Acids, Bases, and Salts: Arrhenius and Brønsted-Lowry Theories, Study Guides, Projects, Research of Chemistry

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Artists often use hydrofluoric acid to etch designs on glass.

Acids, Bases, and Salts

622

CHEMYSTERY

Paper Trail

The invention of the

printing press in 1440

increased the number

of books that could be

printed and the need

for paper on which to

print these books. The main

ingredient in paper was cotton, which is almost pure

cellulose. This type of paper is referred to as rag-based

paper. When books printed on high-quality rag-based

paper are stored correctly, they can last for hundreds

of years.

By 1880, the demand for paper was so great that

printers switched from rag-based paper to wood-based

paper. Over time, the wood-based paper tended to

yellow and crack. Sometimes the paper was so brittle

that it crumbled when touched.

Why did a change in the content of paper cause

such a dramatic change in the properties of the

paper? As you read about acids and bases and

their reactions, think about what could cause paper

to crumble.

Take a video field trip with the Untamed Science crew

to learn about another situation in which acid-base

reactions are important.

TEKS

FO

CUSON

10H

Texas Essential Knowledge and Skills

READINESS TEKS: 10H Understand and differentiate among acid-base reactions, precipitation reactions, and oxidation-reduction reactions.

SUPPORTING TEKS: 10G Define acids and bases and distinguish between Arrhenius and Brønsted-Lowry definitions and predict products in acid base reactions that form water. 10I Define pH and use the hydrogen or hydroxide ion concentrations to calculate the pH of a solution. 10J Distinguish between degrees of dissociation for strong and weak acids and bases.

TEKS: 3E Describe the connection between chemistry and future careers. 2H Organize, analyze, evaluate, make inferences, and predict trends from data.

Acids, Bases, and Salts 623

Acids, Bases, and Salts 625

Figure 18.2 Hydrochloric Acid Hydrochloric acid is actually an aqueous solution of hydrogen chloride. Hydrogen chloride forms hydronium ions, making this compound an acid. Explain Why does hydrogen chloride release a hydrogen ion when dissolved in water?

Arrhenius Acids Table 18.1 lists six common acids. They vary in the

number of hydrogens that can form hydrogen ions. A hydrogen atom that can form a hydrogen ion is ionizable. Nitric acid (HNO 3 ) has one ionizable hydrogen, so nitric acid is classified as a monoprotic acid. The prefix mono- means “one,” and the stem protic means that a hydrogen ion is a proton. Acids that contain two ionizable hydrogens, such as sulfuric acid (H 2 SO 4 ), are called diprotic acids. Acids that contain three ionizable hydrogens, such as phosphoric acid (H 3 PO 4 ), are called triprotic acids. Not all compounds that contain hydrogen are acids. Also, some hydrogens in an acid may not form hydrogen ions. Only a hydrogen that is bonded to a very electronegative element can be released as an ion. Recall that such bonds are highly polar. When a compound that contains such bonds dissolves in water, it releases hydrogen ions. An example is the hydrogen chloride molecule, shown below.

However, in an aqueous solution, hydrogen ions are not present. Instead, the hydrogen ions are joined to water molecules as hydronium ions. A hydronium ion (H 3 O∙)^ is the ion that forms when a water molecule gains a hydrogen ion. As seen in Figure 18.2, hydrogen chloride ionizes to form an aqueous solution of hydronium ions and chloride ions.

Methane (CH 4 ) is an example of a hydrogen-containing com- pound that is not an acid. The four hydrogen atoms in methane are attached to the central carbon atom by weakly polar C (^) — H bonds. Thus, methane has no ionizable hydrogens and is not an acid. Ethanoic acid (CH 3 COOH), which is commonly called acetic acid, is an example of a molecule that contains both hydrogens that do not ionize and a hydrogen that does ionize. Although its molecules contain four hydrogens, etha- noic acid is a monoprotic acid. The structural formula shows why.

Ethanoic acid (CH 3 COOH)

H

H C C O H

H

O

The three hydrogen atoms attached to carbon atom are bound by weakly polar C (^) — H bonds. These hydrogens do not ionize. Only the hydrogen bonded to the highly electronegative oxygen can be ionized. For complex acids, you need to look at the structural formula to recog- nize which hydrogens can be ionized.

H—Cl( g) H+(aq) + Cl−(aq) Hydrogen chloride

Hydrogen ion

Chloride ion (hydrochloric acid)

δ+ δ− (^) H 2 O

HCl Hydrogen chloride

H 2 O Water

H 3 O+ Hydronium ion

Cl– Chloride ion

• –

Table 18.

Some Common Acids

Name Formula Hydrochloric acid HCl Nitric acid HNO 3 Sulfuric acid H 2 SO 4 Phosphoric acid H 3 PO 4 Ethanoic acid CH 3 COOH Carbonic acid H 2 CO 3

626 Chapter 18 • Lesson 1

Figure 18.3 Clogged Drains Sometimes water backs up in a sink because the drain is clogged. A plumber can take apart the pipes to remove a clog, or a drain cleaner containing sodium hydroxide can be used to dissolve the clog.

Table 18. Some Common Bases

Name Formula Solubility in water Sodium hydroxide NaOH High Potassium hydroxide KOH High Calcium hydroxide Ca(OH) 2 Very low Magnesium hydroxide Mg(OH) 2 Very low

Arrhenius Bases Table 18.2 lists four common bases. You may

be familiar with the base sodium hydroxide (NaOH), which is also known as lye. Sodium hydroxide is an ionic solid. It dissoci- ates into sodium ions and hydroxide ions in aqueous solution.

NaOH(s) Sodium hydroxide

Na+(aq) Sodium ion

• (^) OH−(aq) Hydroxide ion

Sodium hydroxide is extremely caustic. A caustic substance can burn or eat away materials with which it comes in con- tact. This property is the reason that sodium hydroxide is a major component of products that are used to clean clogged drains. Figure 18.3 shows a drain cleaner that contains sodium hydroxide. Potassium hydroxide (KOH) is another ionic solid. It dissociates to produce potassium ions and hydroxide ions in aqueous solution.

Sodium and potassium are Group 1A elements. Elements in Group 1A, the alkali metals, react violently with water. The products of these reactions are aqueous solutions of a hydroxide and hydrogen gas. The following equation summarizes the reac- tion of sodium with water.

2Na(s) Sodium metal

• 2H 2 O(l) Water

2NaOH(aq) Sodium hydroxide

• H 2 ( g) Hydrogen

Sodium hydroxide and potassium hydroxide are very soluble in water. Thus, making concentrated solutions of these compounds is easy. The solutions would have the typically bitter taste and slippery feel of a base. However, these are not proper- ties that you would want to confirm. The solutions are extremely caustic to the skin. They can cause deep, painful, slow-healing wounds if not immediately washed off.

H 2 O

KOH(s) K+(aq) + OH−(aq) Potassium hydroxide

Potassium ion

Hydroxide ion

H 2 O Q: Visitors to Bracken Cave wear protective gear to keep ammonia gas out of their eyes and respiratory tracts. Think about the properties of bases. Why are high levels of ammonia harmful?

CHEMISTRY (^) & YOU

628 Chapter 18 • Lesson 1

Figure 18.6 Sulfuric Acid When sulfuric acid and water react, they form hydronium ions and hydrogen sulfate ions. Identify Which product is the conjugate acid, and which is the conjugate base?

Conjugate Acids and Bases All gases become less soluble in water as the

temperature rises. Thus, when the temperature of an aqueous solution of ammonia is increased, ammonia gas is released. This release acts as a stress on the system. In response to this stress, NH 4 +^ reacts with OH−^ to form more NH 3 and H 2 O. In the reverse reaction, ammonium ions donate hydrogen ions to hydroxide ions. Thus, NH 4 +^ (the donor) acts as a Brønsted-Lowry acid, and OH−^ (the acceptor) acts as a Brønsted-Lowry base. In essence, the reversible reaction of ammonia and water has two acids and two bases.

In the equation, the products of the forward reaction are distinguished from the reactants by the use of the adjective conjugate. This term comes from a Latin word meaning “to join together.” A conjugate acid is the ion or molecule formed when a base gains a hydrogen ion. In the reaction above, NH 4 +^ is the conjugate acid of the base NH 3. A conjugate base is the ion or molecule that remains after an acid loses a hydrogen ion. In the reaction above, OH−^ is the conjugate base of the acid H 2 O. Conjugate acids are always paired with a base, and conjugate bases are always paired with an acid. A conjugate acid-base pair consists of two ions or molecules related by the loss or gain of one hydrogen ion. The ammonia molecule and the ammonium ion are a conjugate acid-base pair. The water molecule and the hydroxide ion are also a conjugate acid-base pair.

The dissociation of hydrogen chloride in water provides another example of conjugate acids and bases.

In this reaction, hydrogen chloride is the hydrogen-ion donor. Thus, it is by definition a Brønsted-Lowry acid. Water is the hydrogen-ion acceptor and a Brønsted-Lowry base. The chloride ion is the conjugate base of the acid HCl. The hydronium ion is the conjugate acid of the base water. Figure 18.6 shows the reaction that takes place when sulfuric acid dis- solves in water. The products of this reaction are hydronium ions and hydro- gen sulfate ions. Use the figure to identify the two conjugate acid-base pairs.

NH 3 (aq) + H 2 O(l) NH 4 +(aq) + OH−(aq) Base Acid Conjugate acid

Conjugate base

NH 3 (aq) + H 2 O(l) NH 4 +(aq) + OH−(aq) Base Acid Conjugate acid

Conjugate base

HCl( g) + H 2 O(l) H 3 O+(aq) + Cl−(aq) Acid Base Conjugate acid

Conjugate base

H 2 SO 4 Sulfuric acid

H 2 O Water

H 3 O+ Hydronium ion

HSO 4 – Hydrogen sulfate ion

• –

Acids, Bases, and Salts 629

Amphoteric Substances Look at Table 18.3. Note that water appears

in both the list of acids and the list of bases. Sometimes water accepts a hydrogen ion. At other times, it donates a hydrogen ion. How water behaves depends on the other reactant. A substance that can act as either an acid or a base is said to be amphoteric. Water is ampho- teric. In the reaction with hydrochloric acid, water accepts a proton and is therefore a base. In the reaction with ammonia, water donates a proton and is therefore an acid. Look for two other substances in Table 18.3 that are amphoteric.

Lewis Acids and Bases

How did Lewis define an acid and a base?

The work that Gilbert Lewis (1875–1946) did on bonding led to a new concept of acids and bases. According to Lewis, an acid accepts a pair of electrons and a base donates a pair of electrons during a reaction. This definition is more general than those offered by Arrhenius or by Brønsted and Lowry. A Lewis acid is a substance that can accept a pair of electrons to form a covalent bond. Similarly, a Lewis base is a substance that can donate a pair of electrons to form a covalent bond. The Lewis definitions include all the Brønsted-Lowry acids and bases. Consider the reaction of H+^ and OH−. The hydrogen ion donates itself to the hydroxide ion. Therefore, H+^ is a Brønsted-Lowry acid and OH−^ is a Brønsted-Lowry base. The hydroxide ion can bond to the hydrogen ion because it has an unshared pair of electrons. Thus, OH−^ is also a Lewis base, and H+, which accepts the pair of electrons, is a Lewis acid.

Lewis acid

Lewis base

H+^ + −^ O H O

H H

A second example of a reaction between a Lewis acid and a Lewis base is what happens when ammonia dissolves in water. Hydrogen ions from the dissociation of water are the electron-pair acceptor and the Lewis acid. Ammonia is the electron-pair donor and the Lewis base. Table 18.4 compares the definitions of acids and bases. The Lewis definition is the broadest. It extends to compounds that the Brønsted-Lowry theory does not classify as acids and bases. Sample Problem 18.1 provides some examples of those compounds.

Table 18.

Acid-Base Definitions

Type Acid Base

Arrhenius (^) H+^ producer OH−^ producer

Brønsted-Lowry (^) H+^ donor H+^ acceptor Lewis electron-pair acceptor electron-pair donor

READING SUPPORT

Building Vocabulary: Prefixes The prefix amphi- is from a Greek word meaning “of both kinds.” An amphibian is an animal that is capable of living both on land and in the water. What does it mean to describe an airplane as amphibious_?_

Table 18.

Some Conjugate Acid-Base Pairs

Acid Base

HCl (^) Cl−

H 2 SO 4 HSO 4 −

H 3 O+^ H 2 O

HSO 4 −^ SO 42 − CH 3 COOH (^) CH 3 COO−

H 2 CO 3 HCO 3 −

HCO 3 −^ CO 32 −

NH 4 +^ NH 3

H 2 O (^) OH−

TEKS 10G

CHEMISTRY (^) & YOU

631

18.2 Hydrogen Ions and Acidity

Q: What factors do you need to control so a fish has healthy water to live in?

Goldfish can live for 20 years or more in an aquarium if the conditions are

right. The water in the aquarium must be cleaned regularly. You must also

control the temperature of the water. In this lesson, you will study another

factor that affects the ability of fish to survive.

Hydrogen Ions From Water How are [H ∙ ] and [OH ∙ ] related in an aqueous solution? Water molecules are highly polar and are in constant motion, even at room temperature. On occasion, the collisions between water molecules are ener- getic enough for a reaction to occur. When this happens, a hydrogen ion is transferred from one water molecule to another, as illustrated below. A water molecule that gains a hydrogen ion becomes a hydronium ion (H 3 O+). A water molecule that loses a hydrogen ion becomes a hydroxide ion (OH−).

Self-Ionization of Water The reaction in which water molecules produce

ions is called the self-ionization of water. This reaction can be written as fol- lows, where a water molecule dissociates to form two ions, H+^ and OH−.

In water or in an aqueous solution, hydrogen ions are always joined to water molecules as hydronium ions. Yet chemists may still refer to these ions as hydrogen ions or even protons. In this textbook, either H+^ or H 3 O+^ is used to represent hydrogen ions in aqueous solution. The self-ionization of water occurs to a very small extent. In pure water at 25°C, the concentration of hydrogen ions is only 1 × 10 −^7 M. The concentra- tion of OH−^ is also 1 × 10 −^7 M because the numbers of H+^ and OH−^ ions are equal in pure water. Any aqueous solution in which [H+] and [OH−] are equal is a neutral solution.

H 2 O(l) H+(aq) + OH−(aq) Hydrogen ion Hydroxide ion

Key Questions

How are [H ∙ ] and [OH ∙ ] related in an aqueous solution?

How is pH used to classify a solution as neutral, acidic, or basic?

What are two methods that are used to measure pH?

Vocabulary

• self-ionization
• neutral solution
• ion-product constant for water ( K w )
• acidic solution
• basic solution
• pH

H 2 O Water molecule

H 2 O Water molecule

H 3 O+ Hydronium ion

OH– Hydroxide ion

• (^) –

ELPS 4.G. Read the sections titled “Acidic Solutions” and “Basic Solutions” on page 632. Retell in your own words what makes a solution acidic or basic. Use your summary to explain the pH levels listed for the various solutions in Table 18.5 on page 634.

In this lesson, you will learn about pH, including how to calculate the pH of a solution using the hydrogen-ion or hydroxide-ion concentrations (TEKS 10I). You will also learn about the connection between chemistry and future careers (TEKS 3E).

632 Chapter 18 • Lesson 2

Figure 18.7 Aged by Acid Sometimes guitar players want a new guitar to look like it is old or “vintage.” The guitarist can remove the shiny new metal parts of the guitar and expose them to hydrochloric acid. The acid will make the metal parts look dull. Both of the guitars in the photo below are new, but the bottom one has been aged with acid.

Ion-Product Constant for Water The ionization of water

is a reversible reaction. Adding either hydrogen ions or hydroxide ions to an aqueous solution is a stress to the system. In response, the equilibrium will shift toward the formation of water. The concentration of the other ion will decrease. In any aqueous solution, when [H+] increases, [OH−] decreases. Likewise, when [H+] decreases, [OH−] increases.

H+(aq) + OH−(aq) H 2 O(l)

For aqueous solutions, the product of the hydrogen- ion concentration 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. When substances are added to water, the concentrations of H+^ and OH−^ may change. However, the product of [H+] and [OH−] does not change. The product of the concentrations of the hydrogen ions and hydroxide ions in water is called the ion-product constant for water ( K w ).

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

HCl(aq) H+(aq) + Cl−(aq)

In hydrochloric acid, the hydrogen-ion concentration is greater than the hydroxide-ion concentration. (The hydrox- ide ions come from the self-ionization of water.) A solution in which [H+] is greater than [OH−] is an acidic solution. In acidic solutions, the [H+] is greater than 1 × 10 −^7 M. Figure 18.7 shows a guitar that was artificially aged by using hydrochloric acid.

Basic Solutions When sodium hydroxide dissolves in water, it forms hydroxide ions in solution.

NaOH(aq) Na+(aq) + OH−(aq)

In such a solution, the hydrogen-ion concentration is less than the hydroxide-ion concentration. Remember, the hydrogen ions are present from the self-ionization of water. A basic solution is one in which [H+] is less than [OH−]. The [H+] of a basic solution is less than 1 × 10 −^7 M. Basic solutions are also known as alkaline solutions.

K w ∙ [H∙] ∙ [OH∙] ∙ 1.0 ∙ 10 ∙^14

634 Chapter 18 • Lesson 2

The pH Concept How is pH used to classify a solution as neutral, acidic, or basic? Expressing hydrogen-ion concentration in molarity is not practical. A more widely used system for expressing [H+] is the pH scale, proposed in 1909 by the Danish scientist Søren Sørensen. The pH scale ranges from 0 to 14.

Hydrogen Ions and pH Chemists define the pH of a solution as the nega-

tive logarithm of the hydrogen-ion concentration. The pH may be represented mathematically using the following equation:

In pure water or a neutral solution, the [H+] = 1 × 10 −^7 M, and the pH is 7.

If the [H+] of a solution is greater than 1 × 10 −^7 M, the pH is less than 7.0. If the [H+] of the solution is less than 1^ ×^10 −^7 M, the pH is greater than 7.0. A solution with a pH less than 7.0 is acidic. A solution with a pH of 7.0 is neutral. A solution with a pH greater than 7.0 is basic. Table 18. summarizes the relationship among [H+], [OH−], and pH. It also indicates the pH values of some common aqueous systems, including milk and blood.

Table 18.

Relationships Among [H∙], [OH∙], and pH

[H∙] (mol/L)

[OH∙] (mol/L)

Increasing acidity

1 × 10 0 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 Neutral (^1) × 10 −^7 1 × 10 −^7

Increasing basicity

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 × 10 0

pH ∙ ∙log[H∙]

pH = −log (1 × 10 −^7 ) = −(log 1 + log 10−^7 ) = −(0.0 + (−7.0)) = 7.

pH 1 M HCl

0.1 M HCl Gastric juice Lemon juice

Tomato juice Black coffee

Milk Pure water Blood Seawater

Milk of magnesia Household ammonia 0.1 M NaOH 1 M NaOH

When [H ∙ ] is given in the format 1 ∙ 10 ∙ n, it’s easy to find the pH. It’s just the absolute value of the exponent n. Also, note that [H ∙ ] ∙ [OH ∙ ] always equals 1 ∙ 10 ∙^14.

Q: In an aquarium, the pH of water is another factor that affects the ability of fish to sur- vive. Most freshwater fish need a slightly acidic or neutral pH. For a saltwater tank, the ideal pH is slightly basic. What might explain this difference in the ideal pH range?

CHEMISTRY (^) & YOU

TEKS 10I

Sample Problem 18.

KNOWN

[H ∙ ] ∙ 4.2 ∙ 10 ∙^10 M

UNKNOWN

pH ∙?

Round the pH to two decimal places because the hydrogen- ion concentration has two significant figures.

pH ∙ ∙ log [H ∙ ]

pH ∙ ∙ log (4.2 ∙ 10 ∙^10 ) pH ∙ ∙ ( ∙ 9.37675) pH ∙ 9. pH ∙ 9.

Substitute the known [H∙] and use the log function on your calculator to calculate the pH.

Start with the equation for finding pH from [H∙].

Acids, Bases, and Salts 635

Calculating pH From [H∙] When you use the hydrogen-ion concentration to

calculate the pH of a solution, expressing [H∙] in scientific notation can make the calculation easier. For example, you would rewrite 0.0010M as 1.0 × 10 −^3 M. The coefficient 1.0 has two significant figures. The pH for a solution with this concentration is 3.00. The two numbers to the right of the decimal point represent the two significant figures in the concentration. It is easy to find the pH for solutions when the coefficient is 1.0. The pH of the solution equals the exponent, with the sign changed from minus to plus. For example, a solution with [H+] = 1 × 10 −^2 M has a pH of 2.0. When the coefficient is a number other than 1, you will need to use a calculator with a log function key to calculate pH.

Calculating pH From [H∙] What is the pH of a solution with a hydrogen-ion concentration of 4.2 × 10 −^10 M?

➊ Analyze^ List the known and the unknown.^ To find the pH from the hydrogen-ion concentration, you use the equation pH = −log[H+].

➋ Calculate^ Solve for the unknown.

12. Calculate the pH of each solution. a. [H+] = 0.045M b. [H+] = 8.7 × 10 −^6 M c. [H+] = 0.0015M d. [H+] = 1.2 × 10 −^3 M 13. Use the hydrogen-ion concentrations to calculate the pH of the following solutions. a. [H+] = 1.0 × 10 −^12 M b. [H+] = 1 × 10 −^4 M

➌ Evaluate^ Does the result make sense?^ The value of the hydrogen-ion concentration is between 1 × 10 −^9 M and 1 × 10 −^10 M. So, the calculated pH should be between 9 and 10, which it is.

TEKS 10I

Sample Problem 18.

KNOWNS UNKNOWN

[OH ∙ ] ∙ 4.0 ∙ 10 ∙^11 M

KW ∙ 1.0 ∙ 10 ∙^14

pH ∙?

[H ∙ ] ∙

Kw [OH ∙ ]

Kw ∙ [OH ∙ ] ∙ [H ∙ ]

pH ∙ ∙ log [H ∙ ] ∙ ∙ log (2.5 ∙ 10 ∙^4 )

∙ 2.5 ∙ 10 ∙^4 M

1.0 ∙ 10 ∙^14

4.0 ∙ 10 ∙^11

[H ∙ ] ∙ ∙ 0.25 ∙ 10 ∙^3 M

Acids, Bases, and Salts 637

Round the pH to two decimal places because the [OH ∙ ] has two significant figures.

16. Calculate the pH of each solution. a. [OH−] = 4.3 × 10 −^5 M b. [OH−] = 4.5 × 10 −^11 M 17. Calculate the pH of each solution. a. [OH−] = 5.0 × 10 −^9 M b. [OH−] = 8.3 × 10 −^4 M

Calculating pH From [OH∙] What is the pH of a solution if [OH−] = 4.0 × 10 −^11 M?

➊ Analyze^ List the knowns and the unknown. To find [H+], divide Kw by the known [OH−]. Then, calculate pH as you did in Sample Problem 18.3.

➋ Calculate^ Solve for the unknown.

➌ Evaluate^ Does the result make sense?^ A solution in which [OH−] is less than 1^ ×^10 −^7 M is acidic because [H+] is greater than 1 × 10 −^7 M. The hydrogen-ion concentration is between 1 × 10 −^3 M and 1 × 10 −^4 M. Thus, the pH should be between 3 and 4.

Substitute the values for K (^) w and [OH∙] to find [H∙].

Use a calculator to find the log.

Start with the ion-product constant to find [H∙]. Rearrange the equation to solve for [H∙].

Next, use the equation for finding pH. Substitute the value for [H∙] that you just calculated.

Calculating pH From [OH∙] You can also use the hydroxide-ion concentration

to calculate the pH of a solution. Recall that the ion-product constant for water defines the relationship between [H+] and [OH−]. Therefore, you can use the ion-product constant for water to determine [H+] for a known [OH−]. Then, you use [H+] to calculate the pH. For practice, try doing Sample Problem 18.5.

TEKS 10I

InterpretGraphs

638 Chapter 18 • Lesson 2

Measuring pH What are two methods that are used to measure pH? In many situations, knowing the pH is useful. A custodian might need to maintain the correct acid-base balance in a swimming pool. A gardener may want to know if a certain plant will thrive in a yard. A doctor might be try- ing to diagnose a medical condition. Either acid-base indicators or pH meters can be used to measure pH.

Acid-Base Indicators An indicator is often used for initial pH measure-

ments and for samples with small volumes. An indicator (HIn) is an acid or a base that dissociates in a known pH range. Indicators work because their acid form and base form have different colors in solution. The following general equation represents the dissociation of an acid-base indicator (HIn).

The acid form of the indicator (HIn) is dominant at low pH and high [H+]. The base form (In−) is dominant at high pH and high [OH−]. The change from dominating acid form to dominating base form occurs within a narrow range of about two pH units. Within this range, the color of the solution is a mixture of the colors of the acid and the base forms. If you know the pH range over which this color change occurs, you can make a rough estimate of the pH of a solution. At all pH values below this range, you would see only the color of the acid form. At all pH values above this range, you would see only the color of the base form. For a more precise estimate of the solution’s pH, you could repeat the test with indicators that have different pH ranges for their color change. Many indicators are needed to span the entire pH spectrum. Figure 18.8 shows the pH ranges of some common acid-base indicators.

HIn(aq) H+(aq) + In−(aq) Acid form Base form

Figure 18.8 Each indicator is useful for a specific range of pH values. a. Identify At a pH of 12, which indicator would be yellow? b. Apply Concepts Which indicator could you use to show that the pH of a solution has changed from 3 to 5? c. Make Generalizations What do you notice about the range over which each indicator changes color?

Thymol blue Bromophenol blue Bromocresol green Methyl red Alizarin Bromothymol blue Phenol red Phenolphthalein Alizarin yellow R 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 pH

Color Ranges of Acid-Base Indicators

TEKS 10I

Quick Lab

640 Chapter 18 • Lesson 2

18. Review How are the concentrations of hydrogen ions and hydroxide ions related in an aqueous solution? 19. Identify What is the range of pH values in the following solutions? a. basic b. acidic c. neutral 20. List What methods can you use to measure the pH of a solution? 21. Understand pH What happens to the [H+] as the pH of a solution increases? 22. Calculate pH Use the hydrogen-ion or hydroxide- ion concentration to calculate the pH of each solution. a. [H+] = 1 × 10 −^6 M b. [H+] = 0.00010M c. [OH−] = 1 × 10 −^2 M d. [OH−] = 1 × 10 −^11 M 23. Compare In terms of ion concentrations, how do basic solutions differ from acidic solutions? 24. Calculate Find the hydroxide-ion concentra- tions for solutions with the following pH values: a. 6.00 b. 9.00 c. 12.

Analyze and Conclude

1. Observe What color is the indicator in acidic, neutral, and basic solutions? 2. Relate Cause and Effect What caused the color of the indicator to change when a material was added to a cup? 3. Classify Divide the household materials you tested into three groups—acidic, basic, and neutral. 4. Define pH In your own words, define pH.

Indicators From Natural Sources

5. Add several drops of vinegar to the first cup. Use a spatula to add a pinch of baking soda to the second cup. Add several drops of ammonia to the third cup. The pH values for the solutions of vin- egar, baking soda, and household ammonia are about 3, 9, and 11, respectively. Record the colors you observe at the correct locations on your pH scale. 6. Repeat the procedure for household items such as table salt, milk, lemon juice, laundry deter- gent, milk of magnesia, tooth- paste, shampoo, and carbonated beverages.

Procedure

1. Put one-half cup of finely chopped red cabbage leaves in a jar and add one-half cup of hot water. Stir and crush the leaves with a spoon. Continue this process until the water has a distinct color. 2. Strain the mixture through a piece of clean cheesecloth into a clean jar. The liquid that collects in the jar is your indicator. 3. Tape three sheets of paper end to end. Draw a line along the center of the taped sheets. Label the line at 5-cm intervals with the numbers 1 to 14. This labeled line is your pH scale. 4. Use the permanent marker to label three cups vinegar, baking soda, and ammonia. Pour indicator into each cup to a depth of about 1 cm.

Purpose To measure the pH of household materials using a natural indicator

Materials

**- red cabbage leaves

• 1-cup measure
• hot water
• 2 jars
• spoon
• cheesecloth
• 3 sheets of plain white paper
• transparent tape
• metric ruler
• pencil
• 10 small clear plastic cups
• permanent marker
• white vinegar (CH 3 COOH)
• baking soda (NaHCO 3 )
• spatula
• household ammonia
• dropper
• assorted household materials**

18.2 LessonCheck TEKS 10I

TEKS 10I

Take It Further

1. Apply Concepts The ideal soil pH

for corn is around 6.0. If [H+] equals

2.14 × 10 −^5 M, is the soil too acidic or too basic for growing corn?

2. Describe Connections Do Internet research on the agronomy program at a Texas college or university. Find out what chemistry concepts are studied by students in the program. Prepare an oral report to describe the connection between chemistry and a future career as an agronomist.

FIELD CHEMISTRY Agronomists can use science to help design and maintain golf courses such as the course at Barton Creek in Austin, Texas, which is pictured above.

Agronomist

Do you like the idea of working with plants but find

chemistry more interesting than biology? If so, you

might want to consider a career in agronomy. There are

agronomy programs at universities in Texas, such as

Texas A&M and Texas Tech, where you can learn about

the connection between chemistry and a future career

in agronomy. Agronomy is a branch of agriculture that

deals with the interactions between plants, soils, and the

environment. Agronomists can use their knowledge of

chemistry to help produce healthy crops and increase

yields, while preserving the environment.

The opportunities for agronomists extend beyond

laboratories and greenhouses. Agronomists also work

for business firms, government agencies, conservation

groups, philanthropic organizations, and universities.

Agronomists can also use their knowledge to address

such issues as urban area beautification and golf

course landscaping.

Chemistry & You 641

SOIL ACIDITY The pH of the soil is among the most important factors for growing plants. Agronomists can advise farmers on the right soil pH for a specific crop.

CHEMISTRY (^) & YOU: CAreers TEKS 3E