Chapter 16. Acid-Base Equilibria Common Student Misconceptionstamapchemistryhart.weebly.com/uploads/3/8/0/0/38007377/chapter_16... · Chapter 16. Acid-Base Equilibria . Common Student

AP Chemistry Chapter 16. Acid-Base Equilibria

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Chapter 16. Acid-Base Equilibria

Common Student Misconceptions • Students often confuse a weak acid with a dilute acid. • Students have problems with the numerical parts of this chapter. They should be strongly encouraged to do

many problems on their own. • Students often have difficulty using their calculators to take decimal logarithms and antilogarithms. • Students confuse decimal (log) and natural (ln) logarithms. • Students have a difficult time accepting that acids may be both positively and negatively charged species. • Students tend to forget that [H3O+][OH-] is equal to 1.0 x 10-14 only at 25 °C. • Students often wrestle with neutralization reactions producing solutions with pH different from 7. • Determining pH of salts is often very conceptually and mathematically challenging.

16.1 Acids and Bases: A Brief Review

• Acids: taste sour and cause certain dyes to change color. • Bases: taste bitter and feel soapy. • Arrhenius concept of acids and bases: • An acid is a substance that, when dissolved in water, increases the concentration of H+ ions. • Example: HCl is an acid. • An Arrhenius base is a substance that, when dissolved in water, increases the concentration of OH–. • Example: NaOH is a base. • This definition is quite narrow in scope as it limits us to aqueous solutions.

16.2 Brønsted-Lowry Acids and Bases

• We can use a broader, more general definition for acids and bases that is based on the fact that acid-base

reactions involve proton transfers.

The H+ Ion in Water • The H+(aq) ion is simply a proton with no surrounding valence electrons. • In water, clusters of hydrated H+(aq) ions form. • The simplest cluster is H3O+(aq). • We call this a hydronium ion. • Larger clusters are also possible (such as H5O2

+ and H9O4+).

• Generally we use H+(aq) and H3O+(aq) interchangeably.


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Lewis structures and molecular models for H5O2

+ and H9O4+.

There is good experimental evidence for The existence of both these species.

Proton-Transfer Reactions • We will focus our attention on H+(aq): • According to the Arrhenius definitions, an acid increases [H+] and a base increases [OH–]. • Another definition of acids and bases was proposed by Brønsted and Lowry. • In the Brønsted-Lowry system, a Brønsted-Lowry acid is a species that donates H+ and a Brønsted-Lowry

base is a species that accepts H+. • Therefore a Brønsted-Lowry base does not need to contain OH–.

• Consider HCl(aq) + H2O(l) H3O+(aq) + Cl-

(aq) HCl donates a proton to water. ∴ HCl is an acid. H2O accepts a proton from HCl. ∴ H2O is a base.


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• NH3 is a Brønsted-Lowry base but not an Arrhenius base. • Consider NH3(aq) + H2O(l) NH4

+(aq) + OH–(aq): • H2O donates a proton to ammonia. • ∴ water is acting as an acid. • NH3 accepts a proton from water. • ∴ ammonia is acting as a base.

• Amphoteric substances can behave as acids and bases. • Thus water is an example of an amphoteric species.

Conjugate Acid-Base Pairs • Whatever is left of the acid after the proton is donated is called its conjugate base. • Similarly, a conjugate acid is formed by adding a proton to the base. • Consider HX(aq) + H2O(l) H3O+(aq) + X–(aq): • HX and X– differ only in the presence or absence of a proton. • These are said to be a conjugate acid-base pair. • X– is called the conjugate base. • After HX (acid) loses its proton it is converted into X– (base). • Therefore HX and X– are a conjugate acid-base pair. • After H2O (base) gains a proton it is converted into H3O+ (acid). • H3O+ is the conjugate acid. • Therefore, H2O and H3O+ are a conjugate acid-base pair.


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Sample Exercise 16.1 (p. 675)

a) What is the conjugate base of each of the following acids:

HClO4

H2S

PH4+

HCO3-?

b) What is the conjugate acid of each of the following bases?

CN-

SO42-

H2O

HCO3-

Practice Exercise 1 (16.1)

Consider the following equilibrium reaction: HSO4

-(aq) + OH-

(aq) SO42-

(aq) + H2O(l) Which substances are acting as acids in the reaction?

a) HSO4- and OH-

b) HSO4- and H2O

c) OH- and SO42-

d) SO42- and H2O


Write the formula for the conjugate acid of each of the following:

HSO3-

F-

PO43-

CO


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The hydrogen sulfite ion (HSO3-) is amphoteric.

a) Write an equation for the reaction of HSO3- with water, in which the ion acts as an acid.

b) Write an equation for the reaction of HSO3- with water, in which the ion acts as a base.

In both cases identify the conjugate acid-base pairs.


The dihydrogen phosphate ion, H2PO4-, is amphiprotic. In which of the following reactions is this ion

serving as a base?

(i) H3O+(aq) + H2PO4

- H3PO4(aq) + H2O(l)

(ii) H3O+(aq) + HPO4

2- H2PO4- + H2O(l)

(iii) H3PO4(aq) + HPO42- 2 H2PO4

- + H2O(l)

a) (i) only

b) (i) and (ii)

c) (i) and (iii)

d) (ii) and (iii)

e) (i), (ii) and (iii)


When lithium oxide (Li2O) is dissolved in water, the solution turns basic from the reaction of the oxide ion

(O2-) with water. Write the reaction that occurs, and identify the conjugate acid-base pairs.


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Relative Strengths of Acids and Bases

• The stronger an acid is, the weaker its conjugate base will be. • We can categorize acids and bases according to their behavior in water. • 1. Strong acids completely transfer their protons to water. • No undissociated molecules remain in solution. • Their conjugate bases have negligible tendencies to become protonated. • Example: HCl. • 2. Weak acids only partially dissociate in aqueous solution. • They exist in solution as a mixture of molecules and component ions. • Their conjugate bases show a slight tendency to abstract protons from water. • These conjugate bases are weak bases. • Example: Acetic acid is a weak acid; acetate ion (conjugate base) is a weak base. • 3. Substances with negligible acidity do not transfer a proton to water. • Example: CH4. • In every acid-base reaction, the position of the equilibrium favors the transfer of a proton from the stronger

acid to the stronger base. • H+ is the strongest acid that can exist in equilibrium in aqueous solution. • OH– is the strongest base that can exist in equilibrium in aqueous solution.

Figure 16.4. Relative strengths of select conjugate acid-base pairs. The two memebers of each pair are listed opposite each other in the two columns.


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For the following proton-transfer reaction, use the above figure (Figure 16.4) to predict whether the equilibrium lies to the left (Kc < 1) or to the right (Kc > 1):

HSO4-(aq) + CO3

2-(aq) SO4

2-(aq) + HCO3

-(aq)


Based on the information in Figure 16.4, place the following equilibria in order from smallest to largest value of Kc:

(i) CH3COOH(aq) + HS-(aq) CH3COO-

(aq) + H2S(aq)

(ii) F-(aq) + NH4

+(aq) HF(aq) + NH3(aq)

(iii) H2CO3 (aq) + Cl-(aq) HCO3

-(aq) + HCl(aq)

a) (i) < (ii) < (iii) b) (ii) < (i) < (iii) c) (iii) < (i) < (ii) d) (ii) < (iii) < (i) e) (iii) < (ii) < (i)


For each of the following reactions, use Figure 16.4 to predict whether the equilibrium lies predominantly to the left or to the right:

a) HPO42-

(aq) + H2O(l) H2PO4-(aq) + OH-

(aq)

b) NH4+

(aq) + OH-(aq) NH3(aq) + H2O(l)


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16.3 The Autoionization of Water • In pure water the following equilibrium is established:

2H2O(l) H3O+(aq) + OH–(aq) • This process is called the autoionization of water.

The Ion Product of Water

• We can write an equilibrium constant expression for the autoionization of water.

• Because H2O(l) is a pure liquid, we can simplify this expression:

• Kw is called the ion-product constant. • At 25oC the ion-product of water is:

• This applies to pure water as well as to aqueous solutions. • A solution is neutral if [OH–] = [H3O+]. • If the [H3O+] > [OH–], the solution is acidic. • If the [H3O+] < [OH–], the solution is basic.


Calculate the values of [H+] and [OH-] in a neutral solution at 25oC.


In a certain acidic solution at 25oC, [H+] is 100 times greater than [OH-]. What is the value for [OH-] for the solution?

a) 1.0 x 10-8 M b) 1.0 x 10-7 M c) 1.0 x 10-6 M d) 1.0 x 10-2 M e) 1.0 x 10-9 M

[ ][ ][ ]2

2

3

OHOHOH

Keq

−+

=

[ ] [ ][ ] weq KOHOHKOH == −+3

22

[ ]−

+− ==× OHOHK w 3

14100.1


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Indicate whether solutions with each of the following ion concentrations is neutral, acidic, or basic:

a) [H+] = 4 x 10-9 M

b) [OH-] = 1 x 10-7 M

c) [OH-] = 7 x 10-13 M


Calculate the concentration of H+(aq) in

a) a solution in which [OH-] is 0.010 M (1.0 x 10-12 M)

b) a solution in which [OH-] is 1.8 x 10-9 M (5.0 x 10-6 M)

Assume T = 25oC.


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A solution has [OH-] = 4.0 x 10-8 M. What is the value of [H+] for the solution?

a) 2.5 x 10-8 M

b) 4.0 x 10-8 M

c) 2.5 x 10-7 M

d) 2.5 x 10-6 M

e) 4.0 x 10-6 M


Calculate the concentration of OH-(aq) in a solution in which

a) [H+] = 2 x 10-6 M

b) [H+] = [OH-]

c) [H+] = 100 x [OH-]


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16.4 The pH Scale

• In most solutions [H+] is quite small. • We express the [H+] in terms of pH.

pH = −log[H+] = −log[H3O+]

• Note that this is a logarithmic scale. • Thus a change in [H+] by a factor of 10 causes the pH to change by 1 unit. • Most pH values fall between 0 and 14. • In neutral solutions at 25oC, pH = 7.00. • In acidic solutions, [H+] > 1.0 x 10–7, so pH < 7.00. • As the pH decreases, the acidity of the solution increases. • In basic solutions, [H+] < 1.0 x 10–7, so pH > 7.00. • As the pH increases, the basicity of the solution increases (acidity decreases).


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Calculate the pH values for the two solutions described in Sample Exercise 16.5.

(a) 12.00 b) 5.25)


A solution at 25oC has [OH-] = 6.7 x 10-3 M. What is the pH of the solution?

a) 0.83

b) 2.2

c) 2.17

d) 11.83

e) 12


a) In a sample of lemon juice [H+] is 3.8 x 10-4 M. What is the pH?

(3.42)

b) A commonly available window-cleaning solution has a [OH-] of 1.9 x 10-6 M. What is the pH?

(8.28)


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A sample of freshly pressed apple juice has a pOH of 10.24. Calculate [H+].

(1.7 x 10-4 M)


A solution at 25oC has pOH = 10.53. Which of the following statements is or are true?

(i) The solution is acidic.

(ii) The pH of the solution is 14.00 – 10.53.

(iii) For this solution, [OH-] = 10-10.53 M.

a) Only one of the statements is true.

b) Statements (i) and (ii) are true.

c) Statements (i) and (iii) are true.

d) Statements (ii) and (iii) are true.

e) All three statements are true.


A solution formed by dissolving an antacid tablet has a pOH of 4.82. Calculate [H+].

(6.6 x 10-10 M)


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Other “p” Scales • We can use a similar system to describe the [OH–].

pOH = −log[OH–] • Recall that the value of Kw at 25oC is 1.0 x 10–14. • Thus we can describe a relationship between pH and pOH:

−log[H+]+ (−log[OH–]) = pH + pOH = −logKw = 14.00

Measuring pH

• The most accurate method to measure pH is to use a pH meter. • However, certain dyes change color as pH changes. • They are called acid-base indicators. • Indicators are less precise than pH meters. • Many indicators do not have a sharp color change as a function of pH. • Most acid-base indicators can exist as either an acid or a base. • These two forms have different colors. • The relative concentration of the two different forms is sensitive to the pH of the solution. • Thus, if we know the pH at which the indicator turns color, we can use this color change to

determine whether a solution has a higher or lower pH than this value. • Some natural products can be used as indicators (tea is colorless in acid and brown in base; red cabbage

extract is another natural indicator).


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16.5 Strong Acids and Bases

Strong Acids • The most common strong acids are HCl, HBr, HI, HNO3, HClO3, HClO4, and H2SO4. MEMORIZE THESE! • Strong acids are strong electrolytes. • All strong acids ionize completely in solution: • Example: Nitric acid ionizes completely in solution.

HNO3(aq) + H2O(l) H3O+(aq) + NO3–(aq)

• Since H+ and H3O+ are used interchangeably, we write

HNO3(aq) H+(aq) + NO3–(aq)

• In solution the strong acid is usually the only source of H+. • Therefore, the pH of a solution of a monoprotic acid may usually be calculated directly from the initial

molarity of the acid. • Caution: If the molarity of the acid is less than 10–6 M then the autoionization of water needs to be taken

into account.


What is the pH of a 0.040 M solution of HClO4? (1.40)

Practice Exercise 1 (16.8) Order the following three solutions from smallest to largest pH.

(i) 0.20 M HClO3 (ii) 0.0030 M HNO3 (iii) 1.50 M HCl



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An aqueous solution of HNO3 has a pH of 2.34. What is the concentration of the acid? (0.0046 M)

Strong Bases • The most common strong bases are ionic hydroxides of the alkali metals or the heavier alkaline earth metals

(e.g., NaOH, KOH, and Ca(OH)2 are all strong bases). • Strong bases are strong electrolytes and dissociate completely in solution. • For example:

NaOH(aq) Na+(aq) + OH–(aq) • The pOH (and thus the pH) of a strong base may be calculated using the initial molarity of the base. • Not all bases contain the OH– ion. • Ionic metal oxides, hydrides, and nitrides are basic. • The oxide, hydride and nitride ions are stronger bases than hydroxide. • They are thus able to abstract a proton from water and generate OH–.

O2–(aq) + H2O(l) 2OH–(aq) H–(aq) + H2O(l) H2(g) + OH–(aq)

N3–(aq) + 3H2O(l) NH3(aq) + 3OH–(aq)


What is the pH of

a) a 0.028 M solution of NaOH? (12.45)

b) a 0.0011 M solution of Ca(OH)2? (11.34)


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Order the following three solutions from smallest to largest pH:

(i) 0.030 M Ba(OH)2 (ii) 0.040 M KOH (iii) Pure water



What is the concentration of a solution of

a) KOH for which the pH is 11.89? (7.8 x 10-3 M)

b) Ca(OH)2 for which the pH is 11.68? (2.4 x 10-3 M)

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Chapter 16. Acid-Base Equilibria Common Student Misconceptionstamapchemistryhart.weebly.com/uploads/3/8/0/0/38007377/chapter_16... · Chapter 16. Acid-Base Equilibria . Common Student