Download Chem 1B Chapter 14 Exercises

Chem 1B
Chapter 14 Exercises
Name:_______________________
Exercise #1
Brønsted-Lowry Acids and Bases
1.
Identify the Bronsted-Lowry acids and bases and the conjugate acid base pairs in the following
reactions.
(a) HCl(aq) + H2O(l) ⇌ H3O+(aq) + Cl–(aq)
(b) H2PO4–(aq) + NH3(aq) ⇌ NH4+(aq) + HPO42–(aq)
(c) NH4+(aq) + CN–(aq) ⇌ HCN(aq) + NH3(aq)
(d) H2SO4(aq) + H2O(l) ⇌ H3O+(aq) + HSO4–(aq)
2.
Write an equation for the acid-base reaction between NH4+ and CO32– in aqueous solution and identify
the acids and the bases in a reversible reaction.
3.
Hydrochloric acid is a strong acid. In a solution of 0.10 M HCl, what are the concentrations of HCl,
H3O+ and Cl– ?
4.
Hydrofluoric acid is a weak acid (Ka = 7.2 x 10–4). In 0.10 M HF solution, which is the major species,
HF, H3O+, or F– ? What is the approximate concentration of hydronium ion H3O+ ?
5.
What are the major species present in 0.10 M solutions of each of the following acids?
HNO3, HNO2, H2SO4, H3PO4, HClO4, and CH3COOH?
6.
Given the Ka values (in parentheses) for the following acids:
HCl (Ka is very large); HF (Ka = 7.2 x 10–4); HCN (Ka = 6.2 x 10–10); HNO2 (Ka = 4.0 x 10–4);
CH3COOH (Ka = 1.8 x 10–4), and C6H5COOH (Ka = 6.3 x 10–4).
(a) Rank (or list) these acids in order of decreasing strength (strongest first).
(b) List their conjugate bases in order of increasing strength (weakest base first) and write the Kb value
for each conjugate bases. (Kw = 1.0 x 10–14)
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Chem 1B
Chapter 14 Exercises
7.
Name:_______________________
For each acid-base equilibrium shown below, predict whether the equilibrium shifts to the right (Kc > 1)
or to the left (Kc < 1). (Refer to the Ka values given in No.6)
(a) HCl(aq) + C6H5CO2–(aq) ⇌ C6H5COOH(aq) + Cl–(aq)
(b) CH3COOH(aq) + NO2–(aq) ⇌ HNO2(aq) + CH3CO2–(aq)
(c) HF(aq) + CN–(aq) ⇌ HCN(aq) + F–(aq)
8.
Acetic acid, CH3COOH, has Ka = 1.8 x 10–5. Calculate the percent ionization of acetic acid in 1.0 M,
0.10 M, and 0.010 M solutions, respectively. How does dilution affect the degree of ionization of weak
acids?
9.
Given that pH = –log [H3O+] and pOH = –log [OH–], calculate the pH of the following solutions:
(a) A solution with [H3O+] = 5.0 x 10–3 M
(b) A solution with [OH–] = 2.0 x 10–3 M
(c) A 0.020 M HCl solution.
(d) A 0.020 M NaOH solution.
(e) A 0.020 M Ba(OH)2 solution.
10.
What are [H3O+] and [OH–] in solutions with the following pH?
(a) pH = 2.85
(b) pH = 9.40
11.
How many times is a solution with pH = 2 more acidic (has higher [H3O+]) than one with pH = 4?
12.
How many times is a solution with pH = 13 more basic (has higher [OH-]) than one with pH = 10?
13.
Pure water has an ionization constant Kw = 1.0 x 10–14 at 25oC and Kw = 5.5 x 10–14 at 100oC. What are
[H3O+] and [OH–] in pure water at 25oC and at 100oC, respectively? What are the pH of water at 25oC
and 100oC, respectively? Does the pH at 100 oC imply that hot water is neutral, acidic, or basic?
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Chem 1B
Chapter 14 Exercises
Name:_______________________
Exercises #2
1.
What the pH of 0.10 M HNO2 solution? (Ka = 4.0 x 10–4)
2.
What is the pH of 0.10 M NH3 solution? (NH3, Kb = 1.8 x 10–5)
3.
If the concentration of acetic acid in vinegar is 0.80 mol/L, calculate the [H3O+] and the pH of vinegar.
(Ka = 1.8 x 10–5 for acetic acid)
4.
A solution of 0.100 M lactic acid (HC3H6O3) has a pH = 2.43. Calculate Ka of lactic acid and its degree
of ionization at the above concentration.
5.
A solution of 0.20 M formic acid (HCOOH) is found to be 3.0% ionized. Write an equation for the
ionization of formic acid and calculate the pH of the solution and the Ka of formic acid?
6.
What is the pH of aqueous solution of 0.10 M ethylamine (C2H5NH2) (Kb = 5.6 x 10–4)
7.
Calculate the concentrations of H2SO4, HSO4–, SO42–, and H3O+ in 0.10 M H2SO4 and the pH of the
solution.
8.
What is the expected pH of each of the following solutions?
(a) 1.0 x 10–5 M HCl
(b) 1.0 x 10–7 M HCl
(c) 1.0 x 10–9 M HCl
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Chem 1B
Chapter 14 Exercises
Name:_______________________
Exercise #3
1.
Predict whether aqueous solution of each of the following salts (ionic compounds) would be acidic,
basic, or neutral. Explain your reasoning or support your statements with appropriate chemical
equations. (Ka for HC2H3O2 = 1.8 x 10–5; Ka for HCN = 6.2 x 10–10;
Kb for NH3 = 1.8 x 10–5)
(a) NaNO3
(b) NaCH3CO2
(c) NH4Cl
(d) NH4CH3CO2
(e) NH4CN
2.
(a) Calculate the dissociation constant, Kb, for the hydrolysis of acetate ion according to the following
equation.
C2H3O2–(aq) + H2O ⇌ HC2H3O2(aq) + OH–(aq).
(Ka for HC2H3O2 = 1.8 x 10–5)
(b) What is the pH of 0.050 M sodium acetate solution.
3.
(a) Calculate the ionization constant, Kb, for the following reaction:
PO43–(aq) + H2O ⇌ HPO42–(aq) + OH–(aq);
(HPO42– has Ka = 4.8 x 10–13)
(b) Calculate the pH of solution that is 0.10 M Na3PO4
4.
(a) Calculate the dissociation constant, Ka, for the hydrolysis of ammonium ion according to the
following equation.
NH4+(aq) + H2O ⇌ H3O+(aq) + NH3(aq);
(Kb for NH3 = 1.8 x 10–5)
(b) What is the pH of 0.10 M NH4Cl solution?
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Chem 1B
Chapter 14 Exercises
5.
Name:_______________________
Identify the Lewis acids and Lewis bases in the following reactions:
(a) Cu2+(aq) + 4NH3(aq) ⇌ Cu(NH3)42+(aq)
(b) CO2 + H2O ⇌ H2CO3(aq)
(c) SO2 + H2O ⇌ H2SO3(aq)
(d) Co2+ + 4Cl ⇌ CoCl42 (aq)
–
6.
–
Rank the following acids in order of increasing strength:
(a) HClO4, H3PO4, H2SO4;
(b) H3PO4, H3AsO4, H3SbO4 ;
(c) HOF, HOCl, HOBr, HOI;
(d) HOCl, HClO2, HClO3, HClO4;
(e) CH3COOH, CF3COOH, CCl3COOH;
7.
Rank the following compounds in increasing base strength:
(a) NH3, CH3NH2, C2H5NH3, (C2H5)2NH, and HONH2;
(b) Na2CO3, NaNO2, Na3PO4, NaF, and NaCN.
8.
(a) Identify acidic and basic oxides in the following list. (b) Which oxide will form the most acidic
solution? (c) Which oxide will form the most basic solution? (d) Which of these oxides has an
amphoteric property? (e) Write a balanced equation for the reaction of each of these oxides with water.
Na2O, MgO, Al2O3, SiO2, P4O10, SO3, and Cl2O7;
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Chem 1B
Chapter 14 Exercises
Name:_______________________
Answers
Exercise #1
1. (a) HCl(aq) + H2O(l) ⇌ H3O+(aq) + Cl–(aq)
Acid
Base
Conj.
Conj. Base
Acid
(b) H2PO4–(aq) + NH3(aq) ⇌ NH4+(aq) + HPO42–(aq)
Acid
Base
Conj.
Conj. Base
Acid
(c) PO43–(aq) + H2O(l) ⇌ HPO42–(aq) + Cl– (aq)
Base
Acid
Conj.
Conj. Base
Acid
(d) HPO42–(aq) + OH–(aq) ⇌ H2O(l) + PO43–(aq)
Acid
Base
Conj.
Conj. Base
Acid
2.
NH4+(aq) + CO32–(aq) ⇌ HPO42–(aq) + NH3(aq);
Acid
Base
Conj.
Conj. Base
Acid
3. [HCl] = 0.0 (HCl ionizes completely; [H3O+] = 0.10 M; [OH–] = 0.10 M;
4. HF is the major species (only a small fraction of HF ionizes); [H3O+] = 0.0085 M.
5. Major species in:
HNO3: H3O+ and NO3– (HNO3 is a strong acid; it ionizes completely);
HNO2: HNO2 (a weak acid; only a small fraction of HNO2 ionizes);
H2SO4: H3O+ and HSO4–; H2SO4 is a strong acid; the first hydrogen dissociates completely, but the
second does not;
H3PO4: H3PO4; it is a weak acid, a small fraction ionizes;
HClO4: H3O+ and ClO4– (HClO4 is a strong acid; it ionizes completely);
CH3CO2H: CH3CO2H (it is a weak acid.)
6. (a) Acid strength: HCl > HF > HNO2 > C6 H5COOH > CH3COOH > HCN;
(b) Conjugate base strength: CN– > CH3COO– > C6H5COO– > NO2– > F– > Cl–;
7. (a) HCl(aq) + C6H5CO2–(aq) ⇌ C6H5COOH(aq) + Cl–(aq); equilibrium shifts right; (Kc >> 1)
(b) CH3COOH(aq) + NO2–(aq) ⇌ HNO2(aq) + CH3CO2–(aq); equilibrium shifts left; (Kc < 1)
(c) HF(aq) + CN–(aq) ⇌ HCN(aq) + F–(aq); equilibrium shift right; (Kc >> 1)
8. Percent ionizations = 0.42% in 1.0 M; 1.3% in 0.10 M, and and 4.2% in 0.010 M.
9. (a) pH = 2.30;
(b) pH = 11.30;
(c) pH = 1.70; (d) 12.30;
(e) pH = 12.60
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Chem 1B
Chapter 14 Exercises
Name:_______________________
10. (a) [H3O+] = 1.4 x 10-3 M; [OH–] = 7.1 x 10–12 M
–
–
(b) [H3O+] = 4.0 x 10 12 M; [OH–] = 2.5 x 10 5 M
11. 100 times more;
12. 1000 time higher.
13. At 100 C, [H3O ] = [OH ] = 2.3 x 10–7 M; pH = 6.63; solution is neutral: [H3O+] = [OH–]
o
–
+
Exercises #2
3. [H3O+] = 3.8 x 10-3 M; pH = 2.42;
1. pH = 2.20;
2. pH = 11.13;
4. Ka = 1.4 x 10–4;
% ionization = 3.72%;
5. pH = 2.22;
Ka = 1.8 x 10–4;
6. pH = 11.87;
7. [H2SO4] = 0.00 M; [HSO4–] = 0.090 M; [SO42–] = 0.010 M; [H3O+] = 0.11 M;
8. (a) pH = 5.00;
(b) pH = 6.63;
pH = 0.96;
(c) pH = 7.00;
Exercise #3:
1. (a) Neutral;
(b) Basic
–10
2. (a) Kb = 5.6 x 10
;
(c) Acidic
(d) Neutral
(e) Basic
(b) pH = 8.72
3. (a) Kb = 2.1 x 10–2;
(b) pH = 12.57
4. (a) Ka = 5.6 x 10–10;
(b) pH = 5.13
5. (a) Cu2+(aq) + 4NH3(aq) ⇌ Cu(NH3)42+(aq)
L.A.
L.B.
(b) CO2 + H2O ⇌ H2CO3(aq)
L.A.
L.B.
(c) SO2 + H2O ⇌ H2SO3(aq)
L.A.
L.B.
(d) Co2+ + 4Cl– ⇌ CoCl42-(aq)
L.A.
L.B.
6. (a) H3PO4 < H2SO4 < HClO4;
(b) H3SbO4 < H3AsO4 < H3PO4;
(c) HOI < HOBr < HOCl < HOF;
(d) HOCl < HClO2 < HClO3 < HClO4;
(e) CH3COOH < CCl3COOH < CF3COOH;
7. (a) HONH2 < NH3 < CH3NH2 < C2H5NH3 < (C2H5)2NH;
(b) NaF < NaNO2 < NaCN < Na2CO3 < Na3PO4.
8. (a) Acidic oxides: SiO2, P4O10, SO3, and Cl2O7.
(b) Most acidic oxide: Cl2O7;
Basic oxides: Na2O and MgO
(c) Most basic oxide: Na2O;
(e) Na2O(s) + H2O  2NaOH(aq);
P4O10(s) + 6 H2O  4 H3PO4(aq);
(d) Amphoteric oxide: Al2O3.
MgO(s) + H2O  2Mg(OH)2(aq);
SO3(g) + H2O  H2SO4(aq);
Cl2O7(s) + H2O  4 HClO4(aq);
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