1. Acids and Bases
Chapter 16

2. Homework pages 532 – 535
1-7, 9 – 14, 16, 17, ,
37, 40, 42, 44, 45, 47, 52 – 56,
58, 59, , 66, 68, 71, 72
74, 82.

3. Acids
Have a sour taste. Vinegar owes its taste to acetic acid. Citrus
fruits contain citric acid.
React with certain metals to produce hydrogen gas.
React with carbonates and bicarbonates to produce carbon
dioxide gas
Bases
Have a bitter taste.
Feel slippery. Many soaps contain bases.
4.3

4. Arrhenius acid is a substance that produces H+ (H3O+) in water
Arrhenius base is a substance that produces OH- in water
4.3

5. A Brønsted acid is a proton donor A Brønsted base is a proton acceptor
conjugate acid
conjugate base
base
acid
15.1

6. The conjugate base of an acid is the species formed
when one proton is removed.
The conjugate acid of a base is the species formed
when one proton is added.
A conjugate acid-base pair differ by a proton.
acid + base conjugate base + conjugate acid
What is conjugate base of: HF, H2SO4, NH3?
What is conjugate acid of: O2-, SO42-, NH3?

7. The Conjugate Pairs in Some Acid-Base Reactions
Acid Base Base Acid
Conjugate Pair
Reaction HF H2O F– H3O+
Reaction 2 HCOOH CN– HCOO– HCN
Reaction NH CO32– NH HCO3–
Reaction H2PO4– OH– HPO42– H2O
Reaction H2SO N2H HSO4– N2H62+
Reaction HPO42– SO32– PO43– HSO3–

8. Identifying Conjugate Acid-Base Pairs
Problem: The following chemical reactions are important for industrial
processes. Identify the conjugate acid-base pairs.
HSO4-(aq) + CN-(aq) SO42-(aq) + HCN(aq)
(b) ClO-(aq) + H2O(l) HClO(aq) + OH-(aq)
(c) S2-(aq) + H2O(aq) HS-(aq) + OH-(aq)
(d) HS-(aq) + H2O(aq) H2S(aq) + OH-(aq)

9. [ ] - Acid-Base Properties of Water autoionization of water
H2O (l) H+ (aq) + OH- (aq)
autoionization of water O H + - [ ]
conjugate acid
base
H2O + H2O H3O+ + OH-
conjugate base
acid
15.2

10. The Ion Product of Water
Kc =
[H+][OH-]
[H2O]
H2O (l) H+ (aq) + OH- (aq)
[H2O] = constant
Kc[H2O] = Kw = [H+][OH-]
The ion-product constant (Kw) is the product of the molar concentrations of H+ and OH- ions at a particular temperature.
Solution Is
[H+] = [OH-]
neutral
At 250C
Kw = [H+][OH-] = 1.0 x 10-14
[H+] > [OH-]
acidic
[H+] < [OH-]
basic
15.2

11. What is the [H+] in 0.035M NaOH?
What is the concentration of OH- ions in a HCl solution whose hydrogen ion concentration is 1.3 M?
Kw = [H+][OH-] = 1.0 x 10-14
[H+] = 1.3 M
[OH-] = Kw
[H+]
1 x 10-14
1.3 =
= 7.7 x M
What is the [H+] in 0.035M NaOH?
15.2

12. pH – A Measure of Acidity
pH = -log [H+]
Solution Is
At 250C
neutral
[H+] = [OH-]
[H+] = 1 x 10-7
pH = 7
acidic
[H+] > [OH-]
[H+] > 1 x 10-7
pH < 7
basic
[H+] < [OH-]
[H+] < 1 x 10-7
pH > 7 pH
[H+]
15.3

13. pOH = -log [OH-] [H+][OH-] = Kw = 1.0 x 10-14
-log [H+] – log [OH-] = 14.00
pH + pOH = 14.00
15.3

14. The pH of rainwater collected in a certain region of the northeastern United States on a particular day was What is the H+ ion concentration of the rainwater?
pH = -log [H+]
[H+] = 10-pH =
= 1.5 x 10-5 M
The OH- ion concentration of a blood sample is 2.5 x 10-7 M. What is the pH of the blood?
pH + pOH = 14.00
pOH = -log [OH-]
= -log (2.5 x 10-7)
= 6.60
pH = – pOH = – 6.60 = 7.40
15.3

15. Calculating [H3O+], pH, [OH-], and pOH
Problem: A chemist dilutes concentrated hydrochloric
acid to make two solutions: (a) 3.0 M and (b) M.
Calculate the [H3O+], pH, [OH-], and pOH of the two
solutions at 25°C.
What is the [H3O+], [OH-], and pOH of a solution with
pH = 3.67? With pH = 8.05?

16. Strong Electrolyte – 100% dissociation
NaCl (s) Na+ (aq) + Cl- (aq)
H2O
Weak Electrolyte – not completely dissociated
CH3COOH CH3COO- (aq) + H+ (aq)
Strong Acids are strong electrolytes
HCl (aq) + H2O (l) H3O+ (aq) + Cl- (aq)
HNO3 (aq) + H2O (l) H3O+ (aq) + NO3- (aq)
HClO4 (aq) + H2O (l) H3O+ (aq) + ClO4- (aq)
H2SO4 (aq) + H2O (l) H3O+ (aq) + HSO4- (aq)
15.4

17. Weak Acids are weak electrolytes
HF (aq) + H2O (l) H3O+ (aq) + F- (aq)
HNO2 (aq) + H2O (l) H3O+ (aq) + NO2- (aq)
HSO4- (aq) + H2O (l) H3O+ (aq) + SO42- (aq)
H2O (l) + H2O (l) H3O+ (aq) + OH- (aq)
Strong Bases are strong electrolytes
NaOH (s) Na+ (aq) + OH- (aq)
H2O
KOH (s) K+ (aq) + OH- (aq)
H2O
Ba(OH)2 (s) Ba2+ (aq) + 2OH- (aq)
H2O
15.4

18. Weak Bases are weak electrolytes (NH3)
F- (aq) + H2O (l) OH- (aq) + HF (aq)
NO2- (aq) + H2O (l) OH- (aq) + HNO2 (aq)
Conjugate acid-base pairs:
The conjugate base of a strong acid has no measurable strength.
H3O+ is the strongest acid that can exist in aqueous solution.
The OH- ion is the strongest base that can exist in aqueous solution.
15.4

19. 15.4

21. Strong Acid
Weak Acid
15.4

22. What is the pH of a 2 x 10-3 M HNO3 solution?
HNO3 is a strong acid – 100% dissociation.
Start
0.002 M
0.0 M
0.0 M
HNO3 (aq) + H2O (l) H3O+ (aq) + NO3- (aq)
End
0.0 M
0.002 M
0.002 M
pH = -log [H+] = -log [H3O+] = -log(0.002) = 2.7
What is the pH of a 1.8 x 10-2 M Ba(OH)2 solution?
Ba(OH)2 is a strong base – 100% dissociation.
Start
0.018 M
0.0 M
0.0 M
Ba(OH)2 (s) Ba2+ (aq) + 2OH- (aq)
End
0.0 M
0.018 M
0.036 M
pH = – pOH = log(0.036) = 12.56
15.4

23. Weak Acids (HA) and Acid Ionization Constants
HA (aq) + H2O (l) H3O+ (aq) + A- (aq)
HA (aq) H+ (aq) + A- (aq)
Ka =
[H+][A-]
[HA]
Ka is the acid ionization constant
weak acid
strength Ka
15.5

24. 15.5

25. What is the pH of a 0.5 M HF solution (at 250C)?
Ka =
[H+][F-]
[HF]
= 7.1 x 10-4
HF (aq) H+ (aq) + F- (aq)
HF (aq) H+ (aq) + F- (aq)
Initial (M)
0.50
0.00
0.00
Change (M) -x +x +x
Equilibrium (M) x x x
Ka = x2 x
= 7.1 x 10-4
Ka << 1
0.50 – x  0.50
Ka  x2
0.50
= 7.1 x 10-4
x2 = 3.55 x 10-4
x = M
[H+] = [F-] = M
pH = -log [H+] = 1.72
[HF] = 0.50 – x = 0.48 M
15.5

26. When can I use the approximation?
Ka << 1
0.50 – x  0.50
When x is less than 5% of the value from which it is subtracted.
0.019 M
0.50 M
x 100% = 3.8%
Less than 5%
Approximation ok.
x = 0.019
What is the pH of a 0.05 M HF solution (at 250C)?
Ka  x2
0.05
= 7.1 x 10-4
x = M
0.006 M
0.05 M
x 100% = 12%
More than 5%
Approximation not ok.
Must solve for x exactly using quadratic equation or method of successive approximation.
15.5

27. EH Assignment
Due Mar 7
Unit 18
Minimum score
Sec 1 Acid-Base Reactions 90
Sec 2 Kw and pH
Sec 3 pH meter
Sec 4 Strong Acids and Bases
Sec 5 Weak Acids and Bases

28. Finding the Ka of a Weak Acid from the pH of its Solution and Vice Versa
Problem: The weak acid hypochlorous acid is formed in bleach
solutions. If the pH of a 0.12 M solution of HClO is 4.19, what is the
value of the Ka of this weak acid? What is the % ionization?
Ka = 3.5 x 10-8

29. Solving weak acid ionization problems:
Identify the major species that can affect the pH.
In most cases, you can ignore the autoionization of water.
Ignore [OH-] because it is determined by [H+].
Use ICE to express the equilibrium concentrations in terms of single unknown x.
Write Ka in terms of equilibrium concentrations. Solve for x by the approximation method. If approximation is not valid, solve for x exactly.
Calculate concentrations of all species and/or pH of the solution.
15.5

30. What is the pH of a 0.122 M monoprotic acid whose Ka is 5.7 x 10-4?
HA (aq) H+ (aq) + A- (aq)
Initial (M)
0.122
0.00
0.00
Change (M) -x +x +x
Equilibrium (M) x x x
Ka = x2 x
= 5.7 x 10-4
Ka << 1
0.122 – x  0.122
Ka  x2
0.122
= 5.7 x 10-4
x2 = 6.95 x 10-5
x = M M
0.122 M
x 100% = 6.8%
More than 5%
Approximation not ok.
15.5

31. Ka = x2 x
= 5.7 x 10-4
x x – 6.95 x 10-5 = 0
-b ± b2 – 4ac  2a
x =
ax2 + bx + c =0
x =
x =
HA (aq) H+ (aq) + A- (aq)
Initial (M)
Change (M)
Equilibrium (M)
0.122
0.00 -x +x x x
[H+] = x = M
pH = -log[H+] = 2.09
15.5

32. For a monoprotic acid HA
Ionized acid concentration at equilibrium
Initial concentration of acid
x 100%
percent ionization =
For a monoprotic acid HA
Percent ionization =
[H+]
[HA]0
x 100%
[HA]0 = initial concentration
15.5

33. The Stepwise Dissociation of Phosphoric Acid
Phosphoric acid is a weak acid, and normally only loses one proton
in solution, but it will lose all three when reacted with a strong base
with heat. The ionization constants are given for comparison.
H3PO4 (aq) + H2O(l) H2PO4-(aq) + H3O+(aq)
H2PO4-(aq) + H2O(l) HPO42-(aq) + H3O+(aq)
HPO42-(aq) + H2O(l) PO43-(aq) + H3O+(aq)
H3PO4 (aq) + 3 H2O(l) PO43-(aq) + 3 H3O+(aq)

34. 15.8

35. Terms to know: diprotic and polyprotic
Calculate the concentration of all species in 0.10 M H2S.

36. Weak Bases and Base Ionization Constants
NH3 (aq) + H2O (l) NH4+ (aq) + OH- (aq)
Kb =
[NH4+][OH-]
[NH3]
Kb is the base ionization constant
weak base
strength Kb
Solve weak base problems like weak acids except solve for [OH-] instead of [H+].
15.6

37. Most anions can act as weak bases.
The base-dissociation constant Kb refers to the
equilibrium that occurs when a weak base is
added to water.
Ammonia and amines are the most common
molecules that act as weak bases.
Most anions can act as weak bases.

38. 15.6

39. Determining pH from Kb and Initial [B]
Problem: Ammonia is commonly used cleaning agent in households and
is a weak base, with a Kb of 1.8 x What is the pH of a 1.5 M NH3
solution?

40. Ionization Constants of Conjugate Acid-Base Pairs
HA (aq) H+ (aq) + A- (aq) Ka
A- (aq) + H2O (l) OH- (aq) + HA (aq) Kb
H2O (l) H+ (aq) + OH- (aq) Kw
KaKb = Kw
Weak Acid and Its Conjugate Base
Ka = Kw Kb
Kb = Kw Ka
15.7

41. All anions can act as weak
bases except those that are
the conjugate bases of the
strong acids.

42. Determining the pH of a Salt Solution
Problem: Sodium cyanide in water can produce a basic solution. What
is the pH of a 0.25 M solution of NaCN?
Ka for HCN is 4.9 x 10-10
Plan: We have to find the pH of a solution of the cyanide ion, CN -,
which acts as a base in water:

43. The stronger an acid, the weaker is its conjugate
base and vice versa.
An acid-base reaction proceeds to the greater
extent in the direction to form the weaker acid
and base. If the weaker acid and base are on the
right Kc > 1
Sample Problem
H2PO4-(aq) + NH3(aq) NH4+(aq) + HPO42-(aq)
H2O(l) + HS-(aq) OH-(aq) + H2S(aq)

44. The reaction of an ion with water is frequently
called hydrolysis. Salt solutions are usually acidic
or basic because of hydrolysis of the cation or anion.
Cations: all cations act as acids except those in
groups 1 and 2.
Anions:
A few are acids HSO4-
Those that are conjugate bases of strong acids
are neutral. (Cl-, NO3-, …)
All others are bases.

46. Acid-Base Properties of Salts
Acid Solutions:
It is easy to see how a salt like NH4Cl can produce an acidic solution.
NH4Cl (s) NH4+ (aq) + Cl- (aq)
H2O
NH4+ (aq) NH3 (aq) + H+ (aq)
How can salts Al(NO3)3, CrCl3, or FeBr3 produce acidic solutions?
Al(H2O)6 (aq) Al(OH)(H2O)5 (aq) + H+ (aq) 3+ 2+
15.10

47. Acid Hydrolysis of Al3+
15.10

48. Problem 16.62
Which of these salts will undergo hydrolysis?
KF, NaNO3, NH4I, KCN, C6H5COONa,
RbI, Na2CO3, CaCl2, HCOOK, AlCl3.
Are the solutions acidic, basic or neutral?
Write hydrolysis equation.

49. Definition of An Acid
Arrhenius acid is a substance that produces H+ (H3O+) in water
A Brønsted acid is a proton donor
A Lewis acid is a substance that can accept a pair of electrons
A Lewis base is a substance that can donate a pair of electrons
+ OH- •
H O H • H+
acid
base
N H • H
N H H + H+ +
acid
base
15.12

50. No protons donated or accepted!
Lewis Acids and Bases
F B F
N H • H
F B F
N H H +
acid
base
No protons donated or accepted!
15.12

51. Problem 16.74
Describe this equation according to the
Lewis theory of acids and bases.
AlCl3 + Cl-  AlCl4-