Tải bản đầy đủ - 0 (trang)
1 Brønsted Acids and Bases

1 Brønsted Acids and Bases

Tải bản đầy đủ - 0trang

SEC T I ON 16. 2



The Acid-Base Properties of Water



675





Practice Problem A What is (a) the conjugate acid of ClOϪ

4 , (b) the conjugate acid of S , (c) the



conjugate base of H2S, and (d) the conjugate base of H2C2O4?

Ϫ

Practice Problem B HSOϪ

3 is the conjugate acid of what species? HSO 3 is the conjugate base of



what species?



Sample Problem 16.2

Label each of the species in the following equations as an acid, base, conjugate base, or conjugate

acid:

FϪ(aq) ϩ NHϩ4 (aq)



(a) HF(aq) ϩ NH3(aq)

Ϫ



(b) CH3COO (aq) ϩ H2O(l)



CH3COOH(aq) ϩ OHϪ(aq)



Strategy In each equation, the reactant that loses a proton is the acid and the reactant that gains a



proton is the base. Each product is the conjugate of one of the reactants. Two species that differ only

by a proton constitute a conjugate pair.

Setup (a) HF loses a proton and becomes FϪ; NH3 gains a proton and becomes NHϩ

4.



(b) CH3COOϪ gains a proton to become CH3COOH; H2O loses a proton to become OHϪ.

Solution



FϪ(aq) ϩ NHϩ4 (aq)



(a) HF(aq) ϩ NH3(aq)

acid



base



conjugate

base



(b) CH3COOϪ(aq) ϩ H2O(l)

base



acid



conjugate

acid



CH3COOH(aq) ϩ OHϪ(aq)

conjugate acid



conjugate base



Think About It In a Brønsted



acid-base reaction, there is always

an acid and a base, and whether

a substance behaves as an acid

or a base depends on what it

is combined with. Water, for

example, behaves as a base when

combined with HCl but behaves as

an acid when combined with NH3.



Practice Problem A Identify and label the species in each reaction.



(a) NHϩ4 (aq) ϩ H2O(l)

Ϫ



(b) CN (aq) ϩ H2O(l)



NH3(aq) ϩ H3Oϩ(aq)



Apago PDF Enhancer



HCN(aq) ϩ OHϪ(aq)



Practice Problem B (a) Write an equation in which HSOϪ

4 reacts (with water) to form its conjugate



base. (b) Write an equation in which HSOϪ4 reacts (with water) to form its conjugate acid.



Checkpoint 16.1



Brønsted Acids and Bases



16.1.1 Which of the following pairs of species

are conjugate pairs? (Select all that

apply.)



16.1.2 Which of the following species does not

have a conjugate base? (Select all that

apply.)



a) H2S and S2Ϫ



a) HC2OϪ4



b) NHϪ2 and NH3



b) OHϪ



c) O2 and H2O2



c) O2Ϫ



Ϫ



d) HBr and Br



d) CO2Ϫ

3



e) HCl and OHϪ



e) HClO



16.2 The Acid-Base Properties of Water

Water is often referred to as the “universal solvent,” because it is so common and so important to

life on Earth. In addition, most of the acid-base chemistry that you will encounter takes place in

aqueous solution. In this section, we take a closer look at water’s ability to act as either a Brønsted

acid (as in the ionization of NH3) or a Brønsted base (as in the ionization of HCl). A species that

can behave either as a Brønsted acid or a Brønsted base is called amphoteric.

Water is a very weak electrolyte, but it does undergo ionization to a small extent:

H2O(l)



bur75640_ch16_672-725.indd 675



Hϩ(aq) ϩ OHϪ(aq)



11/19/09 4:01:58 PM



676



C H A P T E R 16



Acids and Bases



This reaction is known as the autoionization of water. Because we can represent the aqueous proton as either Hϩ or H3Oϩ [9 Section 4.3], we can also write the autoionization of water as

H3Oϩ(aq)



2H2O(l)



ϩ



ϩ

Multimedia



acid



Chemical Equilibrium—equilibrium

(interactive).

Student Annotation: Recall that in a

heterogeneous equilibrium such as this, liquids

and solids do not appear in the equilibrium

expression [9 Section 15.3] .



Student Annotation: The constant Kw is

sometimes referred to as the ion-product

constant.



Student Annotation: Recall that we

disregard the units when we substitute

concentrations into an equilibrium expression

[9 Section 15.5] .



base



Ϫ



ϩ



conjugate

acid



conjugate

base



where one water molecule acts as an acid and the other acts as a base.

As indicated by the double arrow in the equation, the reaction is an equilibrium. The equilibrium expression for the autoionization of water is

Kw ϭ [H3Oϩ][OHϪ]



or



Kw ϭ [Hϩ][OHϪ]



Because the autoionization of water is an important equilibrium that you will encounter frequently

in the study of acids and bases, we use the subscript “w” to indicate that the equilibrium constant

is that specifically for the autoionization of water. It is important to realize, though, that Kw is

simply a Kc for a specific reaction. We will frequently replace the c in Kc expressions with a letter

or a series of letters to indicate the specific type of reaction to which the Kc refers. For example, Kc

for the ionization of a weak acid is called Ka, and Kc for the ionization of a weak base is called Kb.

In Chapter 17 we will encounter Ksp, where “sp” stands for “solubility product.” Each specially

subscripted K is simply a Kc for a specific type of reaction.

In pure water, autoionization is the only source of H3Oϩ and OHϪ, and the stoichiometry of

the reaction tells us that their concentrations are equal. At 25°C, the concentrations of hydronium

and hydroxide ions in pure water are [H3Oϩ] ϭ [OHϪ] ϭ 1.0 ϫ 10Ϫ7 M. Using the equilibrium

expression, we can calculate the value of Kw at 25°C as follows:

Kw ϭ [H3Oϩ][OHϪ] ϭ (1.0 ϫ 10Ϫ7)(1.0 ϫ 10Ϫ7) ϭ 1.0 ϫ 10Ϫ14

Furthermore, in any aqueous solution at 25°C, the product of H3Oϩ and OHϪ concentrations is

equal to 1.0 ϫ 10Ϫ14.



Apago KPDF

ϭ [H OEnhancer

][OH ] ϭ 1.0 ϫ 10



Equation 16.1



Student Annotation: Because the product of

H3Oϩ and OHϪ concentrations is a constant, we

cannot alter the concentrations independently.

Any change in one also affects the other.



ϩ OHϪ(aq)



w



3



ϩ



Ϫ



Ϫ14



(at 25°C)



Although their product is a constant, the individual concentrations of hydronium and hydroxide

can be influenced by the addition of an acid or a base. The relative amounts of H3Oϩ and OHϪ

determine whether a solution is neutral, acidic, or basic.

• When [H3Oϩ] ϭ [OHϪ], the solution is neutral.

• When [H3Oϩ] Ͼ [OHϪ], the solution is acidic.

• When [H3Oϩ] Ͻ [OHϪ], the solution is basic.

Sample Problem 16.3 shows how to use Equation 16.1.



Sample Problem 16.3

The concentration of hydronium ions in stomach acid is 0.10 M. Calculate the concentration of

hydroxide ions in stomach acid at 25°C.

Strategy Use the value of Kw to determine [OHϪ] when [H3Oϩ] ϭ 0.10 M.

Setup Kw ϭ [H3Oϩ][OHϪ] ϭ 1.0 ϫ 10Ϫ14 at 25°C. Rearranging Equation 16.1 to solve for [OHϪ],

Think About It Remember

that equilibrium constants are

temperature dependent. The

value of Kw is 1.0 ϫ 10Ϫ14 only

at 25°C.



1.0 ϫ 10Ϫ14

[OHϪ] ϭ __________

[H3Oϩ]

Solution



1.0 ϫ 10Ϫ14

[OHϪ] ϭ __________ ϭ 1.0 ϫ 10Ϫ13 M

0.10



Practice Problem A The concentration of hydroxide ions in the antacid milk of magnesia is

5.0 ϫ 10Ϫ4 M. Calculate the concentration of hydronium ions at 25°C.

Practice Problem B The value of Kw at normal body temperature (37°C) is 2.8 ϫ 10Ϫ14. Calculate



the concentration of hydroxide ions in stomach acid at body temperature. ([H3Oϩ] ϭ 0.10 M.)



bur75640_ch16_672-725.indd 676



11/25/09 1:06:44 PM



SEC T I ON 16. 3



Checkpoint 16.2



16.2.2 Calculate [H3Oϩ] in a solution in which

[OHϪ] ϭ 0.25 M at 25°C.



a) 1.2 ϫ 10Ϫ3 M



a) 4.0 ϫ 10Ϫ14 M



b) 8.3 ϫ 10Ϫ17 M



b) 1.0 ϫ 10Ϫ14 M



c) 1.0 ϫ 10



677



The Acid-Base Properties of Water



16.2.1 Calculate [OHϪ] in a solution in which

[H3Oϩ] ϭ 0.0012 M at 25°C.



Ϫ14



The pH Scale



M



c) 2.5 ϫ 1013 M



d) 8.3 ϫ 10Ϫ12 M



d) 1.0 ϫ 10Ϫ7 M



e) 1.2 ϫ 1011 M



e) 4.0 ϫ 10Ϫ7 M



16.3 The pH Scale

The acidity of an aqueous solution depends on the concentration of hydronium ions, [H3Oϩ]. This

concentration can range over many orders of magnitude, which can make reporting the numbers

cumbersome. To describe the acidity of a solution, rather than report the molar concentration of

hydronium ions, we typically use the more convenient pH scale. The pH of a solution is defined as

the negative base-10 logarithm of the hydronium ion concentration (in mol/L).

pH ϭ Ϫlog[H3Oϩ]



or



pH ϭ Ϫlog[Hϩ]



Equation 16.2



Student Annotation: Equation 16.2 converts

numbers that can span an enormous range

(ϳ101 to 10Ϫ14) to numbers generally ranging

from ϳ1 to 14.



The pH of a solution is a dimensionless quantity, so the units of concentration must be removed

from [H3Oϩ] before taking the logarithm. Because [H3Oϩ] ϭ [OHϪ] ϭ 1.0 ϫ 10Ϫ7 M in pure

water at 25°C, the pH of pure water at 25°C is

Ϫlog (1.0 ϫ 10Ϫ7) ϭ 7.00

ϩ



Ϫ



Remember, too, that a solution in which [H3O ] ϭ [OH ] is neutral. At 25°C, therefore, a neutral

solution has pH 7.00. An acidic solution, one in which [H3Oϩ] Ͼ [OHϪ], has pH Ͻ 7.00, whereas

a basic solution, in which [H3Oϩ] Ͻ [OHϪ], has pH Ͼ 7.00. Table 16.3 shows the calculation of

pH for solutions ranging from 0.10 M to 1.0 ϫ 10Ϫ14 M.

In the laboratory, pH is measured with a pH meter (Figure 16.1). Table 16.4 lists the pH values of a number of common fluids. Note that the pH of body fluids varies greatly, depending on the

location and function of the fluid. The low pH (high acidity) of gastric juices is vital for digestion

of food, whereas the higher pH of blood is required to facilitate the transport of oxygen.



Apago PDF Enhancer



TABLE 16.3

[H 3 O ؉ ](M)



Benchmark pH Values for a Range of Hydronium Ion

Concentrations at 25°C

؊log [H 3 O ؉ ]



pH



Ϫ1



Ϫlog (1.0 ϫ 10 )



0.10



Ϫ2



Ϫlog (1.0 ϫ 10 )



0.010

Ϫ3



1.0 ϫ 10



Ϫ4



1.0 ϫ 10



Ϫ5



1.0 ϫ 10



Ϫ6



1.0 ϫ 10



Ϫ7



1.0 ϫ 10



Ϫ8



1.0 ϫ 10



Ϫ9



1.0 ϫ 10



Ϫ10



1.0 ϫ 10



Ϫ11



Ϫ3



Ϫlog (1.0 ϫ 10 )

Ϫ4



Ϫlog (1.0 ϫ 10 )

Ϫ5



Ϫlog (1.0 ϫ 10 )

Ϫ6



Ϫlog (1.0 ϫ 10 )

Ϫ7



Ϫlog (1.0 ϫ 10 )

Ϫ8



Ϫlog (1.0 ϫ 10 )

Ϫ9



Ϫlog (1.0 ϫ 10 )

Ϫ10



Ϫlog (1.0 ϫ 10



)



Ϫ11



1.00

2.00

3.00

4.00

5.00

6.00



Acidic



7.00



Neutral



8.00



Basic



9.00

10.00



1.0 ϫ 10



Ϫlog (1.0 ϫ 10



)



11.00



1.0 ϫ 10Ϫ12



Ϫlog (1.0 ϫ 10Ϫ12)



12.00



Ϫ13



1.0 ϫ 10



Ϫ14



1.0 ϫ 10



bur75640_ch16_672-725.indd 677



Student Annotation: A word about

significant figures: When we take the log

of a number with two significant figures,

we report the result to two places past

the decimal point. Thus, pH 7.00 has two

significant figures, not three.



Ϫ13



Ϫlog (1.0 ϫ 10



)



Ϫlog (1.0 ϫ 10



Ϫ14



)



13.00

14.00



Figure 16.1 A pH meter is

commonly used in the laboratory

to determine the pH of a solution.

Although many pH meters have a range

of 1 to 14, pH values can actually be

less than 1 and greater than 14.



11/19/09 4:01:59 PM



678



C H A P T E R 16



Acids and Bases



TABLE 16.4



pH Values of Some Common Fluids



Fluid



pH



Fluid



pH



Stomach acid



1.5



Saliva



6.4Ϫ6.9



Lemon juice



2.0



Milk



6.5



Vinegar



3.0



Pure water



7.0



Grapefruit juice



3.2



Blood



7.35Ϫ7.45



Orange juice



3.5



Tears



7.4



Urine



4.8Ϫ7.5



Milk of magnesia



10.6



Rainwater (in clean air)



5.5



Household ammonia



11.5



A measured pH can be used to determine experimentally the concentration of hydronium ion

in solution. Solving Equation 16.2 for [H3Oϩ] gives



Student Annotation: 10x is the inverse

function of log. (It is usually the second

function on the same key.) You must be

comfortable performing these operations on

your calculator.



[H3Oϩ] ϭ 10ϪpH



Equation 16.3



Sample Problems 16.4 and 16.5 illustrate calculations involving pH.



Sample Problem 16.4

Think About It When a hydronium ion

concentration falls between two “benchmark”

concentrations in Table 16.3, the pH falls between

the two corresponding pH values. In part (c),

for example, the hydronium ion concentration

(8.8 ϫ 10Ϫ11 M) is greater than 1.0 ϫ 10Ϫ11 M

but less than 1.0 ϫ 10Ϫ10 M. Therefore, we expect

the pH to be between 11.00 and 10.00.



[H3O؉] (M)

Ϫ10



1.0 ϫ 10

8.8 ϫ 10Ϫ11*

1.0 ϫ 10Ϫ11



؊log [H3O؉]

Ϫ10



Ϫlog (1.0 ϫ 10 )

Ϫlog (8.8 ϫ 10Ϫ11)

Ϫlog (1.0 ϫ 10Ϫ11)



Determine the pH of a solution at 25°C in which the hydronium ion concentration is

(a) 3.5 ϫ 10Ϫ4 M, (b) 1.7 ϫ 10Ϫ7 M, and (c) 8.8 ϫ 10Ϫ11 M.

Strategy Given [H3Oϩ], use Equation 16.2 to solve for pH.

Setup



(a) pH ϭ Ϫlog (3.5 ϫ 10Ϫ4)



Apago

PDF

)

(b) pH ϭ Ϫlog

(1.7 ϫ 10Enhancer

Ϫ7



(c) pH ϭ Ϫlog (8.8 ϫ 10



pH



Solution



10.00

10.06†

11.00



(b) pH ϭ 6.77



Ϫ11



)



(a) pH ϭ 3.46

(c) pH ϭ 10.06



*[H3Oϩ] between two benchmark values

†pH between two benchmark values



Recognizing the benchmark concentrations

and corresponding pH values is a good way to

determine whether or not your calculated result

is reasonable.



Practice Problem A Determine the pH of a solution at 25°C in which the hydronium ion

concentration is (a) 3.2 ϫ 10Ϫ9 M, (b) 4.0 ϫ 10Ϫ8 M, and (c) 5.6 ϫ 10Ϫ2 M.

Practice Problem B Determine the pH of a solution at 25°C in which the hydroxide ion

concentration is (a) 8.3 ϫ 10Ϫ15 M, (b) 3.3 ϫ 10Ϫ4 M, and (c) 1.2 ϫ 10Ϫ3 M.



Sample Problem 16.5

Calculate the hydronium ion concentration in a solution at 25°C in which the pH is (a) 4.76,

(b) 11.95, and (c) 8.01.

Strategy Given pH, use Equation 16.3 to calculate [H3Oϩ].

Setup



(a) [H3Oϩ] ϭ 10Ϫ4.76

(b) [H3Oϩ] ϭ 10Ϫ11.95

(c) [H3Oϩ] ϭ 10Ϫ8.01



bur75640_ch16_672-725.indd 678



11/25/09 1:07:27 PM



SEC T I ON 16. 3



Solution



(a) [H3Oϩ] ϭ 1.7 ϫ 10Ϫ5 M

(b) [H3Oϩ] ϭ 1.1 ϫ 10Ϫ12 M

(c) [H3Oϩ] ϭ 9.8 ϫ 10Ϫ9 M



Practice Problem A Calculate the hydronium ion concentration in a solution at 25°C in which the

pH is (a) 9.90, (b) 1.45, and (c) 7.01.

Practice Problem B Calculate the hydroxide ion concentration in a solution at 25°C in which the pH

is (a) 11.89, (b) 2.41, and (c) 7.13.



A pOH scale analogous to the pH scale can be defined using the negative base-10 logarithm

of the hydroxide ion concentration of a solution, [OHϪ].

pOH ϭ Ϫlog [OHϪ]



Equation 16.4



The pH Scale



679



Think About It If you use

the calculated hydronium ion

concentrations to recalculate pH,

you will get numbers slightly

different from those given in the

problem. In part (a), for example,

Ϫlog (1.7 ϫ 10Ϫ5) ϭ 4.77. The

small difference between this

and 4.76 (the pH given in the

problem) is due to a rounding error.

Remember that a concentration

derived from a pH with two digits

to the right of the decimal point can

have only two significant figures.

Note also that the benchmarks

can be used equally well in this

circumstance. A pH between 4 and

5 corresponds to a hydronium ion

concentration between 1 ϫ 10Ϫ4 M

and 1 ϫ 10Ϫ5 M.



Rearranging Equation 16.4 to solve for hydroxide ion concentration gives

[OHϪ] ϭ 10ϪpOH



Equation 16.5



Now consider again the Kw equilibrium expression for water at 25°C:

[H3Oϩ][OHϪ] ϭ 1.0 ϫ 10Ϫ14

Taking the negative logarithm of both sides, we obtain

Ϫlog ([H3Oϩ][OHϪ]) ϭ Ϫlog (1.0 ϫ 10Ϫ14)



Apago PDF Enhancer



Ϫ(log [H3Oϩ] ϩ log [OHϪ]) ϭ 14.00



Ϫlog [H3Oϩ] Ϫ log [OHϪ] ϭ 14.00

(Ϫlog [H3Oϩ]) ϩ (Ϫlog [OHϪ]) ϭ 14.00

And from the definitions of pH and pOH, we see that at 25°C

pH ϩ pOH ϭ 14.00



Equation 16.6



Equation 16.6 provides another way to express the relationship between the hydronium ion concentration and the hydroxide ion concentration. On the pOH scale, 7.00 is neutral, numbers greater

than 7.00 indicate that a solution is acidic, and numbers less than 7.00 indicate that a solution is

basic. Table 16.5 lists pOH values for a range of hydroxide ion concentrations at 25°C.



Benchmark pOH Values for a Range of Hydroxide Ion

Concentrations at 25°C



TABLE 16.5



[OH ؊ ] (M)



0.10



1.00

Ϫ3



1.0 ϫ 10



Ϫ5



1.0 ϫ 10



Ϫ7



1.0 ϫ 10



Ϫ9



1.0 ϫ 10



Ϫ11



1.0 ϫ 10



Ϫ13



1.0 ϫ 10



bur75640_ch16_672-725.indd 679



pOH



3.00

5.00



Basic



7.00



Neutral



9.00



Acidic



11.00

13.00



11/25/09 1:08:29 PM



680



C H A P T E R 16



Acids and Bases



Sample Problems 16.6 and 16.7 illustrate calculations involving pOH.



Sample Problem 16.6

Think About It Remember that

the pOH scale is, in essence, the

reverse of the pH scale. On the

pOH scale, numbers below 7

indicate a basic solution, whereas

numbers above 7 indicate an

acidic solution. The pOH

benchmarks (abbreviated in

Table 16.5) work the same way

the pH benchmarks do. In part (a),

for example, a hydroxide ion

concentration between 1 ϫ 10Ϫ4 M

and 1 ϫ 10Ϫ5 M corresponds to a

pOH between 4 and 5:



[OH؊] (M)



pOH



1.0 ϫ 10Ϫ4

3.7 ϫ 10Ϫ5*

1.0 ϫ 10Ϫ5



4.00

4.43†

5.00



*[OHϪ] between two benchmark values

†pOH between two benchmark values



Determine the pOH of a solution at 25°C in which the hydroxide ion concentration is (a) 3.7 ϫ 10Ϫ5 M,

(b) 4.1 ϫ 10Ϫ7 M, and (c) 8.3 ϫ 10Ϫ2 M.

Strategy Given [OHϪ], use Equation 16.4 to calculate pOH.

Setup



(a) pOH ϭ Ϫlog (3.7 ϫ 10Ϫ5)

(b) pOH ϭ Ϫlog (4.1 ϫ 10Ϫ7)

(c) pOH ϭ Ϫlog (8.3 ϫ 10Ϫ2)

Solution



(a) pOH ϭ 4.43

(b) pOH ϭ 6.39

(c) pOH ϭ 1.08



Practice Problem A Determine the pOH of a solution at 25°C in which the hydroxide ion

concentration is (a) 5.7 ϫ 10Ϫ12 M, (b) 7.3 ϫ 10Ϫ3 M, and (c) 8.5 ϫ 10Ϫ6 M.

Practice Problem B Determine the pH of a solution at 25°C in which the hydroxide ion

concentration is (a) 2.8 ϫ 10Ϫ8 M, (b) 9.9 ϫ 10Ϫ9 M, and (c) 1.0 ϫ 10Ϫ11 M.



Sample Problem 16.7

Calculate the hydroxide ion concentration in a solution at 25°C in which the pOH is (a) 4.91,

(b) 9.03, and (c) 10.55.



Apago PDF Enhancer



Strategy Given pOH, use Equation 16.5 to calculate [OHϪ].

Setup



(a) [OHϪ] ϭ 10Ϫ4.91

(b) [OHϪ] ϭ 10Ϫ9.03

(c) [OHϪ] ϭ 10Ϫ10.55

Solution

Think About It Use the



benchmark pOH values to

determine whether these solutions

are reasonable. In part (a), for

example, the pOH between 4 and

5 corresponds to [OHϪ] between

1 ϫ 10Ϫ4 M and 1 ϫ 10Ϫ5 M.



(a) [OHϪ] ϭ 1.2 ϫ 10Ϫ5 M

(b) [OHϪ] ϭ 9.3 ϫ 10Ϫ10 M

(c) [OHϪ] ϭ 2.8 ϫ 10Ϫ11 M



Practice Problem A Calculate the hydroxide ion concentration in a solution at 25°C in which the

pOH is (a) 13.02, (b) 5.14, and (c) 6.98.

Practice Problem B Calculate the hydronium ion concentration in a solution at 25°C in which the

pOH is (a) 2.74, (b) 10.31, and (c) 12.40.



Bringing Chemistry to Life

Antacids and the pH Balance in Your Stomach

An average adult produces between 2 and 3 L of gastric juice daily. Gastric juice is an acidic

digestive fluid secreted by glands in the mucous membrane that lines the stomach. It contains

hydrochloric acid (HCl), among other substances. The pH of gastric juice is about 1.5, which

corresponds to a hydrochloric acid concentration of 0.03 M—a concentration strong enough to

dissolve zinc metal!



bur75640_ch16_672-725.indd 680



11/19/09 4:02:03 PM



SEC T I ON 16. 3



The pH Scale



681



The inside lining of the stomach is made up of parietal cells, which are fused together to

form tight junctions. The interiors of the cells are protected from the surroundings by cell membranes. These membranes allow water and neutral molecules to pass in and out of the stomach, but

they usually block the movement of ions such as Hϩ, Naϩ, Kϩ, and ClϪ. The Hϩ ions come from

carbonic acid (H2CO3) formed as a result of the hydration of CO2, an end product of metabolism:

CO2(g) ϩ H2O(l)



H2CO3(aq)

Hϩ(aq) ϩ HCOϪ3 (aq)



H2CO3(aq)



These reactions take place in the blood plasma bathing the cells in the mucosa. By a process known

as active transport, Hϩ ions move across the membrane into the stomach interior. (Active transport

processes are aided by enzymes.) To maintain electrical balance, an equal number of ClϪ ions

also move from the blood plasma into the stomach. Once in the stomach, most of these ions are

prevented by cell membranes from diffusing back into the blood plasma.

The purpose of the highly acidic medium within the stomach is to digest food and to activate

certain digestive enzymes. Eating stimulates Hϩ ion secretion. A small fraction of these ions normally are reabsorbed by the mucosa, causing a number of tiny hemorrhages. About half a million

cells are shed by the lining every minute, and a healthy stomach is completely relined every few

days. However, if the acid content is excessively high, the constant influx of Hϩ ions through the

membrane back to the blood plasma can cause muscle contraction, pain, swelling, inflammation,

and bleeding.

One way to temporarily reduce the Hϩ ions concentration in the stomach is to take an antacid.

The major function of antacids is to neutralize excess HCl in gastric juice. The following table lists

the active ingredients of some popular antacids. The reactions by which these antacids neutralize

stomach acid are as follows:

NaHCO3(aq) ϩ HCl(aq)



NaCl(aq) ϩ H2O(l) ϩ CO2(g)



CaCO3(aq) ϩ 2HCl(aq)



CaCl2(aq) ϩ H2O(l) ϩ CO2(g)



MgCO3(aq) ϩ 2HCl(aq)



MgCl2(aq) ϩ H2O(l) ϩ CO2(g)



Mg(OH)2(s) ϩ 2HCl(aq)



MgCl (aq)PDF

ϩ 2H O(l)Enhancer

Apago



Al(OH)2NaCO3(s) ϩ 4HCl(aq)



2



2



AlCl3(aq) ϩ NaCl(aq) ϩ 3H2O(l) ϩ CO2(g)



Active Ingredients in Some Common Antacids

Commercial Name

Active Ingredients

Alka-Seltzer

Aspirin, sodium bicarbonate, citric acid

Milk of magnesia

Magnesium hydroxide

Rolaids

Dihydroxy aluminum sodium carbonate

TUMS

Calcium carbonate

Maalox

Sodium bicarbonate, magnesium carbonate

The CO2 released by most of these reactions increases gas pressure in the stomach, causing the person to belch. The fizzing that takes place when Alka-Seltzer dissolves in water is caused by carbon

dioxide, which is released by the reaction between citric acid and sodium bicarbonate:

3NaHCO3(aq) ϩ H3C6H5O7(aq)



3CO2(g) ϩ 3H2O(l) ϩ Na3C6H5O7(aq)



This effervescence helps to disperse the ingredients and enhances the palatability of the solution.



Checkpoint 16.3



The pH Scale



16.3.1 Determine the pH of a solution at 25°C in

which [Hϩ] ϭ 6.35 ϫ 10Ϫ8 M.



16.3.2 Determine [Hϩ] in a solution at 25°C if

pH ϭ 5.75.



a) 7.65



a) 1.8 ϫ 10Ϫ6 M



b) 6.80



b) 5.6 ϫ 10Ϫ9 M



c) 7.20



c) 5.8 ϫ 10Ϫ6 M



d) 6.35



d) 2.4 ϫ 10Ϫ9 M



e) 8.00



e) 1.0 ϫ 10Ϫ6 M

(Continued)



bur75640_ch16_672-725.indd 681



11/19/09 4:02:05 PM



682



C H A P T E R 16



Acids and Bases



Checkpoint 16.3



The pH Scale (continued)



16.3.3 Determine the pOH of a solution at 25°C

in which [OHϪ] ϭ 4.65 ϫ 10Ϫ3 M.



16.3.4 Determine [OHϪ] in a solution at 25°C if

pH ϭ 10.50.



a) 11.67



a) 3.2 ϫ 10Ϫ11 M



b) 13.68



b) 3.2 ϫ 10Ϫ4 M



c) 0.32



c) 1.1 ϫ 10Ϫ2 M



d) 4.65



d) 7.1 ϫ 10Ϫ8 M



e) 2.33



e) 8.5 ϫ 10Ϫ7 M



16.4 Strong Acids and Bases

Student Annotation: We indicate that

ionization of a strong acid is complete by using

) instead of the double,

a single arrow (

) in the equation.

equilibrium arrow (



Most of this chapter and Chapter 17 deal with equilibrium and the application of the principles of

equilibrium to a variety of reaction types. In the context of our discussion of acids and bases, however, it is necessary to review the ionization of strong acids and the dissociation of strong bases.

These reactions generally are not treated as equilibria but rather as processes that go to completion.

This makes the determination of pH for a solution of strong acid or strong base relatively simple.



Strong Acids

There are many different acids, but as we learned in Chapter 4, relatively few qualify as strong.

Multimedia

Acids and Bases—the dissociation of

strong and weak acids (interactive).



Strong Acid

Hydrochloric acid

Hydrobromic acid

Hydroiodic acid

Nitric acid

Chloric acid

Perchloric acid

Sulfuric acid



Ionization Reaction

H3Oϩ(aq) ϩ ClϪ(aq)

HCl(aq) ϩ H2O(l)

HBr(aq) ϩ H2O(l)

H3Oϩ(aq) ϩ BrϪ(aq)

HI(aq) ϩ H2O(l)

H3Oϩ(aq) ϩ IϪ(aq)

HNO3(aq) ϩ H2O(l)

H3Oϩ(aq) ϩ NOϪ3 (aq)

HClO3(aq) ϩ H2O(l)

H3Oϩ(aq) ϩ ClOϪ3 (aq)

HClO4(aq) ϩ H2O(l)

H3Oϩ(aq) ϩ ClOϪ4 (aq)

H2SO4(aq) ϩ H2O(l)

H3Oϩ(aq) ϩ HSOϪ4 (aq)



Apago PDF Enhancer



Student Annotation: Remember that

although sulfuric acid has two ionizable

protons, only the first ionization is complete.



Student Annotation: As we will see in

Section 16.5, a solution of equal concentration

but containing a weak acid has a higher pH.



It is a good idea to commit this short list of strong acids to memory.

Because the ionization of a strong acid is complete, the concentration of hydronium ion at

equilibrium is equal to the starting concentration of the strong acid. For instance, if we prepare a

0.10 M solution of HCl, the concentration of hydronium ion in the solution is 0.10 M. All the HCl

ionizes, and no HCl molecules remain. Thus, at equilibrium (when the ionization is complete),

[HCl] ϭ 0 M and [H3Oϩ] ϭ [ClϪ] ϭ 0.10 M. Therefore, the pH of the solution (at 25°C) is

pH ϭ Ϫlog (0.10) ϭ 1.00

This is a very low pH, which is consistent with a relatively concentrated solution of a strong acid.

Sample Problems 16.8 and 16.9 let you practice relating the concentration of a strong acid

to the pH of an aqueous solution.



Sample Problem 16.8

Calculate the pH of an aqueous solution at 25°C that is (a) 0.035 M in HI, (b) 1.2 ϫ 10Ϫ4 M in HNO3,

and (c) 6.7 ϫ 10Ϫ5 M in HClO4.

Strategy HI, HNO3, and HClO4 are all strong acids, so the concentration of hydronium ion in each

solution is the same as the stated concentration of the acid. Use Equation 16.2 to calculate pH.

Setup (a) [H3Oϩ] ϭ 0.035 M



(b) [H3Oϩ] ϭ 1.2 ϫ 10Ϫ4 M

(c) [H3Oϩ] ϭ 6.7 ϫ 10Ϫ5 M



bur75640_ch16_672-725.indd 682



11/19/09 4:02:05 PM



SEC T I ON 16. 4



Strong Acids and Bases



683



Solution (a) pH ϭ Ϫlog (0.035) ϭ 1.46



(b) pH ϭ Ϫlog (1.2 ϫ 10Ϫ4) ϭ 3.92

(c) pH ϭ Ϫlog (6.7 ϫ 10Ϫ5) ϭ 4.17



Practice Problem A Calculate the pH of an aqueous solution at 25°C that is

(a) 0.081 M in HI, (b) 8.2 ϫ 10Ϫ6 M in HNO3, and (c) 5.4 ϫ 10Ϫ3 M in HClO4.

Practice Problem B Calculate the pOH of an aqueous solution at 25°C that is

(a) 0.011 M in HNO3 , (b) 3.5 ϫ 10Ϫ3 M in HBr, and (c) 9.3 ϫ 10Ϫ10 M in HCl.



Think About It Again, note that when a

hydronium ion concentration falls between two

of the benchmark concentrations in Table 16.3,

the pH falls between the two corresponding pH

values. In part (b), for example, the hydronium ion

concentration of 1.2 ϫ 10Ϫ4 M is greater than 1.0 ϫ

10Ϫ4 M and less than 1.0 ϫ 10Ϫ3 M. Therefore, we

expect the pH to be between 4.00 and 3.00.



[H3O؉] (M)

Ϫ3



1.0 ϫ 10

1.2 ϫ 10Ϫ4*

1.0 ϫ 10Ϫ4



؊log [H3O؉]



pH



Ϫlog (1.0 ϫ 10Ϫ3)

Ϫlog (1.2 ϫ 10Ϫ3)

Ϫlog (1.0 ϫ 10Ϫ4)



3.00

3.92†

4.00



*[H3Oϩ] between two benchmark values

†pH between two benchmark values



Sample Problem 16.9

Calculate the concentration of HCl in a solution at 25°C that has pH (a) 4.95, (b) 3.45,

and (c) 2.78.

Strategy Use Equation 16.3 to convert from pH to the molar concentration of



Being comfortable with the benchmark

hydronium ion concentrations and the

corresponding pH values will help you avoid

some of the common errors in pH calculations.



hydronium ion. In a strong acid solution, the molar concentration of hydronium ion is

equal to the acid concentration.

Setup (a) [HCl] ϭ [H3Oϩ] ϭ 10Ϫ4.95

ϩ



Student Annotation: When we take the

inverse log of a number with two digits to the

right of the decimal point, the result has two

significant figures.



Ϫ3.45



(b) [HCl] ϭ [H3O ] ϭ 10



(c) [HCl] ϭ [H3Oϩ] ϭ 10Ϫ2.78

Solution (a) 1.1 ϫ 10Ϫ5 M



(b) 3.5 ϫ 10Ϫ4 M

(c) 1.7 ϫ 10Ϫ3 M



Think About It As pH decreases,

acid concentration increases.



Apago PDF Enhancer

Practice Problem A Calculate the concentration of HNO3 in a solution at 25°C that has pH



(a) 2.06, (b) 1.77, and (c) 6.01.

Practice Problem B Calculate the concentration of HBr in a solution at 25°C that has pOH

(a) 9.19, (b) 12.18, and (c) 10.96.



Strong Bases

The list of strong bases is also fairly short. It consists of the hydroxides of alkali metals (Group

1A) and the hydroxides of the heaviest alkaline earth metals (Group 2A). The dissociation of a

strong base is, for practical purposes, complete. Equations representing dissociations of the strong

bases are as follows:



Multimedia

Acids and Bases—ionization of a strong

base and weak base (interactive).



Group 1A hydroxides

LiOH(aq)

Liϩ(aq) ϩ OHϪ(aq)

NaOH(aq)

Naϩ(aq) ϩ OHϪ(aq)

KOH(aq)

Kϩ(aq) ϩ OHϪ(aq)

RbOH(aq)

Rbϩ(aq) ϩ OHϪ(aq)

CsOH(aq)

Csϩ(aq) ϩ OHϪ(aq)

Group 2A hydroxides

Ca(OH)2(aq)

Ca2ϩ(aq) ϩ 2OHϪ(aq)

Sr(OH)2(aq)

Sr2ϩ(aq) ϩ 2OHϪ(aq)

Ba(OH)2(aq)

Ba2ϩ(aq) ϩ 2OHϪ(aq)



Student Annotation: Recall that Ca(OH)2

and Sr(OH)2 are not very soluble, but what

does dissolve dissociates completely

[9 Section 4.3, Table 4.4] .



Again, because the reaction goes to completion, the pH of such a solution is relatively easy to

calculate. In the case of a Group 1A hydroxide, the hydroxide ion concentration is simply the

starting concentration of the strong base. In a solution that is 0.018 M in NaOH, for example,



bur75640_ch16_672-725.indd 683



11/19/09 4:02:08 PM



684



C H A P T E R 16



Acids and Bases



[OHϪ] ϭ 0.018 M. Its pH can be calculated in two ways. We can either use Equation 16.1 to

determine hydronium ion concentration,

[H3Oϩ][OHϪ] ϭ 1.0 ϫ 10Ϫ14

1.0 ϫ 10Ϫ14 __________

1.0 ϫ 10Ϫ14

[H3Oϩ] ϭ __________

ϭ

ϭ 5.56 ϫ 10Ϫ13 M

Ϫ

[OH ]

0.018

and then Equation 16.2 to determine pH:

pH ϭ Ϫlog (5.56 ϫ 10Ϫ13 M) ϭ 12.25

or we can calculate the pOH with Equation 16.3,

pOH ϭ Ϫlog (0.018) ϭ 1.75

and use Equation 16.6 to convert to pH:

pH ϩ pOH ϭ 14.00

pH ϭ 14.00 Ϫ 1.75 ϭ 12.25

Both methods give the same result.

In the case of a Group 2A metal hydroxide, we must be careful to account for the reaction

stoichiometry. For instance, if we prepare a solution that is 1.9 ϫ 10Ϫ4 M in barium hydroxide,

the concentration of hydroxide ion at equilibrium (after complete dissociation) is 2(1.9 ϫ 10Ϫ4 M)

or 3.8 ϫ 10Ϫ4 M—twice the original concentration of Ba(OH)2. Once we have determined the

hydroxide ion concentration, we can determine pH as before:

1.0 ϫ 10Ϫ14 __________

1.0 ϫ 10Ϫ14

[H3Oϩ] ϭ __________

ϭ

ϭ 2.63 ϫ 10Ϫ11 M

Ϫ

Ϫ4

[OH ]

3.8 ϫ 10

and

pH ϭ Ϫlog (2.63 ϫ 10Ϫ11 M) ϭ 10.58

or



Apago PDF Enhancer



pOH ϭ Ϫlog (3.8 ϫ 10Ϫ4) ϭ 3.42



pH ϩ pOH ϭ 14.00

pH ϭ 14.00 Ϫ 3.42 ϭ 10.58

Sample Problems 16.10 and 16.11 illustrate calculations involving hydroxide ion concentration, pOH, and pH.



Sample Problem 16.10

Calculate the pOH of the following aqueous solutions at 25°C: (a) 0.013 M LiOH,

(b) 0.013 M Ba(OH)2, (c) 9.2 ϫ 10Ϫ5 M KOH.

Strategy LiOH, Ba(OH)2, and KOH are all strong bases. Use reaction stoichiometry to determine



hydroxide ion concentration and Equation 16.4 to determine pOH.

Setup (a) The hydroxide ion concentration is simply equal to the concentration of the base.

Therefore, [OHϪ] ϭ [LiOH] ϭ 0.013 M.



(b) The hydroxide ion concentration is twice that of the base:

Think About It These are basic

pOH values, which is what we

should expect for the solutions

described in the problem. Note

that while the solutions in parts

(a) and (b) have the same base

concentration, they do not have

the same hydroxide concentration

and therefore do not have the

same pOH.



bur75640_ch16_672-725.indd 684



Ba(OH)2(aq)



Ba2ϩ(aq) ϩ 2OHϪ(aq)



Therefore, [OHϪ] ϭ 2 ϫ [Ba(OH)2] ϭ 2(0.013 M) ϭ 0.026 M.

(c) The hydroxide ion concentration is equal to the concentration of the base. Therefore,

[OHϪ] ϭ [KOH] ϭ 9.2 ϫ 10Ϫ5 M.

Solution (a) pOH ϭ Ϫlog (0.013) ϭ 1.89



(b) pOH ϭ Ϫlog (0.026) ϭ 1.59

(c) pOH ϭ Ϫlog (9.2 ϫ 10Ϫ5) ϭ 4.04



11/19/09 4:02:09 PM



SEC T I ON 16. 4



Strong Acids and Bases



685



Practice Problem A Calculate the pOH of the following aqueous solutions at 25°C:

(a) 0.15 M NaOH, (b) 8.4 ϫ 10Ϫ3 M RbOH, (c) 1.7 ϫ 10Ϫ5 M CsOH.

Practice Problem B Calculate the pH of the following aqueous solutions at 25°C:

(a) 9.5 ϫ 10Ϫ8 M NaOH, (b) 6.1 ϫ 10Ϫ2 M LiOH, (c) 6.1 ϫ 10Ϫ2 M Ba(OH)2.



Sample Problem 16.11

An aqueous solution of a strong base has pH 8.15 at 25°C. Calculate the original concentration of

base in the solution (a) if the base is NaOH and (b) if the base is Ba(OH)2.

Strategy Use Equation 16.6 to convert from pH to pOH and Equation 16.5 to determine the



hydroxide ion concentration. Consider the stoichiometry of dissociation in each case to determine the

concentration of the base itself.

Setup



Student Annotation: Remember to keep an

additional significant figure or two until the

end of the problem—to avoid rounding error

[9 Section 1.5] .



pOH ϭ 14.00 Ϫ 8.15 ϭ 5.85

(a) The dissociation of 1 mole of NaOH produces 1 mole of OHϪ. Therefore, the concentration of the

base is equal to the concentration of hydroxide ion.

(b) The dissociation of 1 mole of Ba(OH)2 produces 2 moles of OHϪ. Therefore, the concentration of

the base is only one-half the concentration of hydroxide ion.

Solution



Think About It Alternatively, we

could determine the hydroxide ion

concentration using Equation 16.3,



[H3Oϩ] ϭ 10Ϫ8.15 ϭ 7.1 ϫ 10Ϫ9 M



[OHϪ] ϭ 10Ϫ5.85 ϭ 1.41 ϫ 10Ϫ6 M

(b) [Ba(OH)2] ϭ −3 [OHϪ] ϭ 7.1 ϫ 10Ϫ7 M



and Equation 16.1,

1.0 ϫ 10Ϫ14

[OHϪ] ϭ ____________

7.1 ϫ 10Ϫ9 M

ϭ 1.4 ϫ 10Ϫ6 M



Apago PDF Enhancer

Practice Problem A An aqueous solution of a strong base has pH 8.98 at 25°C. Calculate the



Once [OHϪ] is known, the solution

is the same as shown previously.



(a) [NaOH] ϭ [OHϪ] ϭ 1.4 ϫ 10Ϫ6 M



concentration of base in the solution (a) if the base is LiOH and (b) if the base is Ba(OH)2.

Practice Problem B An aqueous solution of a strong base has pOH 1.76 at 25°C. Calculate the

concentration of base in the solution (a) if the base is NaOH and (b) if the base is Ba(OH)2.



Checkpoint 16.4



Strong Acids and Bases



16.4.1 Calculate the pH of a 0.075 M solution of perchloric acid

(HClO4) at 25°C.



16.4.3 What is the pOH of a solution at 25°C that is 1.3 ϫ 10Ϫ3 M in

Ba(OH)2?



a) 12.88



a) 2.89



b) 7.75



b) 2.59



c) 6.25



c) 3.19



d) 1.12



d) 11.11



e) 7.00



e) 11.41



16.4.2 What is the concentration of HBr in a solution with pH 5.89 at

25°C?



16.4.4 What is the concentration of KOH in a solution at 25°C that has

pOH 3.31?



a) 7.8 ϫ 10Ϫ9 M



a) 2.0 ϫ 10Ϫ11 M



b) 1.3 ϫ 10Ϫ6 M



b) 3.3 ϫ 10Ϫ1 M



c) 5.9 ϫ 10



Ϫ14



M



c) 3.3 ϫ 10Ϫ7 M



d) 8.1 ϫ 10Ϫ7 M



d) 4.5 ϫ 10Ϫ4 M



e) 1.0 ϫ 10Ϫ7 M



e) 4.9 ϫ 10Ϫ4 M

(Continued)



bur75640_ch16_672-725.indd 685



11/19/09 4:02:09 PM



Tài liệu bạn tìm kiếm đã sẵn sàng tải về

1 Brønsted Acids and Bases

Tải bản đầy đủ ngay(0 tr)

×