Tải bản đầy đủ - 0 (trang)
5: Reactions of Metals with Nonmetals (Oxidation–Reduction)

5: Reactions of Metals with Nonmetals (Oxidation–Reduction)

Tải bản đầy đủ - 0trang

C H E M I S T R Y I N F OCUS

Garden-Variety Acid–Base Indicators



When placed in an acidic solution, most of

the basic form of the indicator is converted to

the acidic form by the reaction



What



can flowers tell us about acids and

bases? Actually, some flowers can tell us whether

the soil they are growing in is acidic or basic. For

example, in acidic soil, bigleaf hydrangea blossoms will be blue; in basic (alkaline) soil, the

flowers will be red. What is the secret? The pigment in the flower is an acid–base indicator.

Generally, acid–base indicators are dyes that

are weak acids. Because indicators are usually

complex molecules, we often symbolize them as

HIn. The reaction of the indicator with water can

be written as

HIn(aq) ϩ H2O(l) 3

4 H3Oϩ(aq) ϩ InϪ(aq)

To work as an acid–base indicator, the conjugate acid–base forms of these dyes must have different colors. The acidity level of the solution will

determine whether the indicator is present mainly

in its acidic form (HIn) or its basic form (InϪ).



InϪ(aq) ϩ Hϩ(aq) n HIn(aq)

When placed in a basic solution, most of the

acidic form of the indicator is converted to the

basic form by the reaction

HIn(aq) ϩ OHϪ(aq) n InϪ(aq) ϩ H2O(l)

It turns out that many fruits, vegetables,

and flowers can act as acid–base indicators. Red,

blue, and purple plants often contain a class of

chemicals called anthocyanins, which change

color based on the acidity level of the surroundings. Perhaps the most famous of these plants is

red cabbage. Red cabbage contains a mixture of

anthocyanins and other pigments that allow it

to be used as a “universal indicator.” Red cabbage juice appears deep red at a pH of 1–2, purple at a pH of 4, blue at a pH of 8, and green at

a pH of 11.



16.5 Calculating the pH of Strong

Acid Solutions

OBJECTIVE:



To learn to calculate the pH of solutions of strong acids.

In this section we will learn to calculate the pH for a solution containing a

strong acid of known concentration. For example, if we know a solution contains 1.0 M HCl, how can we find the pH of the solution? To answer this

question we must know that when HCl dissolves in water, each molecule dissociates (ionizes) into Hϩ and ClϪ ions. That is, we must know that HCl is a

strong acid. Thus, although the label on the bottle says 1.0 M HCl, the solution contains virtually no HCl molecules. A 1.0 M HCl solution contains Hϩ

and ClϪ ions rather than HCl molecules. Typically, container labels indicate

the substance(s) used to make up the solution but do not necessarily describe

the solution components after dissolution. In this case,

1.0 M HCl S 1.0 M Hϩ and 1.0 M ClϪ

Therefore, the [Hϩ] in the solution is 1.0 M. The pH is then

pH ϭ Ϫlog[Hϩ] ϭ Ϫlog(1.0) ϭ 0



532



variety of flower petals, including delphiniums,

geraniums, morning glories, and, of course, hydrangeas.



Royalty-free Corbis



Neil Holmes/Homes Garden Photos/Alamy



Other natural indicators include the skins

of beets (which change from red to purple in

very basic solutions), blueberries (which change

from blue to red in acidic solutions), and a wide



EXAMPLE 16.10



Calculating the pH of Strong Acid Solutions

Calculate the pH of 0.10 M HNO3.



SOLUTION

HNO3 is a strong acid, so the ions in solution are Hϩ and NO3Ϫ. In this case,

0.10 M HNO3 S 0.10 M Hϩ and 0.10 M NOϪ

3

Thus

[Hϩ] ϭ 0.10 M



and



pH ϭ Ϫlog(0.10) ϭ 1.00



Self-Check EXERCISE 16.7 Calculate the pH of a solution of 5.0 ϫ 10Ϫ3 M HCl.

See Problems 16.57 and 16.58. ■



533



534 Chapter 16 Acids and Bases



16.6 Buffered Solutions

OBJECTIVE:



Water: pH ϭ 7

0.01 M HCl: pH ϭ 2



To understand the general characteristics of buffered solutions.

A buffered solution is one that resists a change in its pH even when a

strong acid or base is added to it. For example, when 0.01 mole of HCl is

added to 1 L of pure water, the pH changes from its initial value of 7 to 2, a

change of 5 pH units. However, when 0.01 mole of HCl is added to a solution containing both 0.1 M acetic acid (HC2H3O2) and 0.1 M sodium acetate

(NaC2H3O2), the pH changes from an initial value of 4.74 to 4.66, a change

of only 0.08 pH unit. The latter solution is buffered—it undergoes only a

very slight change in pH when a strong acid or base is added to it.

Buffered solutions are vitally important to living organisms whose cells

can survive only in a very narrow pH range. Many goldfish have died because their owners did not realize the importance of buffering the aquarium

water at an appropriate pH. For humans to survive, the pH of the blood must

be maintained between 7.35 and 7.45. This narrow range is maintained by

several different buffering systems.

A solution is buffered by the presence of a weak acid and its conjugate

base. An example of a buffered solution is an aqueous solution that contains

acetic acid and sodium acetate. The sodium acetate is a salt that furnishes acetate ions (the conjugate base of acetic acid) when it dissolves. To see how

this system acts as a buffer, we must recognize that the species present in this

solution are

Naϩ,



HC2H3O2,



C2H3O2–



Hans Reinhard/Bruce Coleman, Inc.



When NaC2H3O2 is

dissolved, it produces

the separated ions



What happens in this solution when a strong acid such as HCl is added?

In pure water, the Hϩ ions from the HCl would accumulate, thus lowering

the pH.

100%



HCl ¡ Hϩ ϩ ClϪ

For goldfish to survive, the pH

of the water must be carefully

controlled.



However, this buffered solution contains C2H3O2Ϫ ions, which are basic.

That is, C2H3O2Ϫ has a strong affinity for Hϩ, as evidenced by the fact that

HC2H3O2 is a weak acid. This means that the C2H3O2Ϫ and Hϩ ions do not

exist together in large numbers. Because the C2H3O2Ϫ ion has a high affinity

for Hϩ, these two combine to form HC2H3O2 molecules. Thus the Hϩ from

the added HCl does not accumulate in solution but reacts with the C2H3O2Ϫ

as follows:

H ϩ (aq) ϩ C2H3OϪ

2 (aq) S HC2H3O2(aq)

Next consider what happens when a strong base such as sodium hydroxide is added to the buffered solution. If this base were added to pure water, the OHϪ ions from the solid would accumulate and greatly change (raise)

the pH.

100%



NaOH ¡ Naϩ ϩ OHϪ

However, in the buffered solution the OHϪ ion, which has a very strong affinity for Hϩ, reacts with HC2H3O2 molecules as follows:

HC2H3O2(aq) ϩ OHϪ(aq) S H2O(l) ϩ C2H3OϪ

2 (aq)



Chapter Review



535



Table 16.3 The Characteristics of a Buffer

1. The solution contains a weak acid HA and its conjugate base AϪ.

2. The buffer resists changes in pH by reacting with any added Hϩ or OHϪ so

that these ions do not accumulate.

3. Any added Hϩ reacts with the base AϪ.

Hϩ(aq) ϩ AϪ(aq) n HA(aq)

4. Any added OHϪ reacts with the weak acid HA.

OHϪ(aq) ϩ HA(aq) n H2O(l) ϩ AϪ(aq)



This happens because, although C2H3O2Ϫ has a strong affinity for Hϩ, OHϪ

has a much stronger affinity for Hϩ and thus can remove Hϩ ions from acetic

acid molecules.

Note that the buffering materials dissolved in the solution prevent

added Hϩ or OHϪ from building up in the solution. Any added Hϩ is trapped

by C2H3O2Ϫ to form HC2H3O2. Any added OHϪ reacts with HC2H3O2 to form

H2O and C2H3O2Ϫ.

The general properties of a buffered solution are summarized in

Table 16.3.



C H A P T E R



16



REVIEW



F



Key Terms

acid (16.1)

base (16.1)

Arrhenius concept of

acids and bases (16.1)

Brønsted–Lowry

model (16.1)

conjugate acid (16.1)

conjugate base (16.1)

conjugate acid–base

pair (16.1)

hydronium ion (16.1)

completely ionized

(dissociated) (16.2)

strong acid (16.2)

weak acid (16.2)



diprotic acid (16.2)

oxyacid (16.2)

organic acid (16.2)

carboxyl group (16.2)

amphoteric

substance (16.3)

ionization of water (16.3)

ion-product constant,

Kw (16.3)

neutral solution (16.3)

acidic solution (16.3)

basic solution (16.3)

pH scale (16.4)

buffered solution (16.6)

buffered (16.6)



Summary

1. Acids or bases in water are commonly described by

two different models. Arrhenius postulated that acids

produce Hϩ ions in aqueous solutions and that bases

produce OHϪ ions. The Brønsted–Lowry model is



VP



directs you to the Chemistry in Focus feature in the chapter

indicates visual problems

interactive versions of these problems are assignable in OWL



more general: an acid is a proton donor, and a base is

a proton acceptor. Water acts as a Brønsted–Lowry

base when it accepts a proton from an acid to form a

hydronium ion:

HA(aq) ϩ H2O(l) 3

4 H3Oϩ(aq) ϩ AϪ(aq)

Acid



Base



Conjugate

acid



Conjugate

base



A conjugate base is everything that remains of the

acid molecule after the proton is lost. A conjugate

acid is formed when a proton is transferred to the

base. Two substances related in this way are called a

conjugate acid–base pair.

2. A strong acid or base is one that is completely ionized

(dissociated). A weak acid is one that is ionized (dissociated) only to a slight extent. Strong acids have

weak conjugate bases. Weak acids have relatively

strong conjugate bases.

3. Water is an amphoteric substance—it can behave either as an acid or as a base. The ionization of water reveals this property; one water molecule transfers a



536 Chapter 16 Acids and Bases

proton to another water molecule to produce a hydronium ion and a hydroxide ion.

H2O(l) ϩ H2O(l) 3

4 H3Oϩ(aq) ϩ OHϪ(aq)

The expression

ϩ



Ϫ



ϩ



Ϫ



Kw ϭ [H3O ][OH ] ϭ [H ][OH ]

is called the ion-product constant. It has been shown

experimentally that at 25 °C,

[Hϩ] ϭ [OHϪ] ϭ 1.0 ϫ 10Ϫ7 M

so Kw ϭ 1.0 ϫ 10Ϫ14.

4. In an acidic solution, [Hϩ] is greater than [OHϪ]. In a

basic solution, [OHϪ] is greater than [Hϩ]. In a neutral

solution, [Hϩ] ϭ [OHϪ].

ϩ



5. To describe [H ] in aqueous solutions, we use the pH

scale.

pH ϭ Ϫlog[Hϩ]

Note that the pH decreases as [Hϩ] (acidity) increases.

6. The pH of strong acid solutions can be calculated directly from the concentration of the acid, because

100% dissociation occurs in aqueous solution.

7. A buffered solution is one that resists a change in its

pH even when a strong acid or base is added to it. A

buffered solution contains a weak acid and its conjugate base.



Active Learning Questions

These questions are designed to be considered by groups of

students in class. Often these questions work well for introducing a particular topic in class.

1. You are asked for the Hϩ concentration in a solution

of NaOH(aq). Because sodium hydroxide is a strong

base, can we say there is no Hϩ, since having Hϩ

would imply that the solution is acidic?

2. Explain why ClϪ does not affect the pH of an aqueous

solution.



8. Can the pH of a solution be negative? Explain.

9. Stanley’s grade-point average (GPA) is 3.28. What is

Stanley’s p(GPA)?

10. A friend asks the following: “Consider a buffered solution made up of the weak acid HA and its salt NaA.

If a strong base like NaOH is added, the HA reacts

with the OHϪ to make AϪ. Thus, the amount of acid

(HA) is decreased, and the amount of base (AϪ) is increased. Analogously, adding HCl to the buffered solution forms more of the acid (HA) by reacting with

the base (AϪ). How can we claim that a buffered solution resists changes in the pH of the solution?” How

would you explain buffering to your friend?

11. Mixing together aqueous solutions of acetic acid and

sodium hydroxide can make a buffered solution.

Explain.

12. Could a buffered solution be made by mixing aqueous solutions of HCl and NaOH? Explain.

13. Consider the equation: HA(aq) ϩ H2O 3

4 H3Oϩ(aq) ϩ

Ϫ

A (aq).

a. If water is a better base than AϪ, which way will

equilibrium lie?

b. If water is a better base than AϪ, does this mean

that HA is a strong or a weak acid?

c. If water is a better base than AϪ, is the value for Ka

greater or less than 1?

14. Choose the answer that best completes the following

statement and defend your answer. When 100.0 mL

of water is added to 100.0 mL of 1.00 M HCl,

a. the pH decreases because the solution is diluted.

b. the pH does not change because water is neutral.

c. the pH is doubled because the volume is now

doubled.

d. the pH increases because the concentration of Hϩ

decreases.

e. the solution is completely neutralized.

15. You mix a solution of a strong acid with a pH of 4 and

an equal volume of a strong acid solution with a pH

of 6. Is the final pH less than 4, between 4 and 5, 5,

between 5 and 6, or greater than 6? Explain.



3. Write the general reaction for an acid acting in water.

What is the base in this case? The conjugate acid? The VP 16. The following figures are molecular-level representaconjugate base?

tions of acid solutions. Label each as a strong acid or

a weak acid.

4. Differentiate among the terms concentrated, dilute,

weak, and strong in describing acids. Use molecularlevel pictures to support your answer.

5. What is meant by “pH”? True or false: A strong acid

always has a lower pH than a weak acid does. Explain.

6. Consider two separate solutions: one containing a

weak acid, HA, and one containing HCl. Assume that

you start with 10 molecules of each.

a. Draw a molecular-level picture of what each solution looks like.

b. Arrange the following from strongest to weakest

base: ClϪ, H2O, AϪ. Explain.

7. Why is the pH of water at 25 °C equal to 7.00?



H+

B–



H+

A–



17. Answer the following questions concerning buffered

solutions.

a. Explain what a buffered solution does.

b. Describe the substances that make up a buffered

solution.

c. Explain how a buffered solution works.



Tài liệu bạn tìm kiếm đã sẵn sàng tải về

5: Reactions of Metals with Nonmetals (Oxidation–Reduction)

Tải bản đầy đủ ngay(0 tr)

×