices? Why would a chemist want to have a pH meter that can be cali-brated at different temperatures? (Hand-held pH meters are portable and, therefore, convenient to use in the field. Outdoor temperatures may vary widely.)

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Answers to...Figure 19.13 vinegar: 2; soda water: 4; ammonia solution: 10 to 11Figure 19.15 A pH meter is more accurate, faster, and provides con-tinuous readings.


pH Meters A pH meter makes rapid, accurate pH measurements. Mostchemistry laboratories have a pH meter. A pH meter connected to a com-puter or chart recorder can be used to make a continuous recording of pHchanges. As you can see in Figure 19.15, the pH meter gives a direct readoutof pH.

A pH meter is often easier to use than liquid indicators or indicatorstrips. Measurements of pH obtained with a pH meter are typically accu-rate to within 0.01 pH unit of the true pH. The color and cloudiness of theunknown solution do not affect the accuracy of the pH value obtained.Hospitals use pH meters to find small but meaningful changes of pH inblood and other body fluids. Sewage, industrial effluents, and soil pH arealso easily monitored with a pH meter.

Figure 19.14 Altering soil pH can affect the development of plants. In acidic soils, hydrangeas produce blue flowers. In basic soils, hydrangeas produce pink flowers. Evergreen plants

suffer from chlorosis, a yellowing of the foliage if soil pH is too basic.

a

b

c

d

Figure 19.15 A pH meter provides a quick and accurate way to measure the pH of a solution. Water is neutral, having a pH of 7. The pH of vinegar, a dilute aqueous solution of ethanoic (acetic) acid, is about 3. The pH of milk of magnesia, an aqueous suspension of magnesium hydroxide, is 10.5.Applying Concepts What are some advantages of using a pH meter rather than an indicator?

a

b

c

a b

c d

a b c

604 Chapter 19


ASSESSEvaluate UnderstandingHave students use equations to describe the relationship between the concentrations of hydrogen and hydroxide ions in pure water and to show what happens to the equilibrium when HCl is added to water. Repeat the procedure for NaOH. (At 25˚C, [H+] × [OH-] = 1.0 × 10-14. When acid is added, the [H+] increases and the [OH-] decreases. When base is added the [OH-] increases and the [H+] decreases.)

Quick LABQuick LAB

Indicators from Natural SourcesObjective Students will measure the pH of various household materials using a natural indicator and an indicator chart.

Prep Time 20 minutes

Class Time If the period allotted to lab work is short, spread the work over two days.

Safety Perform this lab in a well-ventilated room.Expected Outcome Students will use a natural indicator to determine if tested materials are acidic, basic, or neutral.

Analyze and Conclude1. The initial color is purple. In an acidic

solution, the indicator is red; in a neutral solution, it is blue-purple; and in a basic solution, it is green.

2. Changes in the relative number of H+ and OH– ions present are respon-sible for the color changes.

3. Answers will vary depending on the materials chosen.

4. Personal hygiene items mainly test neutral; cleaning materials, such as soap, mainly test basic.

Elements HandbookStudents will note that the pH affects the relative concentrations of ClO– and HClO in pool water.

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Section 19.2 Assessment17. [H+] × [OH-] = 1.0 × 10-14; when [H+] in a

solution increases, the [OH-] decreases.18. a. Hydroxide ion concentration is greater.

b. Hydrogen ion concentration is greater. c. The concentrations are equal.

19. The color of HIn (aq) must be different than the color of In- (aq)

20. a. 6.0 21. b. 4.00 22. c. 12.0 23. d. 3.0 21. a. 1.0 × 10-8M 22. b. 1.0 × 10-5M

c. 1.0 × 10-2M

604 Chapter 19

Quick LABQuick LAB

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Indicators from Natural Sources

PurposeTo measure the pH of vari-ous household materials by using a natural indicator to make an indicator chart.

Materials

• knife

• red cabbage leaves

• 1-cup measure

• hot water

• 2 jars

• clean white cloth

• teaspoon

• tape

• 3 sheets of plain white paper

• pencil

• ruler

• 10 clear plastic cups

• white vinegar (CH3COOH)

• baking soda (NaHCO3)

• household ammonia

• dropper

• various household items listed in Step 5

Procedure

1. Put cup of finely chopped red cab-bage leaves in a jar and add cup of hot water. Stir and crush the leaves with a spoon. Continue the extraction until the water is distinctly colored.

2. Strain the extract through a piece of cloth into a clean jar. This liquid is your natural indicator.

3. Tape three sheets of paper end to end. Draw a line along the center and label it at 5-cm intervals with the numbers 1 to 14. This is your pH scale.

4. Pour your indicator to about 1-cm depth into each of three plastic cups. To one cup, add several drops of vine-gar, to the second add a pinch of bak-ing soda, and to the third add several drops of ammonia. The resulting colors indicate pH values of about 3, 9, and 11, respectively. Place these colored positions on your pH scale.

5. Repeat Step 4 for household items such as table salt, borax, milk, lemon juice, laundry detergent, dish deter-gent, milk of magnesia, mouthwash,

12 1

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toothpaste, shampoo, and carbonated beverages.

Analyze and Conclude

1. What was the color of the indicator at acidic, neutral, and basic conditions?

2. What chemical changes were responsi-ble for the color changes?

3. Label the materials you tested as acidic, basic, or neutral.

4. Which group contains items used for cleaning or for personal hygiene?


17. Key Concept What is the relationship between [H�] and [OH�] in an aqueous solution?

18. Key Concept What is true about the relative concentrations of hydrogen ions and hydroxide ions in each kind of solution?a. basic b. acidic c. neutral

19. Key Concept What is true about the colors of a pH indicator?

20. Determine the pH of each solution. a. [H�] � 1 � 10�6M b. [H�] � 0.00010M c. [OH�] � 1 � 10�2M d. [OH�] � 1 � 10�11M

21. What are the hydroxide-ion concentrations for solutions with the following pH values?a. 6.00 b. 9.00 c. 12.00


Handbook

Group 7A Elements Hypochlorite salts are used to disinfect swimming pools. Use page R32 of the Elements Handbook to research how the pH of swimming pool water is regulated to maintain the necessary concentration of hypochlorous acid, HOCl.




Review• Transparencies, T223–T224


Problem-Solving 19.23, Assessment 19.3• Go Online, Section 19.3

19.3

FOCUSObjectives19.3.1 Define strong acids and weak

acids.19.3.2 Describe how an acid’s

strength is related to the value of its acid dissociation constant.

19.3.3 Calculate an acid dissociation constant (Ka) from concentra-tion and pH measurements.

19.3.4 Order acids by strength accord-ing to their acid dissociation constants (Ka).

19.3.5 Order bases by strength according to their base dissoci-ation constants (Kb).

Guide for Reading

Build VocabularyGraphic Organizer Students can make two columns labeled Acid and Base. Under each column, write the related vocabulary words and their definitions.

Reading StrategyPreview Students can preview the information in this section by skim-ming the headings, visuals, and bold-faced material.

INSTRUCT

Have students study the photograph and read the text that opens the sec-tion. Ask, What is an example of a weak acid? (lemon or grapefruit)

What is an example of a strong acid? (sulfuric acid)

Strong and Weak Acids and BasesUse VisualsTable 19.6 Have students study the table. Ask, Which is the weakest acid in the table? (hypochlorous acid) Which is the weakest base? (ammonia)

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Section Resources

Connecting to Your World

Section 19.3 Strengths of Acids and Bases 605

Strong Acid

Strong Base

Neutral Solution

Incr

easi

ng

stre

ng

th o

f aci

dIn

crea

sin

gst

ren

gth

of b

ase

s

19.3 Strengths of Acids and Bases

Lemons and grapefruits have a sour taste because they contain citric acid. When you make lemonade, or cut up a grapefruit, you probably do not wear safety goggles or chemical-

resistant clothing even though you are working with an acid. However, some acids require

such precautions. For example, sulfuric acid is a widely used industrial chemical

that can quickly cause severe burns if it comes into contact with skin. In this

section, you will learn what makes some acids weak acids and other acids

strong acids.

Guide for Reading

Key Concepts• How does the value of an acid

dissociation constant relate to the strength of an acid?

• How can you calculate an acid dissociation constant (Ka) of a weak acid?

Vocabularystrong acids

weak acids

acid dissociation constant (Ka)

strong bases

weak bases

base dissociation constant (Kb)

Reading StrategyComparing and ContrastingCompare the concentrations of H�

and OH� in solutions of strong acids and weak acids. Explain these differences. Do the same for strong bases and weak bases.

Strong and Weak Acids and BasesAcids are classified as strong or weak depending on the degree to whichthey ionize in water. In general, strong acids are completely ionized inaqueous solution. Hydrochloric acid and sulfuric acid are strong acids.

HCl(g) � H2O(l)¡ H3O�(aq) � Cl�(aq) (100% ionized)

Weak acids ionize only slightly in aqueous solution. The ionization of etha-

noic acid (acetic acid), a typical weak acid, is not complete.

Table 19.6 shows the relative strengths of some common acids and bases.

H3O+ 1aq 2 + CH3COO- 1aq 2EPJCH3COOH 1aq 2 + H2O 1 l 2

Ethanoic acid Water Hydronium ion Ethanoate ion

Table 19.6

Relative Strengths of Common Acids and Bases

s

Substance Formula Relative Strength

Hydrochloric acid HClNitric acid HNO3

Sulfuric acid H2SO4

Phosphoric acid H3PO4

Ethanoic acid CH3COOH

Carbonic acid H2CO3

Hypochlorous acid HClONeutral Solution

Ammonia NH3

Sodium silicate Na2SiO3

Calcium hydroxide Ca(OH)2

Sodium hydroxide NaOHPotassium hydroxide KOH

606 Chapter 19


Use VisualsFigure 19.16. Explain to students that the bar graphs compare the extent of ionization or dissociation of strong to weak acids. Ask, What is the primary difference between strong and weak acids? (Strong acids are nearly 100% dissociated, while weak acids are only partially dissociated.)

Interpreting Graphsa. The strong acid dissociates com-

pletely, forming equal amounts of ions.

b. The amount of acid that dissociates is equal to the amount of the two resulting ions.

c. The bar graph for the first dissocia-tion of oxalic acid should show low concentrations of H+ and HOOCCOO–. The bar graph for the second dissociation should show even lower concentrations of H+ and OOCCOO2–. Perceptive students may note that [H+] in the second bar graph actually should be much higher than [OOCCOO2–] because of the H+ ions produced in the first dissociation.

Enrichment QuestionRelate the amount of dissociation with the properties of strong and weak acids. (With increasing amount of dissociation, the acid becomes stronger.)

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606 Chapter 19

Dissociation of a Strong Acid, Weak Acid, and Very Weak Acid

Acid Dissociation Constant In an aqueous solution of ethanoic acid,fewer than 1% of ethanoic acid molecules are ionized at any instant.Therefore, ethanoic acid is considered a weak acid. Figure 19.16 comparesthe extent of dissociation of strong, weak, and very weak acids. A strongacid completely dissociates in water. As a result, [H3O�] is high. Hydrochlo-ric acid and sulfuric acid are examples of strong acids. On the other hand,weak acids, such as boric acid and carbonic acid, remain largely undissoci-ated. The [H3O�] of a weak acid is low.

You can write the equilibrium-constant expression from the balancedchemical equation. The equilibrium-constant expression for ethanoic acidis shown below.

Keq �3H3O+ 4 � 3CH3COO- 4

3CH3COOH 4 � 3H2O 4

Figure 19.16 Dissociation of an acid (HA) in water yields H3O� and A�. The bar graphs compare the extent of dissociation of strong, weak, and very weak acids.

HA(aq) � H2O(l)uy H3O�(aq) � A�(aq)

uuuuuyComplete

dissociation

Rel

ativ

e nu

mbe

r of

mol

es HA H3O� A�

Strong Acid

HA(aq) � H2O(l)E

H3O�(aq) � A�(aq)

DJ

uuuuuyModerate

dissociation

Rel

ativ

e nu

mbe

r of

mol

es HA

HA

H3O� A�

Weak Acid

HA(aq) � H2O(l)E

H3O�(aq) � A�(aq)

DJ

uuuuuyVery little

dissociation

Rel

ativ

e nu

mbe

r of

mol

es HA

Very Weak Acid

HA

H3O� A�

INTERPRETING GRAPHS

a. Explain In the graph for the strong acid, why are the heights of the H3O� and A�

bars the same as the height of HA bar?b. Inferring In the graph for the weak acid, why is the height of the H3O� bar the same as the distance from the top of the second HA bar to the dotted line?c. Apply Concepts Draw a bar graph for the dissociation of the weak diprotic acid, oxalic acid. Be sure to include the first and second dissociation.


DiscussRemind students that, in solution, the negative ion of an acid is a base. Explain that strong acids, such as nitric acid, dis-sociate almost completely because the negative ion is a very weak base. That is, its tendency to combine with a hydro-gen ion is slight. In contrast, the dissoci-ation of a weak acid, such as carbonic acid, is limited, because its negative ion is a relatively strong base. These ions combine with most of the available hydrogen ions to form the undissoci-ated form of the acid again.

The terms strong and weak, as applied to acids, are often confused with the concept of concentration. Stress that the dissociation of an acid or base into ions involves the establishment of equilibrium. The terms strong and weak refer to the position of the equi-librium. When a strong acid or base dis-solves, the equilibrium favors the products. When a weak acid or base dissolves, the equilibrium favors the reactants.

Show students that the extent to which products or reactants are favored can be determined from the values of Ka or Kb. Also, discuss how val-ues of Ka and Kb can be used to com-pare the strengths of acids and bases.

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Answers to...

Checkpoint

Stronger acids have larger Ka values and ionize more completely to produce H3O+ ions.

Losing IonsDiprotic and triprotic acids lose their hydro-gen ions one at a time. Each successive ion is more difficult to lose from the negative ion. Phosphoric acid (H3PO4) loses one hydrogen ion relatively easily, leaving an H2PO4

– ion. This negatively charged ion loses fewer

hydrogen ions than phosphoric acid, and even fewer hydrogen ions leave HPO4

2–. Although phosphoric acid contains three acidic hydrogens, it does not have a high degree of ionization and is, therefore, a rela-tively weak acid.

Facts and Figures


For dilute solutions, the concentration of water is a constant. It can becombined with Keq to give an acid dissociation constant. An acid dissocia-tion constant (Ka) is the ratio of the concentration of the dissociated (orionized) form of an acid to the concentration of the undissociated (nonion-ized) form. The dissociated form includes both the H3O� and the anion.

Keq � 3H2O 4 � Ka �3H3O+ 4 � 3CH3COO- 4

3CH3COOH 4

The acid dissociation constant (Ka) reflects the fraction of an acid in theionized form. For this reason, dissociation constants are sometimes calledionization constants. If the value of the dissociation constant is small, thenthe degree of dissociation or ionization of the acid in the solution is small.

Weak acids have small Ka values. The stronger an acid is, the larger is itsKa value. A larger value of Ka means the dissociation or ionization of the acidis more complete. For example, nitrous acid (HNO2) has a Ka of 4.4 � 10�4,whereas ethanoic acid (acetic acid) has a Ka of 1.8 � 10�5. This means thatnitrous acid is more ionized in solution than ethanoic acid. Nitrous acid isa stronger acid than ethanoic acid.

Therefore, a strong acid has a higher [H3O�] and a large dissociationconstant. Conversely, a weak acid has a low [H3O�] and a small dissociationconstant. Diprotic and triprotic acids lose their hydrogens one at a time.Each ionization reaction has a separate dissociation constant. Thus, phos-phoric acid (H3PO4) has three dissociation constants to go with its threeionizable hydrogens. Table 19.7 shows the ionization reactions and disso-ciation constants of some common weak acids, ranked by the value of thefirst dissociation constant of each acid.

Checkpoint What distinguishes a strong acid from a weak acid?

Figure 19.17 A solution of phosphoric acid is used to remove lime deposits from plumbing fixtures. Lime deposits are calcium carbonate.

Table 19.7

Dissociation Constants of Weak Acids

LIME-GONE

Acid Ionization Ka (25�C)

Oxalic acid HOOCCOOH(aq)∆∆∆∆H�(aq) � HOOCCOO�(aq) 5.6 � 10�2

HOOCCOO�(aq)∆∆∆∆H�(aq) � OOCCOO2�(aq) 5.1 � 10�5

Phosphoric acid H3PO4(aq)∆∆∆∆H�(aq) � H2PO4�(aq) 7.5 � 10�3

H2PO4�(aq)∆∆∆∆H�(aq) � HPO4

2�(aq) 6.2 � 10�8

HPO42�(aq)∆∆∆∆H�(aq) � PO4

3�(aq) 4.8 � 10�13

Methanoic acid HCOOH(aq)∆∆∆∆H�(aq) � HCOO�(aq) 1.8 � 10�4

Benzoic acid C6H5COOH(aq)∆∆∆∆H�(aq) � C6H5COO�(aq) 6.3 � 10�5

Ethanoic acid CH3COOH(aq)∆∆∆∆H�(aq) � CH3COO�(aq) 1.8 � 10�5

Carbonic acid H2CO3(aq)∆∆∆∆H�(aq) � HCO3�(aq) 4.3 � 10�7

HCO3�(aq)∆∆∆∆H�(aq) � CO3

2�(aq) 4.8 � 10�11

608 Chapter 19


CLASS ActivityCLASS

Shampoo SurveyPurpose Students compare sham-poos for normal, dry, or oily hair.

Materials different brands and types of shampoo; 10-mL graduated cylin-der; beakers; universal indicator paper

Procedure Have groups of students conduct a survey of shampoo prod-ucts. Each group should gather data on a different brand and type of shampoo. Decide as a class what information should be included in the survey and how the data will be reported. To assure uniform pH testing, have groups prepare 1% shampoo solutions and use universal indicator paper.

Expected Outcome Shampoos labeled for use on normal, dry, or oily hair are formulated by controlling the strength and amount of the synthetic detergent. The quantity of the active ingredient controls the “defatting” action, which removes oil from the hair.

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608 Chapter 19

Base Dissociation Constant Just as there are strong acids and weakacids, there are also strong bases and weak bases. Strong bases dissociatecompletely into metal ions and hydroxide ions in aqueous solution. Somestrong bases, such as calcium hydroxide and magnesium hydroxide, arenot very soluble in water. The small amounts of these bases that do dissolvedissociate completely.

Weak bases react with water to form the hydroxide ion and the conju-

gate acid of the base. Ammonia is an example of a weak base.

The equilibrium of this equation greatly favors the reverse reaction. Only

about 1% of the ammonia is present as NH4�, the conjugate acid of NH3.

The concentrations of NH4� and OH� are low and equal. The equilibrium-

constant expression for the reaction of ammonia with water is

The concentration of water is constant in dilute solutions. It can be com-bined with Keq to give a base dissociation constant (Kb).

In general, the base dissociation constant (Kb) is the ratio of the concentra-

tion of the conjugate acid times the concentration of the hydroxide ion tothe concentration of the conjugate base. The general form of this equationis as follows.

The magnitude of Kb indicates the ability of a weak base to compete with

the very strong base OH� for hydrogen ions. Because bases such as ammo-nia are weak relative to the hydroxide ion, Kb for such bases is usually small.

The Kb for ammonia is 1.8 � 10�5. The smaller the value of Kb, the weaker is

the base.

Checkpoint Which is larger, the Ka of a weak acid or the Ka ofa strong acid?

NH4+1aq2 + OH-1aq2EPJ

NH3 1aq 2 + H2O 1l 2Ammonia Water Ammonium Hydroxide

ion ion

Keq �3NH4

+ 4 � 3OH- 4

3NH3 4 � 3H2O 4

Keq � 3H2O 4 � Kb �3NH4

+ 4 � 3OH- 4

3NH3 4

Kb �3conjugate acid 4 � 3OH- 4

3conjugate base 4

Figure 19.18 Window cleaners use an ammonia solution to clean glass because it is a weak base. Being a weak base also makes ammonia relatively safe to use.


Calculating Dissociation ConstantsDiscussExplain how the dissociation constant of an acid or base can be determined experimentally from the concentra-tion of the solution and its pH. The cal-culation is based on the assumption that the acid molecule forms an equal number of hydrogen ions and negative ions when it dissociates. This is true only if there is no additional source of either hydrogen ions or the negative ions.

Point out that calculating Ka is much more complicated in the case of a polyprotic acid. The dissociation of H3PO4, for example, involves three sep-arate ionizations, each with its own Ka.

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Answers to...

Checkpoint

The Ka of a strong acid is larger.


Concentration

Acid or base Moles/liter (molarity) Grams/liter

Concentrated hydrochloric acid 12 438

Dilute hydrochloric acid 6 219

Concentrated sulfuric acid 18 1764

Dilute sulfuric acid 6 588

Concentrated phosphoric acid 15 1470

Concentrated nitric acid 16 1008

Dilute nitric acid 6 378

Ethanoic acid, glacial 17 1020

Ethanoic acid, dilute 6 360

Dilute sodium hydroxide 6 240

Concentrated aqueous ammonia 15 255

Dilute aqueous ammonia 6 102

Concentration and Strength The words concentrated and dilute indi-cate how much of an acid or base is dissolved in solution. These terms referto the number of moles of the acid or base in a given volume. The wordsstrong and weak refer to the extent of ionization or dissociation of an acidor base. They indicate how many of the particles ionize or dissociate intoions. Hydrochloric acid (HCl)(aq) is a strong acid; it is completely dissoci-ated into ions. Gastric juice in the stomach is a dilute solution of hydro-chloric acid. A relatively small number of HCl molecules are present in agiven volume of gastric juice, but they are all dissociated into ions. A sam-ple of hydrochloric acid added to a large volume of water becomes moredilute, but it is still a strong acid. Vinegar is a dilute solution of a weak acid,ethanoic acid. Pure ethanoic acid (glacial acetic acid) is still a weak acid,even though it is highly concentrated. Solutions of ammonia can be diluteor concentrated, depending on the amount of ammonia dissolved in agiven volume of water. In any solution of ammonia, however, whether con-centrated or dilute, ammonia will be a weak base because the amount ofionization will be small. Table 19.8 lists the concentrations of acids andbases commonly found in the laboratory.

Calculating Dissociation ConstantsYou can calculate the acid dissociation constant (Ka) of a weak acid or thebase dissociation constant (Kb) of a weak base from experimental data.

To find the Ka of a weak acid or the Kb of a weak base, substitute the mea-sured concentrations of all the substances present at equilibrium into theexpression for Ka or Kb. For a weak acid, you can determine these concentra-tions experimentally if you know the initial molar concentration of the acidand the pH (or [H3O�]) of the solution at equilibrium.

Table 19.8

Concentration of Some Common Laboratory Acids and Bases

610 Chapter 19


Sample Problem 19.5

Answers22. Ka = (4.2 × 10-3)(4.2 × 10-3)/(9.58 ×

10-2) = 1.8 × 10-4

23. Ka = (9.86 × 10-4)(9.86 × 10-4)/(2.0 × 10-1) = 4.86 × 10-6

Practice Problems PlusA solution of a weak acid, exactly 0.500M, has a [H+] = 5.77 × 10-6Ma. What is the pH of this solution? (5.239)b. What is the value of Ka for this acid? (6.66 × 10-11)

Math HandbookFor a math refresher and practice, direct students to scientific notation, page R56.

Stone ConservatorHave students do research about the sources and causes of acid rain. (Acid rain is thought to be due principally to the release of sulfur oxides and nitrogen oxides into the atmosphere. ) Ask, How do weather patterns affect the distri-bution of environmental damage due to acid rain? (Prevailing winds generally carry pollutants from west to east.) What types of vegetation, if any, are most resistant to acid rain, or tend to thrive in acidic soils? (Students may wish to speak to staff at nurseries for help answering this question.) What types of conservation efforts are used to adjust the pH of lakes and rivers that become too acidic? (A special slurry of CaCO3, or limestone, is sometimes added to lakes and rivers to control pH.)

Have students research chemistry-related careers in the library or on the Internet. Students can then con-struct a table that describes the nature of the work, educational and training requirements, employ-ment outlook, working conditions, and other necessary information.

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610 Chapter 19

In general, for an acid in water you can find Ka by substituting the con-centrations of the acid, [HA], the negative ion from the dissociation of theacid, [A�], and the hydrogen ion, [H�], into the equation below.

Ka �3H+ 4 3A- 4

3HA 4

Math Handbook

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Practice Problems

Practice Problems

SAMPLE PROBLEM 19.5

Calculating a Dissociation ConstantA 0.1000M solution of ethanoic acid is only partially ionized. Frommeasurements of the pH of the solution, [H�] is determined to be1.34 � 10�3M. What is the acid dissociation constant (Ka) of ethanoic acid?


Knowns• [ethanoic acid] � 0.1000M

• [H�] � 1.34 � 10�3M

• CH3COOH(aq) � H2O(l)∆ H3O�(aq) � CH3COO�(aq)

•

Unknown• Ka � ?


Each molecule of CH3COOH that ionizes gives an H� and a CH3COO�

ion. Therefore, at equilibrium [H�] � [CH3COO�] � 1.34 � 10�3M.The equilibrium concentration of CH3COOH is the initial con-centration minus the concentration of the ionized acid or (0.1000 � 0.00134)M � 0.0987M.

Substitute the equilibrium values into the expression for Ka.

Evaluate Does the result make sense?

The value of Ka is consistent with that of a weak acid.

Concentration [CH3COOH] [H�] [CH3COO�]

Initial 0.1000 0 0

Change �1.34 � 10�3 1.34 � 10�3 1.34 � 10�3

Equilibrium 0.0987 1.34 � 10�3 1.34 � 10�3

Ka �3H+ 4 � 3CH3COO- 4

3CH3COOH 4

� 1.82 � 10-5

K a �3H+ 4 � 3CH3COO- 4

3CH3COOH 4 �11.34 � 10-32 � 11.34 � 10-32

0.0987

22. In an exactly 0.1M solution of methanoic acid, [H�] �

4.2 � 10�3M. Calculate the Ka

of methanoic acid.

23. In an exactly 0.2M solution of a monoprotic weak acid, [H�] � 9.86 � 10�4M.What is the Ka for this acid?


For help with scientific notation, go to page R56.


ASSESSEvaluate UnderstandingAsk, List the two conditions needed to calculate the dissociation con-stant of a weak acid. (You must know the initial molar concentration of the acid and the pH or [H+] of the solution at equilibrium.)

ReteachTable 19.7, may pose difficulties for some students. Use an example to explain the information in the table. Point out that the Ka for each acid is equal to the product of the concentra-tions of the ions on the right of the equation, divided by the concentration of the non-ionized acid. Also explain that the Ka value is a quantitative indi-cation of acid strength. A lower Ka value means the reactants are favored in the equilibrium; thus, the acid is weaker.

As a class activity, have students pre-pare a numbered list of steps for calcu-lating dissociation constants. Give them a different example than those in Sample Problem 19.5, and have them use the numbered procedure to work through the problem.

To thwart tooth decay, reduce sugar intake; brush and floss regularly; use a fluoride toothpaste or rinse; and have regular dental check-ups to catch minor decay early.

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Section 19.3 Assessment24. A strong acid is completely ionized in

aqueous solution and has a large Ka. A weak acid is ionized only slightly in aque-ous solution and has a small Ka.

25. Substitute the measured concentrations of all the substances present at equilib-rium into the expressions for Ka or Kb.

26. hypochlorous acid27. The [HX] is much greater than the [H+].28. a. HNO3 + H2O → H3O+ + NO3

–

b.

c.

d. Mg(OH)2 → Mg2+ + 2OH–

29. Strong acids and bases are completely ionized in aqueous solution. Weak acids and bases ionize only slightly in aqueous solution. Concentrated acid or base solu-tions contain large amounts (high con-centrations) of acid or base. Dilute acid or base solutions contain small amounts (low concentrations) of acid or base.

CH3COOH + H2O CH3COO– + H3O+

NH3 + H2O NH4+ + OH–


Cause and Effect Paragraph The chief cause of tooth decay is the weak acid called lactic acid (C3H6O3). Lactic acid is formed in the mouth by the action of specific bacteria, such as Streptococcus mutans, on sugars present in sticky plaque on tooth surfaces. Research current efforts to thwart tooth decay. Write a report summarizing your findings.

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Stone ConservatorAs more is learned about

the damaging effects of air-borne pollutants, such as acid rain, there is a growing con-cern to preserve historic buildings and pieces of sculp-ture from damage. Stone conservators work to prevent and repair the damage to stone used in buildings and sculptures.

Stone conservators must clean statues properly to remove pollution deposits. One method they use is to apply a thin, clay mudpack to the stone’s surface to pull out the deposits. Lasers are also used to remove pollution from stone. When making repairs, stone conservators sometimes use surgical microscopes to examine the surfaces of statues. If the stone

was originally painted, the origi-nal paint is chemically analyzed to determine its composition. In addition, conservators research and apply methods to preserve the stone once it has been repaired. One preservation tech-nique is to seal the stone to pre-vent water, which carries dissolved gases and salts, from seeping into the pores of the stone.

Stone conservators often work for museums, universi-ties, or private companies that specialize in the cleaning and restoration of stone objects. They must have knowledge of both chemistry and art because they often work closely with both chem-ists and museum personnel.

A stone conservator’s level of edu-cation may range from a bache-lor’s degree to a doctorate.


24. Key Concept Compare a strong acid and a weak acid in terms of the acid dissociationconstant.

25. Key Concept How do you determine the Ka

of a weak acid or the Kb of a weak base?

26. Which acid in Table 19.6 would you expect to have the lowest ionization constant?

27. Acid HX has a very small value of Ka. How do the relative amounts of H� and HX compare at equi-librium?

28. Write the equations for the ionization or dissoci-ation of the following acids and bases in water.

a. nitric acid b. ethanoic acid c. ammonia d. magnesium hydroxide

29. Compare the terms strong/weak and concen-trated/dilute as they pertain to acids and bases.


For: Careers in ChemistryVisit: PHSchool.comWeb Code: cdb-1193

612 Chapter 19



Review• Transparencies, T225–T226• Laboratory Manual, Labs 41, 42, 43• Small-Scale Chemistry Laboratory

Manual, Lab 31

• Probeware Lab Manual, The Neutralizing Power of Antacids, Small-Scale Titrations


Problem-Solving 19.30, 19.33, Simulation 26, Assessment 19.4

19.4

FOCUSObjectives19.4.1 Define the products of an acid-

base reaction.19.4.2 Explain how acid-base titra-

tion is used to calculate the concentration of an acid or a base.

19.4.3 Explain the concept of equiva-lence in neutralization reactions.

19.4.4 Describe the relationship between equivalence point and the end point of a titration.

Guide for Reading

Build VocabularyGraphic Organizers Have students write definitions of standard solution, titration, and end point in the first col-umn of a two column table. Have them describe how these three terms are related in the second column.

Reading StrategyRelate Text and Visuals Have students reference the visuals as they learn about titration.

INSTRUCT

Have students study the photograph and read the text that opens the section. Ask, What is a neutralization reaction? (the reaction between an acid and a base, producing a salt and water)

Why not use a stronger base to pro-vide quicker relief from excess stom-ach acid? (The stronger base would be too corrosive for the digestive tract.)

Acid-Base ReactionsDiscuss Ask students if they have witnessed or read a news report about people try-ing to clean up a chemical spill on a highway or at a factory. Ask, If a spill involves an acid, what type of treat-ment is possible? (A weak base in pow-der form, such as sodium carbonate, can be spread over the spill to absorb, react with, and neutralize the acid.)

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612 Chapter 19

19.4 Neutralization Reactions

Nearly all of the adult popula-tion suffers from acid indigestion at some time. Although hydrochloric acid is always present in the stomach, an excess can cause heartburn and a feeling of nausea. A common way to relieve the pain of acid indigestion is to take antacids to neutral-ize the stomach acid. The active ingredient in many antacids is sodium hydrogen car-bonate, aluminum hydroxide, or magnesium hydroxide. In this section, you will learn what a neutralization reaction is.

Guide for Reading

Key Concepts• What are the products of the

reaction of an acid with a base?• What is the endpoint of a

titration?

Vocabularyneutralization reactions

equivalence point

standard solution

titration

end point

Reading StrategyIdentifying a SequenceA sequence is the order in which a series of events occurs. As you read about acid-base titrations, list the steps that should be used in carrying out a precise titration. Include the reactants and how they are measured, the indicator, and what to look for as the titration nears its end point.

Figure 19.19 For fish and other aquatic animals to survive, the water in which they live must be maintained at the proper pH.

Table salt (sodium chloride) is one example of a salt,but there are many more. Salts are compounds

consisting of an anion from an acid and a cation froma base.

Acid-Base ReactionsIf you mix a solution of a strong acid containing hydronium (hydrogen)ions with a solution of a strong base that has an equal number of hydroxideions, a neutral solution results. The final solution has properties that arecharacteristic of neither an acidic nor a basic solution. Consider theseexamples:

HCl(aq) � NaOH(aq)¡ NaCl(aq) � H2O(l)

H2SO4(aq) � 2KOH(aq)¡ K2SO4(aq) � 2H2O(l)

In each example, a strong acid reacts with a strong base. If solutions ofthese substances are mixed in the mole ratios specified by the balancedequation, neutral solutions will result. Similar reactions of weak acids and/or weak bases do not usually produce neutral solutions. In general, how-ever, reactions in which an acid and a base react in an aqueous solution toproduce a salt and water are called neutralization reactions. The formationof water in a neutralization reaction is shown in Figure 19.20.

Neutralization reactions are one way to prepare pure samples of salts.You could prepare potassium chloride, for example, by mixing equal molarquantities of hydrochloric acid and potassium hydroxide. An aqueoussolution of potassium chloride would result. You could heat the solution to

evaporate the water, leaving the salt potassium chloride. Table 19.9 listssome common salts and their applications.

In general, the reaction of an acid with a base pro-duces water and one of a class of compounds called salts.

When you hear the word salt, you may think of the sub-stance that flavors your French fries or scrambled eggs.


TitrationDiscuss Point out that neutralization is a pro-cess that occurs whenever an acid reacts with a base in the mole ratios specified by the balanced equation. Explain that not all neutralization reac-tions produce neutral solutions. The relative strengths of the reactants determines whether the solution will be acidic, basic, or neutral.


Titration Using IndicatorsPurpose Students observe properties of titration.

Materials equimolar solutions of HCl and NaOH, phenolphthalein, electric stirrer with a magnetic stir bar, beaker, table salt

Safety Wear protective goggles and laboratory apron.

Procedure Write the following equa-tion on the board: HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)

Prepare equimolar solutions of HCl and NaOH. Add a few drops of phenolphtha-lein indicator to the NaOH solution. Slowly add the HCl to the NaOH while carefully stirring. If possible, use an elec-tric stirrer with a magnetic stir bar. Ask, What happens to the color of the solu-tion as the acid is added? (The deep pink color slowly fades until the solution becomes colorless.) Remind students that phenolphthalein is an acid-base indica-tor that is colorless at pH lower than 8. Have a few volunteers come up and gently touch the outside of the reac-tion beaker. Ask them to describe their observations to the class.

Ask, Is the neutralization reaction endothermic or exothermic? (exo-thermic)

Display a quantity of table salt equiva-lent to the amount of salt that could be retrieved if the neutral solution were evaporated.

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Answers to...

Checkpoint

the point at which mol H+ = mol OH-

Section 19.4 Neutralization Reactions 613

Name Formula Applications

Ammonium sulfate (NH4)2SO4 Fertilizer

Barium sulfate BaSO4 Gastrointestinal studies; white pigment

Calcium chloride CaCl2 De-icing roadways and sidewalks

Potassium chloride KCl Sodium-free salt substitute

Silver bromide AgBr Photographic emulsions

Sodium hydrogen carbonate (baking soda) NaHCO3 Antacid

Sodium carbonate decahydrate (washing soda) Na2CO3 · 10H2O Glass manufacture; water softener

Sodium chloride (table salt) NaCl Body electrolyte; chlorine manufacture

The properties of acids, bases, and salts help explain many diversephenomena. The usefulness of antacids, for example, depends on the pro-cess of acid-base neutralization. Farmers use a similar process to controlthe pH of soil.

TitrationAcids and bases sometimes, but not always, react in a 1:1 mole ratio.

When sulfuric acid reacts with sodium hydroxide, however, the ratio is1:2. Two moles of the base sodium hydroxide are required to neutralize onemole of H2SO4.

Similarly, hydrochloric acid and calcium hydroxide react in a 2:1 ratio.

You will notice in the preceding examples that the number of moles ofhydrogen ions provided by the acid are equivalent to the number ofhydroxide ions provided by the base. When an acid and base are mixed, theequivalence point is when the number of moles of hydrogen ions equals thenumber of moles of hydroxide ions.

Checkpoint What is the equivalence point of a reaction between an acid and a base?

1 mol 1 mol 1 mol 1 mol

HCl 1aq 2 + NaOH 1aq 2¡ NaCl 1aq 2 + H2O 1l 2


H2SO4 1aq 2 + 2NaOH 1aq 2¡ Na2SO4 1aq 2 + 2H2O 1 l 2


2HCl 1aq 2 + Ca(OH22 1aq 2¡ CaCl2 1aq 2 + 2H2O 1l 2

Table 19.9

Some Salts and Their Applications

Figure 19.20 In a neutralization reaction, hydronium ions (H3O�)combine with hydroxide ions (OH�) to form neutral water.

�

�

—→

—→

�

�H2OWater

H2OWater

H3O�

Hydronium ionOH�

Hydroxide ion

��

614 Chapter 19


Sample Problem 19.6

Answers30. H3PO4 + 3KOH → K3PO4 + 3H2O;

1.56 mol H3PO4 × 3 mol KOH/1 molH3PO4 = 4.68 mol KOH

31. HNO3 + NaOH → NaNO3 + H2O;0.20 mol HNO3 × 1 mol NaOH/1 molHNO3 = 0.20 mol NaOH

Practice Problems PlusHow many moles of aluminum hydroxide are needed to neutralize 2.30 mol of sulfuric acid? (1.53 moles)

Math HandbookFor a math refresher and practice, direct students to dimensional analysis, page R66.

Use VisualsTable 19.5 and Figure 19.21 Have students study Table 19.5 and Figure 19.21. Red cabbage, when used as an acid-base indicator, turns red between pH 2 and 3, pink between pH 3 and 4, and green between pH 8-12. Ask, What color will emerge when testing household ammonia? (green) Lemon juice? (red) Seawater? (green)

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Acids, Bases, and Curly HairCurly hair from a permanent wave involves an acid-base neutralization reaction. The waving solution is basic and has a pH of about 9.20. The active ingredient in the wav-ing solution is often an ammonium com-pound that is applied once the hair has been rolled. After a specified time, the hair is

rinsed to remove the excess base. Then a neutralizer is applied. The neutralizer is an acid solution that has a pH range of 2 to 6. If the hair is exposed to the waving solution for too long a period of time, the hair’s texture and appearance are damaged.

Facts and Figures

614 Chapter 19

Math Handbook

Practice Problems

Practice Problems

SAMPLE PROBLEM 19.6

Finding the Number of Moles of an Acid in NeutralizationHow many moles of sulfuric acid are required to neutralize 0.50 mol ofsodium hydroxide?


Knowns• mol NaOH � 0.50 mol• H2SO4(aq) � 2NaOH(aq)¡Na2SO4(aq) � 2H2O(l)

•

Unknown• moles H2SO4 � ? mol


The mole ratio of H2SO4 to NaOH is 1:2. The necessary number of moles

of H2SO4 is calculated using this ratio.


Because the mole ratio of H2SO4 to NaOH is 1:2, the expected numberof moles of H2SO4 should be half the given number of moles of NaOH.The answer should have two significant figures.

mol H2SO4

mol NaOH � 12

0.50 mol NaOH �1 mol H2SO4

2 mol NaOH � 0.25 mol H2SO4

30. How many moles of potas-sium hydroxide are needed to completely neutralize 1.56 mol of phosphoric acid?

31. How many moles of sodium hydroxide are required to neutralize 0.20 mol of nitric acid?

For help with dimensional analysis, go to page R66.

Figure 19.21 Red cabbage juice is used as an acid-base indicator. As the solution changes from highly acidic to basic, the color changes from red to violet to green to yellow. Predicting Would the yellow solution have a high or low pH?

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You can determine the concentration of acid (or base) in a solution byperforming a neutralization reaction. You must use an appropriate acid-base indicator to show when neutralization has occurred. As you can see inFigure 19.21, the juice of the red cabbage is an acid-base indicator. In thelaboratory, phenolphthalein is often the preferred indicator for acid-baseneutralization reactions. Solutions that contain phenolphthalein turn fromcolorless to deep pink as the pH of the solution changes from acidic tobasic. In slightly basic solutions, the indicator is very faintly pink.



Titration Using a pH MeterPurpose Students observe properties of the equivalence point.

Materials equimolar solutions of HCl and NaOH, phenolphthalein, beaker, buret, pH meter, graph paper

Procedure Repeat the demo described on page 613. This time place the base in a beaker and the acid in a buret. Use a pH meter to monitor the titration. Construct a graph similar to the one in Figure 19.23

Expected Outcome Compared with Figure 19.23, the graph will be reversed because an acidic solution is being added to a basic solution.

Ask, What happens to the graph as the titration nears its equivalence point? (The slope becomes very steep.) What would happen to the pH if a small amount of base were added to the beaker after the equivalence point is reached? (The pH would increase dramatically.)

Add some base to the beaker and observe the pH. Ask, What can be done to regain the equivalence point? (Add acid from the buret.)

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Answers to...Figure 19.21 highFigure 19.23 They are equal for this titration.

Section 19.4 Neutralization Reactions 615

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The steps in a neutralization reaction are as follows.

1. A measured volume of an acid solution of unknown concentration is added to a flask.

2. Several drops of the indicator are added to the solution while the flask is gently swirled.

3. Measured volumes of a base of known concentration are mixed into the acid until the indicator just barely changes color.

The process of adding a known amount of solution of known concentra-tion to determine the concentration of another solution is called titration.The solution of known concentration is called the standard solution. Use aburet to add the standard solution. A titration is continued until the indica-tor shows that neutralization has just occurred. The point at which theindicator changes color is the end point of the titration. The titration of anacid of unknown concentration with a standard base is shown inFigure 19.22. You can use a similar procedure to find the concentration of abase using a standard acid.

Figure 19.23 shows how the pH of a solution changes during the titra-tion of a strong acid (HCl) with a strong base (NaOH). The pH of the initialacid solution is low. As the base is added, the pH increases because some ofthe acid is neutralized. As the titration approaches the point of neutraliza-tion, at a pH of 7, the pH increases dramatically as hydrogen ions are usedup. Once past the point of neutralization, additional base produces a fur-ther increase of pH. The point of neutralization is the end point of thetitration. At this point, the contents of the beaker consist of only H2O andNaCl, which is the salt derived from the strong acid HCl and the strong baseNaOH, plus a trace of indicator.

Figure 19.22 The titration of an acid with a base is shown here. A known volume of an acid (plus a few drops of phenolphthalein indicator) in a flask is placed beneath a buret filled with a base of known concentration. Base is slowly added from the buret to the acid while the flask is gently swirled.

A change in the color of the indicator signals that neutralization has occurred.

a

b

c

Simulation 26 Simulate the titration of several acids and bases and observe patterns in the pH at equivalence.

Figure 19.23 In this titration of a strong acid with a strong base, 0.10M NaOH is slowly added from a buret to 50.0 mL of 0.10M HCl in the beaker. The equivalence point, the midpoint on the vertical portion of the pH titration curve, occurs at 50.0 mL of NaOH added. Interpreting Diagrams What is true con-cerning [H�] and [OH�] at the equivalence point?

ca b

0

2

4

6

8

10

12

14

50 1000.10M NaOH added (mL)

Titration of a Strong Acid

with a Strong Base

pH Equivalence

pointz

616 Chapter 19


Sample Problem 19.7

Answers32. 25.0 mL KOH × 1.00 molKOH/1000mL KOH × 1 mol HCl/1 mol KOH × 1000 mL HCl/0.45 mol HCl = 56 mL HCl33. 38.5 mL NaOH × 0.150 molNaOH/1000 mL NaOH × 1 molH3PO4/3 mol NaOH = 0.00193 molH3PO4; 0.001923 mol H3PO4/0.0150 LH3PO4 = 0.129M H3PO4

Practice Problems PlusWhat is the molarity of sulfuric acid if 25.0 mL of the solution is neutral-ized by 25.5 mL of 0.50M KOH?(0.26M)

Math HandbookFor a math refresher and practice, direct students to dimensional analysis, page R66.

ASSESSEvaluate UnderstandingAsk students to summarize the titra-tion process in their own words. (Dur-ing a titration, equivalent volumes of an acid and of a base are determined. One is a standard solution with a known con-centration; the concentration of the other solution is not known.)

ReteachHave students use Figures 19.22 and 19.23 to describe how titration works.

Connecting Concepts

Acid-base neutralizations are double displacement reactions. The positive ions are exchanged between the acid and base reactants; the products are a salt and water.

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If your class subscribes to the Inter-active Textbook, use it to review key concepts in Section 19.4.

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Section 19.4 Assessment34. a salt and water35. neutralization36. a. 2 mol b. 0.2 mol37. a. H2SO4 + 2KOH→ 2H2O + K2SO4

b. 2H3PO4 + 3Ca(OH)2 →6H2O + Ca3(PO4)2c. 2HNO3 + Mg(OH)2 →2H2O + Mg(NO3)2

616 Chapter 19

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Math Handbook

Practice Problems

Practice Problems

SAMPLE PROBLEM 19.7

Determining the Concentration of an Acid by TitrationA 25-mL solution of H2SO4 is completely neutralized by 18 mL of1.0M NaOH. What is the concentration of the H2SO4 solution?


Knowns• molarity base � 1.0M NaOH • volume base � 18 mL � 0.018 L• volume acid � 25 mL � 0.025 L• H2SO4(aq) � 2NaOH(aq)¡Na2SO4(aq) � 2H2O(l)

Unknown• molarity acid � ?M H2SO4

Use the molarity to convert the volume of base to moles of base. Usethe mole ratio to find the number of moles of H2SO4. Calculate themolarity by dividing the number of moles of H2SO4 by the number ofliters of H2SO4.The conversion steps are as follows:

L NaOH ¡ mol NaOH ¡ mol H2SO4¡M H2SO4


The [H2SO4] is 0.36M.


Because the volume of acid was greater than the volume of base, theconcentration is less than 1.0M. The answer has two significant figures.

molarity � molesliters � 0.0090 mol

0.025 L � 0.36 M

0.018 L NaOH � 1.0 mol NaOH1 L NaOH �

1 mol H2SO4

2 mol NaOH � 0.0090 mol H2SO4

32. How many milliliters of 0.45M HCl will neutralize 25.0 mL of 1.00M KOH?

33. What is the molarity of H3PO4 if15.0 mL is completely neutral-ized by 38.5 mL of 0.150M NaOH?


34. Key Concept What are the products of a reaction between an acid and a base?

35. Key Concept What occurs at the endpoint of a titration?

36. How many moles of HCl are required to neutralize aqueous solutions of these bases?

a. 2 mol NH3 b. 0.1 mol Ca(OH)2

37. Write complete balanced equations for the following acid-base reactions?

a. H2SO4(aq) � KOH(aq)¡ b. H3PO4(aq) � Ca(OH)2(aq)¡ c. HNO3(aq) � Mg(OH)2(aq)¡

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Problem-Solving 19.33Solve Problem 33 with the help of an interactive guided tutorial.

For help with dimensional analysis, go to page R66.

Types of Reactions Reread the information on types of chemical reactions in Section 11.2. Which type of reaction are all acid-base neutralizations? Explain your answer.



Small-ScaleLAB

Small-ScaleLAB

Ionization Constants of Weak AcidsObjective After completing this activ-ity, students will be able to: measure ionization constants of weak acids

Prep Time 1 hour

Advance Prep

*Can dissolve sodium salt directly in 250 mL of distilled water.

Obtain commercially-prepared pH buffers or prepare pH buffer solutions according to instructions in a reference such as the Handbook of Chemistry and Physics.

Class Time 30 minutes

Safety Make sure that students wear aprons, goggles, and disposable gloves throughout the lab.

Expected Outcome Students mea-sure the pH at which common indica-tors change colors and use their data to calculate the ionization constants of weak acids.

Solution Preparation

*0.04% BTB 100 mg in 16.0 mL 0.01M NaOH, dilute to 250 mL

*0.04% BCG 100 mg in 14.3 mL 0.01M NaOH, dilute to 250 mL

*0.04% BPB 100 mg in 14.9 mL 0.01M NaOH, dilute to 250 mL

*0.04% MCP 100 mg in 26.2 mL 0.01M NaOH, dilute to 250 mL

*0.04% TB 100 mg in 21.5 mL 0.01M NaOH, dilute to 250 mL

0.1% phenol-phthalein

250 mg in 250 mL of 70% 2-propanol

0.05% MO 125 mg in 250 mL

0.02% AYR 50 mg in 250 mL

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Analyze1. pH solutions 1 to 3 are yellow.2. pH solutions 5 to 12 are blue.3. pH solution 4 is green (an intermediate

between yellow and blue).4. The conjugate acid, HBCG, is yellow.5. The conjugate base, BCG-, is blue.6. The equal mixture of HBCG and BCG- is green

at pH = 4.

You’re the Chemist1. Results will vary depending on the indicators

chosen.2. To measure the Ka of a colored acid-base

indicator, mix one drop of the indicator solu-tion with one drop of each pH 1–12 buffer solution. Look for the pH of the color change. This pH is close to the Ka of the acid.

Small-Scale Lab 617

Small-ScaleLAB

Small-ScaleLAB

PurposeTo measure ionization constants of weak acids such as bromocresol green (BCG).

Materials

• ruler • reaction surface • pH buffer and BCG

Procedure1. On separate sheets of paper, draw two grids similar to

the one below. Make each square 2 cm on each side.

2. Place a reaction surface over one of the grids and place one drop of BCG in each square.

3. Place one drop of pH buffer in each square correspond-ing to its pH value.

4. Use the second grid as a data table to record your observations for each solution.

AnalyzeUsing your experimental data, record the answers to the following questions below your data table.

1. What is the color of the lowest-pH solutions?

2. What is the color of the highest-pH solutions?

3. At which pH does the bromcresol green change from one color to the other? At which pH does an intermedi-ate color exist?

Ionization Constants of Weak Acids

An acid-base indicator is usually a weak acid with a char-acteristic color. Because bromcresol green is an acid, it is convenient to represent its rather complex formula as HBCG. HBCG ionizes in water according to the following equation.

The Ka expression is

When [BCG�] � [HBCG], Ka � [H3O�].

4. What color is the conjugate acid of BCG�?

5. What color is the conjugate base of HBCG?

6. What color is an equal mixture of the conjugate acid and conjugate base of bromcresol green? At what pH does this equal mixture occur?

You’re The ChemistThe following small-scale activities allow you to develop your own procedures and analyze the results.

1. Design It! Design and carry out an experiment to measure the Ka of several more acid-base indicators. Record the colors of the conjugate acids (low pH) and the conjugate bases (high pH). Determine the Ka foreach acid.

2. Analyze It! Explain in your own words how to measure the Ka of a colored acid-base indicator. Explain what you would do and how you would interpret your results.

HBCG + H2O � BCG-+ H3O+

(yellow) (blue)

Ka �3BCG- 4 � 3H3O+ 4

3HBCG 41

pH

2 3

4 5 6

7

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618 Chapter 19



Review, Interpreting Graphics• Transparencies, T227–T229• Laboratory Manual, Labs 44, 45• Small-Scale Chemistry Laboratory Manual,

Labs 32, 33


Animation 26, Problem-Solving 19.39, Assessment 19.5

19.5

FOCUSObjectives19.5.1 Describe when a solution of a

salt is acidic or basic.19.5.2 Demonstrate with equations

how buffers resist changes in pH.

Guide for Reading

Build VocabularyParaphrase Students can look up the definition of a buffer in the dictionary. Based on this definition, they can define buffer capacity in their own words.

Reading StrategyKWL Have students construct a KWL chart about buffers, with the three col-umn titles, What I Know about Buffers; What I Want to Know about Buffers; and What I Learned about Buffers. Students will write in the appropriate columns before, during, and after reading this section.

INSTRUCT

Have students read the opening para-graph. Ask, What is the normal pH of human blood? (7.4)

What are the pH limits of human blood before causing death? (between 6.8 and 7.8)

Salt Hydrolysis

Interpreting Graphsa. (a) 8.7, (b) 7b. The number of equivalents of acid

and base are equal.c. In (a), CH3COO– ions react with

water to produce OH– ions. In (b), Cl– ions do not react to produce either H3O+ or OH– ions.

Enrichment QuestionAre there limits to the pH of salt solutions? (In general, salt solutions would not have pH values in the range of strong acids or bases.)

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618 Chapter 19

Titrations of Weak Acid–Strong Base and Strong Acid–Strong Base

a. Identify What is the pH of the equivalence point of each titration?b. Describe Why is the same amount of base used in each titration?c. Apply Concepts Explainwhy the equivalence points of the two titrations are different.

INTERPRETING GRAPHS

19.5 Salts in Solution

The internal pH of most living cells is close to 7. Because the chemical processes of the cell are very sensitive to pH levels, even a slight change in pH can be harmful. Human blood, for exam-ple, normally maintains a pH very close to 7.4. A person cannot survive for more than a few minutes if the blood pH drops to 6.8 or rises to 7.8. In this section, you will learn about chemical processes that ensure that blood is kept near 7.4.

Guide for Reading

Key Concepts• When is the solution of a salt

acidic or basic?• What are the components of a

buffer?

Vocabularysalt hydrolysis

buffer

buffer capacity

Reading StrategyPredicting Based on what you have learned about acid-base neutralization reaction, predict the possible pH range of solutions of salts of strong acids and weak bases. Account for this pH range. Do the same for salts of weak acids and strong bases.

Salt HydrolysisA salt consists of an anion from an acid and a cation from a base. It formsas a result of a neutralization reaction. Although solutions of many saltsare neutral, some are acidic and others are basic. Solutions of sodiumchloride and of potassium sulfate are neutral. A solution of ammoniumchloride is acidic. A solution of sodium ethanoate (sodium acetate) is basic.Figure 19.24 shows a titration curve obtained by adding a solution ofsodium hydroxide, a strong base, to a solution of ethanoic (acetic) acid, aweak acid. An aqueous solution of sodium ethanoate exists at the equiva-lence point.

The pH at the equivalence point is 8.7—basic.

Ethanoic Sodium Sodium Wateracid hydroxide ethanoate

CH3COOH 1aq 2 + NaOH 1aq 2¡ CH3COONa 1aq 2 + H2O 1 l 2

Figure 19.24 The titration curve for a weak acid and a strong base is compared with the titration curve for a strong acid and a strong base.

0

2

4

6

8

10

12

14

(a) Weak acid–strong base

(b) Strong acid–strong base

pH

50 100

Equivalencepoint (a)

Equivalencepoint (b)

y

y

0.10M NaOH added (mL)

CH3COOH

CH3COO�

OH�

Cl�

Na�

Na�



Predicting pH of SolutionsPurpose Students predict the pH of different salt solutions and and give reasons for their predictions.

Materials 1M solutions of NH4NO3, KCl, NaHCO3, Na2SO4; pH meter or uni-versal pH paper

Procedure Have students speculate about the pH of these solutions and explain their reasoning. After some dis-cussion, check the pH of the solutions with a pH meter or universal pH paper. After you obtain the experimental results, have students try to explain the data. (NH4NO3 is acidic because the salt results from a strong acid and a weak base; KCl is neutral because the salt results from a strong acid and a strong base; NaHCO3 is basic because the salt results from a strong base and a weak acid; Na2SO4 is neutral because the salt results from a strong acid and a strong base.)

DiscussPoint out that the salts that give non-neutral solutions in the demo above and in the illustrations are classified as Brønsted-Lowry acids and bases because the cations and anions of these salts accept or donate hydrogen ions from water (hydrolysis). Have stu-dents learn the following:

Strong acid + Strong base→Neutral solution

Strong acid + Weak base→Acidic solution

Weak acid + Strong base→Basic solution

Use VisualsFigures 19.24 and 19.26 Have students compare the figures. Explain that the equivalence point is not always at pH 7. The salts that form may hydrolyze, which results in a reversible reaction that produces a higher concen-tration of hydrogen or hydroxide ions. Ask, What is different about neutral-ization reactions that produce a neu-tral solution at the equivalence point and those that produce acidic or basic solutions? (The strengths of the reactants; when a strong acid and base react, the solution is neutral; when a strong base and a weak acid react, the solution is basic; when a strong acid and a weak base react, the solution is acidic.)

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Section 19.5 Salts in Solution 619

For a strong acid–strong base titration, the pH at the equivalence pointis 7, or neutral. This difference exists because some salts promote hydroly-sis. In salt hydrolysis, the cations or anions of a dissociated salt removehydrogen ions from or donate hydrogen ions to water. Depending on thedirection of the hydrogen-ion transfer, solutions containing hydrolyzingsalts may be either acidic or basic. Hydrolyzing salts are usually derivedfrom a strong acid and a weak base, or from a weak acid and a strong base.Sodium carbonate, washing soda, is the salt of the strong base sodiumhydroxide and carbonic acid, a weak acid. Ammonium nitrate, used in fer-tilizers, is the salt of the weak base ammonia and nitric acid, a strong acid.Soap is the salt of a strong base, usually sodium hydroxide, and stearic acid,a weak acid present in fats. In general, salts that produce acidic solu-tions contain positive ions that release protons to water. Salts that producebasic solutions contain negative ions that attract protons from water.

Sodium ethanoate (CH3COONa) is the salt of a weak acid (ethanoic

acid, CH3COOH) and a strong base (sodium hydroxide, NaOH). In solution,

the salt is completely ionized.

The ethanoate ion is a Brønsted-Lowry base, which means it is a hydrogen-ion acceptor. It establishes an equilibrium with water, forming electrically neu-tral ethanoic acid and negative hydroxide ions.

This process is called hydrolysis because it splits a hydrogen ion off a water mol-ecule. The resulting solution contains a hydroxide-ion concentration greaterthan the hydrogen-ion concentration. Thus the solution is basic.

Ammonium chloride (NH4Cl) is the salt of a strong acid (hydrochloricacid, HCl) and a weak base (ammonia, NH3). It is completely ionized insolution.

NH4Cl(aq)¡ NH4�(aq) � Cl�(aq)

Sodium ethanoate Ethanoate ion Sodium ion

CH3COONa 1aq 2¡ CH3COO- 1aq 2 + Na+ 1aq 2

CH3COOH1aq2 + OH-1aq2EPJCH3COO- 1aq 2 + H2O 1l 2(hydrogen-ion (hydrogen-ion (makes the

acceptor, Brønsted- donor, Brønsted- solutionLowry base) Lowry acid) basic)

Figure 19.25 Vapors of the strong acid HCl(aq) and the weak base NH3(aq) combine to form the acidic white salt ammonium chloride (NH4Cl).

Figure 19.26 A few drops of universal indicator solution have been added to each of these 0.10M aqueous salt solutions. The color of the indicator shows the following. Ammonium chloride (NH4Cl), is acidic (pH about 5.3). Sodium chloride (NaCl) is neutral (pH 7).

Sodium ethanoate (CH3COONa) is basic (pH about 8.7).

a

b

c

a b c

620 Chapter 19


Facts and FiguresBlood BuffersYour body functions properly only when the pH of your blood lies between 7.35 and 7.45. A blood pH outside these narrow limits can be life-threatening! To keep your blood at the proper pH, your blood contains buffers. The most important blood buffer consists of carbonic acid and hydrogen carbonate ions. Hydrogen ions produced by chemical reactions in your body are the main threat to your blood’s pH. This buffer maintains the

desired blood pH by removing excess hydrogen ions.

As long as there are hydrogen carbonate ions available, the excess hydrogen ions are removed, and the pH of the blood changes very little.

HCO3-(aq) + H+(aq) H2CO3(aq)

Hydrogencarbonate ion

Hydrogenion

Carbonicacid

Buffers

Word OriginsA buffer zone is a neutral area set up around a site to prevent any changes or to the site being investi-gated. In order for the correct con-clusions to be drawn, no evidence can be corrupted.


Comparing Commerical BuffersPurpose To demonstrate that the neu-tralizing capacities of commercial buff-ers differ, students compare antacid tablets, buffered aspirin, and aspirin.

Materials Antacid tablets, flasks and burets, 0.5M HCl and NaOH, distilled water, methyl red, bromothymol blue

Procedure Set up pairs of flasks con-taining 100 mL of distilled water to test each product. Dissolve equal masses of the tablets in each flask. Add a few drops of methyl red to one of the paired flasks and titrate with 0.5M NaOH. To the other flask, add a few drops of bromothymol blue and titrate with 0.5M HCl. Titrate until the end point is reached. Explain that the buff-ering capacity is directly related to the amount of acid or base required to reach the end point. Have students compare the buffering capacities of the tablets tested.

Use VisualsFigure 19.27 Explain to students that buffers are solutions in which the pH stays relatively constant when small amounts of either acid or base are added. Point out that in the buffered solution on the left, there is little change in pH.

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620 Chapter 19

Word Origins

The ammonium ion (NH4�) is a strong enough acid to donate a hydrogen ion

to a water molecule, although the equilibrium is strongly to the left.

NH31aq2 + H3O+(aq)EPJNH4

+ 1aq 2 + H2O 1l 2(hydrogen-ion (hydrogen-ion (makes the

donor, Brønsted- acceptor, Brønsted- solutionLowry acid) Lowry base) acidic)

This process is also called hydrolysis. It results in the formation of unionizedammonia and hydronium (hydrogen) ions. The [H3O�] is greater than the[OH�]. Thus, a solution of ammonium chloride is acidic. To determine if asalt solution is acidic or basic, remember the following rules:

Strong acid � Strong base¡ Neutral solution

Strong acid � Weak base¡ Acidic solution

Weak acid � Strong base¡ Basic solution

BuffersThe addition of 10 mL of 0.10M sodium hydroxide to 1 L of pure waterincreases the pH by 4.0 pH units (from 7.0 to 11.0). A solution containing0.20 mol/L each of ethanoic acid and sodium ethanoate has a pH of 4.76.When moderate amounts of either acid or base are added to this solution,however, the pH changes little. The addition of 10 mL of 0.10M sodiumhydroxide to 1 L of this solution, for example, increases the pH by only0.01 pH unit, from 4.76 to 4.77. Figure 19.27 shows what happens when1.0 mL of 0.01M HCl solution is added to an unbuffered solution.

The solution of ethanoic acid and sodium ethanoate is an example of atypical buffer. A buffer is a solution in which the pH remains relatively con-stant when small amounts of acid or base are added. A buffer is a solu-tion of a weak acid and one of its salts, or a solution of a weak base and one ofits salts.

A buffer solution is better able to resist drastic changes in pH than ispure water. Figure 19.28 illustrates how a buffer works. Ethanoic acid(CH3COOH) and its anion (CH3COO�) act as reservoirs of neutralizingpower. They react with any hydroxide ions or hydrogen ions added to thesolution. For example, consider the buffer solution in which the sodiumethanoate (CH3COONa) is completely ionized.

Sodium Ethanoate Sodiumethanoate ion ion

CH3COONa 1aq 2¡ CH3COO- 1aq 2 + Na+ 1aq 2

Figure 19.27 A buffer is a solution in which the pH remains relatively constant. The indicator shows that the buffered solution on the left and the unbuffered solution on the right are basic—pH about 8. After the addition of 1.0 mL of 0.01MHCl solution, the pH of the buffered solution shows no visible change. The pH of the unbuffered solution, however, is now about 3—the solution is acidic.

a

b

Buffer comes from the Old English word buff, meaning “firmly or sturdily.” A buffer solution resists changes in its pH even when acidic or basic solutions are added to it. What does it mean if authorities set up a buffer zone around a fire inves-tigation site?

a

buffered unbuffered

b

buffered unbuffered


DiscussExplain that buffers present a practical application of concepts in this chapter. The limited dissociation of weak acids and bases gives these substances the ability to act as buffers. This ability is greatly expanded if the salt of the acid or base is added to the solution. Buffer action utilizes the equilibrium estab-lished when a weak acid or base is combined with its salt. When a base is added to a buffered solution, the acidic form removes hydroxide ions from the solution. When an acid is addedto a buffered solution, the basic form removes hydrogen ions from the solution.

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Section 19.5 Salts in Solution 621

When an acid is added to the solution, the ethanoate ions (CH3COO�) actas a hydrogen-ion “sponge.” This creates ethanoic acid, which does notionize extensively in water; thus the pH does not change appreciably.

Ethanoate Hydrogen Ethanoic acidion ion

CH3COO-1aq 2 + H+1aq 2 3rs CH3COOH1aq 2

A base is a source of hydroxide ions. When a base is added to the solution,the ethanoic acid and the hydroxide ions react to produce water and theethanoate ion.

The ethanoate ion is not a strong enough base to accept hydrogen ions

Ethanoic acid Hydroxide Ethanoate Waterion ion

CH3COOH1aq 2 + OH- 1aq 2 3rs CH3COO- 1aq 2 + H2O(l)

from water extensively. Again, the pH does not change very much.An ethanoate buffer solution cannot control the pH when too much

acid is added, because no more ethanoate ions are present to accepthydrogen ions. The ethanoate buffer also becomes ineffective when toomuch base is added. Do you know why this is so? When too much base isadded, no more ethanoic acid molecules are present to donate hydrogenions. When too much acid or base is added, the buffer capacity of a solutionis exceeded. The buffer capacity is the amount of acid or base that can beadded to a buffer solution before a significant change in pH occurs.

Two buffer systems are crucial in maintaining human blood pH within avery narrow range (pH 7.35–7.45). One is the carbonic acid–hydrogen carbo-nate buffer system. The other is the dihydrogen phosphate–monohydrogenphosphate buffer system. Table 19.10 lists several important buffer systems.

Figure 19.28 The buffer described here is made of ethanoic acid (CH3COOH) and sodium ethanoate, which is the source of ethanoate ions (CH3COO�). To begin, the concentrations of ethanoic acid and ethanoate ions are equal. When either H� or OH� is added, the buffer produces additional ethanoic acid or ethanoate ions. In both situations, the ratio of [CH3COOH] to [CH3COO�] and consequently, the pH, changes very little.

CH3COO� CH3COOHH�

CH3COO� CH3COOH CH3COO� CH3COOH

CH3COOH H� � CH3COO�

OH�

CH3COOH � OH� CH3COO� � H2O

Table 19.10

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Animation 26 Discover the chemistry behind buffer action.

Important Buffer Systems

Buffer pH

Buffer name Buffer species (components 0.1M)

Ethanoic acid–ethanoate ion CH3COOH/CH3COO� 4.76

Dihydrogen phosphate ion–hydrogenphosphate ion H2PO4�/HPO4

2� 7.20

Carbonic acid–hydrogen carbonate ion(solution saturated with CO2)

H2CO3/HCO3� 6.46

Ammonium ion–ammonia NH4�/NH3

9.25

622 Chapter 19



Answers38. a. HPO4

2- + H+ → H2PO4-

b. H2PO4- + OH- → HPO4

2- + H2O39. CH3COO- + H+ → CH3COOH

ASSESSEvaluate UnderstandingHave students explain in their own words what determines whether a solu-tion containing a hydrolyzing salt is acidic or basic. (Salts formed from a strong acid and strong base are neutral. Those formed from a strong base and weak acid are basic; those formed from a strong acid and weak base are acidic.) Ask students to explain how ethanoate ions can effectively remove hydrogen ions from solution. (The ethanoate ions are a Brønsted–Lowry base that accept the hydrogen ions to form ethanoic acid.)

ReteachPoint out that a hydrolysis reaction always uses one part of the water mol-ecule and leaves behind the other. Use equations to show how the salt of a weak acid always reacts with water to produce hydroxide ions and the salt of a weak base always reacts to produce hydrogen ions. Remind students that chemical buffers are equilibrium sys-tems that tend to resist change due to external influences.

The most important blood buffer consists of carbonic acid and hydro-gen carbonate ions. HCO3

–(aq) + H+ (aq) EP

H2CO3(aq) The buffer maintains the desired blood pH by removing or providing hydrogen ions.

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Section 19.5 Assessment40. acidic solution: salt of a strong acid and

weak base; basic solution: salt of a weak acid and a strong base

41. a weak acid and one of its salts or a weak base and one of its salts

42. d43. NH3 + H+ → NH4

+

NH4+ + OH- → NH3 + H2O

622 Chapter 19

Practice Problems


Using Equations to Illustrate the Action of a BufferShow how the carbonic acid–hydrogen carbonate buffer can “mop up” added hydrogen ions and hydroxide ions.

When an acid is added to the buffer, it reacts with HCO3

� (a proton acceptor) to produce un-ionized carbonic acid. Again the pH changes very little.

Carbonic Hydroxide Hydrogen Wateracid ion carbonate ion

H2CO31aq 2 + 1OH- 2 1aq 2 3rs HCO3- 1aq 2 + H2O1l 2

Hydrogen Hydrogen Carboniccarbonate ion ion acid

HCO3-1aq 2 + H+1aq 2 3rs H2CO31aq 2

38. Write reactions to show what happens when the following occur.a. Acid is added to a solution of HPO4

2�.b. Base is added to a solution of H2PO4

�.

39. Write an equation that shows what happens when acid is added to the ethanoic acid–ethanoate buffer.

40. Key Concept What type of salt produces an acidic solution? A basic solution?

41. Key Concept What substances are combined to make a buffer?

42. Which of these salts would form an acidic aqueous solution?

a. KC2H3O2 b. LiCl c. NaHCO3 d. (NH4)2SO4

43. Using equations, show what happens when acid is added to an ammonium ion–ammonia buffer. What happens when base is added?

Analyze Identify the relevant concepts.

The definition of a buffer, in terms of its ability to act as a reservoir of neutralizing power, should be illustrated. One compo-nent in a buffer can react with H�; this com-ponent is a proton acceptor. Another component can react with OH�; this compo-nent is a proton donor.

Solve Apply concepts to this situation.

The carbonic acid–hydrogen carbonate buff-er is a solution of carbonic acid (H2CO3) and hydrogen carbonate ions (HCO3

�). When a base is added to this buffer, it reacts with H2CO3 (a proton donor) to produce neutral water. The pH changes very little.

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Practice Problems


Explanatory Paragraph To maintain a proper pH, your blood contains buffers. Research the buffer systems in your blood and write a paragraph on how one of them works.


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Rescuing Crumbling BooksExplain that when aluminum sulfate [Al2(SO4)3)] combines with water vapor in the atmosphere, small amounts of sulfuric acid form on the pages of books, causing the paper to turn yellow and brittle. Humidity has an important influence on how much acid forms and how quickly books deteriorate. There-fore, it is important for libraries to have adequate climate control systems.

After students have read the feature, lead them in a discussion about the advantages and disadvantages of hav-ing the reference materials in your library available as bound books, on microfilm, and on audiotapes.

Answers to...InferringLibrary books are handled by a lot of people over time. It is important that the books are preserved. Also, it would cost less to do many books at one time, as done in the deacidifica-tion method described, than indi-vidual treatments using a different method.

Treating Books in MassFor an estimated cost of $6 to $10 per book, hundreds of volumes can be treated at once. If books had to be treated page by page, deacidification could cost $1000 per book.

Facts and Figures

Technology and Society 623

Rescuing Crumbling Books

Millions of books printed since the mid-nineteenth

century slowly decay as they sit on library shelves.

The cause is the acidity of the paper, which comes

from alum (aluminum sulfate), used to prevent inks

from soaking into paper. Libraries around the world

are working with chemists to find ways to remove

the acid from these books. One such method

involves placing a book in a vacuum chamber,

removing all the moisture from the pages, and

applying a gas to neutralize the acid. Inferring Why

might a library decide to use the method shown below?

Deacidification In the tank system above, many books are treated at the same time. The machine moves the books gently in deacidification fluid, bathing each page. After the fluid is pumped away, a vacuum dries the treated books.

Deterioration The pages of this book are so fragile that they crumble at a touch. Deacidifying the paper will extend the life of the book.

Restoration King James of England issued this charter to establish a

colony in Virginia and find a water passage to the Pacific Ocean. It was severely damaged from prolonged exposure to moisture and a poor storage

environment, after which chemists washed and

deacidified each page.

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