1. Chapter 16: ACIDS & BASES 16.1 Brønsted Acids and Bases 16.2
SECTIONS TO STUDY:
16.1
Brønsted Acids and Bases
16.2
The Acid-Base Properties of Water
16.3
The pH Scale
16.4
16.5
16.6
16.10
Strong Acids and Bases
Weak Acids and Acid Ionization Constant
Weak Bases and Base Ionization Constant
Acid-Base Properties of Salt Solutions

2. Acids & Bases, and Environment
Chemical industries, food industries, drug industries, paper industry etc.
Changes in the pH value of water are important to many organisms. Most organisms have adapted to life in water of a specific pH and may die if it changes even slightly. This is especially true of aquatic macroinvertebrates and fish eggs
Acids (lower pH) help in digestion, while bases (higher pH) helps to transport oxygen in the body.

3. Brønsted Acids and Bases
A Brønsted acid is a species that donates a proton.
(a proton donor).
HCl(aq) H2O(l) H3O+(aq) Cl-(aq)
A Conjugate base is what remains of the acid after the donation of a proton.
A Brønsted base is a species that accepts a proton.
(a proton acceptor).
NH3(aq) H2O(l) NH4+(aq) OH- (aq)
A Conjugate acid is a newly formed protonated species.
hydronium ion
acid
conjugate base
base
conjugate acid

4. Brønsted Acids and Bases
Any reaction, using the Brønsted theory, involves both an acid and a base.
Water is an amphoteric species which can act as a Brønsted acid or as a Brønsted base.

5. Brønsted Acids and Bases

6. Brønsted Acids and Bases
Exercise:
Identify acids, conjugate acids, bases and conjugate bases in the reactions below:
(a) H2SO4(aq) + H2O(l) HSO4‒ (aq) H3O+(aq)
(b) NH4+(aq) + H2O(l) H3O+(aq) NH3(aq)

7. Exercise:
(a) the conjugate base of H2PO4‒ ?
(b) the conjugate acid of O2–,
(c) the conjugate base of HSO4–
(d) the conjugate acid of HCO3– ?

8. The Acid-Base Properties of Water
Water is a very weak electrolyte. It undergoes ionization as:
H2O (l) H+ (aq) + OH‒ (aq)
2H2O (l) H3O+ (aq) + OH‒ (aq)
The equilibrium expression of water autoionization is:
Kw = [H+] [OH‒] or Kw = [H3O+] [OH‒]
Autoionization of water

9. Equilibrium Constant of Water Autoionization Reaction
Kw = [H+] [OH‒] or Kw = [H3O+] [OH‒]
At 25C for pure water:
[H3O+] = [OH‒] = 1.0 × 10-7 M.
Kw = (1.0 × 10-7)(1.0 × 10-7) = 1.0 × 10-14
One can influence the concentrations of [H3O+] or [OH‒] by adding an acid or base to pure water.
However, at 25C, the product of the hydronium ion and hydroxide ion concentrations is always constant.
If you add an acid to pure water, [H3O+] increases and [OH-] decreases (a new equilibrium position)  Kw remains constant.

10. Equilibrium Constant in Aqueous Solutions
For any aqueous solution:
Kw = [H3O+] [OH‒] = 1.0 × at 25C
When [H3O+] = [OH‒] ; the solution is neutral.
When [H3O+] > [OH‒] ; the solution is acidic.
When [H3O+] < [OH‒] ; the solution is basic.

11. Equilibrium Constant in Aqueous Solutions
16.2
Equilibrium Constant in Aqueous Solutions
Exercise:
Calculate [OH‒] in a pure water solution in which the concentration of protons is M at 25C. Is the solution acidic, basic or neutral?
[OH‒] = 8.33 * M

12. Equilibrium Constant in Aqueous Solutions
Exercise:
At 60C, the value of Kw is 1 ×
2H2O (l) H3O+ (aq) + OH‒ (aq)
Calculate [H+] and [OH‒] in a neutral solution at 60C.
3 × 10-7 M

13. The pH Meter
16.3
A pH meter is commonly used in laboratories to determine the pH of solutions.
The pH meter is an electrical device with a probe that measures the proton concentration in a solution and displays the pH value.

14. pH = ‒log [H3O+] or pH = ‒log [H+]
The pH Scale
16.3
The pH scale is a convenient way to express the acidity and basicity of a solution (the concentration of H3O+ ions).
The pH of a solution is defined as:
pH = ‒log [H3O+] or pH = ‒log [H+]
[H3O+] = 10-pH OR
[H+] = 10-pH

15. Higher pH values ( > 7.00 )
16.3
The pH Scale
Exercise:
Find the pH of a neutral solution.
A neutral solution has [H3O+] = [OH‒] .
Since Kw = [H3O+] [OH‒] = 1.0 × 10-14
then [H3O+] = 1.0 × 10-7 M
pH = - log (1.0 × 10-7) = 7.00
Lower pH values ( < 7.00 )
Neutral
Higher pH values ( > 7.00 )

16. 16.3
The pOH Scale
A pOH scale is analogous to the pH scale. It is defined as:
pOH = ‒ log [OH‒]
Also, [OH‒] = 10-pOH

17. What is the hydroxide ion concentration of a solution with pOH 8.35?
16.3
Exercise:
What is the pOH of a solution that has a hydroxide ion concentration of 4.3 x 10-2 M?
What is the hydroxide ion concentration of a solution with pOH 8.35?

18. The pH and pOH Scale Exercise:
16.3
The pH and pOH Scale
Exercise:
What is the hydroxide ion concentration of a solution that has pH of 9.45 at 25C?
[OH-] = 2.8 × 10-5 M

19. pH = ‒log [H3O+] or pH = ‒log [H+]
Summary of pH and pOH scale
pH = ‒log [H3O+] or pH = ‒log [H+]
[H3O+] = 10-pH
[H+] = 10-pH
pOH = ‒ log [OH‒]
Also, [OH‒] = 10-pOH
[H3O+] = [OH–] solution is neutral and pH is 7.
[H3O+] > [OH–], solution is acidic and has pH < 7.00.
[H3O+] < [OH–], solution is basic and has pH > 7.00.

20. Strong Acids and Bases
16.4
Strong acids and strong bases are completely ionized when dissolved in water.
How can you determine that this is a strong acid?

21. [H3O+] = [strong acid]initial
16.4
Strong Acids
Because strong acids are completely ionized in aqueous solutions, it is easy to determine the pH of that solution.
[H3O+] = [strong acid]initial
HNO3(aq) + H2O(l) NO3‒ (aq) H3O+(aq)
What will be the concentration of HNO3, NO3-, and H3O+, if 1 M of HNO3 solution has been prepared?

22. Exercise:
Calculate the pH of an aqueous solution that contains 0.1 M HNO3 at 25C.
What concentration of HNO3 should we take to prepare a solution of pH = 1 at 25C.
pH = 1
0.1M

23. Strong Bases
Like strong acids, strong bases are completely ionized when dissolved in water.
Group 1A Hydroxides
Group 2A Hydroxides

24. Concentration of Strong Bases
16.4
NaOH(aq) Na+ (aq) OH-(aq)
What will be the concentration of NaOH, OH-, and Na+, if 1 M of NaOH solution has been prepared?
[HO-] = [strong base]initial
Ca(OH)2(aq) Ca2+ (aq) OH-(aq)
What will be the concentration of Ca(OH)2, OH-, and Ca2+, if 1 M of Ca(OH)2 solution has been prepared?

25. pH = 13 0.1M Exercise: Find the pH of 0.1mol of NaOH at 25C?
What concentration of NaOH should we take to prepare a solution of pH = 13 at 25C?
pH = 13
0.1M

26. pH of Strong Bases pH = 13.3 = 0.1M Exercise:
16.4
pH of Strong Bases
Exercise:
At 25C, what is the pH of a solution that is 0.1 M in Ba(OH)2?
pH = 13.3
Exercise:
At 25C, determine the concentration of Ba(OH)2 if the pH of a solution is 13.3?
= 0.1M

27. Weak Acids Most of the acids are weak acids.
16.5
Weak Acids
Molar concentrations
Most of the acids are weak acids.
Weak acids are NOT completely ionized in water. They are ionized in water only to a limited extent.
The degree of dissociation of a weak acid depends on:
acid concentration.
equilibrium constant (T dependent)
Strong acids
Weak acids

28. The Acid Ionization Constant
16.5
The Acid Ionization Constant
The larger the Ka value, the stronger the acid is.

29. The Acid Ionization Constant
16.5
The Acid Ionization Constant
For a weak mono-protic acid HA:
HA(aq) + H2O (l) H3O+(aq) + A‒ (aq)
or:
HA(aq) H+(aq) A‒ (aq)
the equilibrium expression is:
Ka =
[H3O+] [A‒]
[HA]
Ka =
[H+] [A‒]
[HA] or
Ka is called the acid ionization constant. Its magnitude indicates how strong a weak acid is?
Larger Ka values indicates a stronger acid.
Smaller Ka values indicates a weaker acid.

30. Calculating pH from Ka Exercise:
16.5
Exercise:
Calculate the pH at 25°C of a 0.18 M solution of a weak acid that has Ka = 9.2 x 106.
Ka =
[H+] [A‒]
[HA]
= 9.2 x 106
We use the same method we studied in the previous chapter.
0.18 +x
− x
0.18 − x x
Continue on the next slide

31.  Ka = [H+] [A‒] [HA] = (x) (x) 0.18 - x = x2 0.18 - x = 9.2 x 106
16.5
Ka =
[H+] [A‒]
[HA] =
(x) (x) x = x2 x
= 9.2 x 106
Since Ka is very small compared to the initial acid concentration, we can make use of a useful approximation instead of solving the problem using a quadratic equation. ≈ x2 x x2
0.18
= 9.2 x 106
The 5% rule
Solving for x:
x = 1.3 x 103 M
pH = -log [H+]
= -log (1.3 x 103)
= 2.89
The above approximation is valid only when x is significantly smaller than the [HA]initial .
( x must be less than 5% of [HA]initial in order for the above approximation to be valid).
X 100% = 0.72%
1.3 x 103
0.18
Check the validity =>
< 5%  OK

32. Weak Bases
16.6
Most of the bases are weak bases. Weak bases are not ionized completely when dissolved in water.
The equilibrium expression for the ionization of a weak base is:
Kb is called the base ionization constant. Its magnitude indicates how strong a weak base is.
Larger Kb values indicates a stronger base.
Smaller Kb values indicates a weaker base.
weak base
conjugate acid
conjugate base
Kb =
[HB+] [OH‒]
[B]

33. The larger the Kb value, the stronger the base is.
Weak Bases
16.6
The larger the Kb value, the stronger the base is.

34. Solving Problems Involving Weak Bases
Exercise:
Determine the Kb of a weak base if a 0.50 M solution of the base has a pH of 9.59 at 25°C.
At 25°C:
pOH = – pH = – 9.59 = 4.41
[OH‒] = 10 4.41 = 3.89 × 105 M.
Kb =
[BH+][OH‒]
[B] =
(3.89 × 105)2
Kb = 3.0 × 109

35. For strong acids and weak bases
SUMMARY
For strong acids and weak bases
[H+] = [strong acid]initial
[OH-] = [strong base]initial
Be careful about stoichiometry of bases
For weak acids and weak bases
[H+]eq = [Calculate from ICE table]
[OH-]eq = [Calculate from ICE table]

36. Acid-Base Properties of Salt Solutions
16.10
Acid-Base Properties of Salt Solutions
Basic Salt Solutions
Acidic Salt Solutions
Neutral Salt Solutions
Salts in Which Both the Cation and the Anion Hydrolyze
Ka  Kb = Kw
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37. Acid-Base Properties of Salt Solutions
16.10
Acid-Base Properties of Salt Solutions
When a specific salt (an ionic compound) is dissolved in water, it breaks up into its ions (salt hydrolysis).
Under some circumstances, these ions behave like acids or bases.
NaF (s) Na+(aq) + F‒(aq)
The F‒ ion is the conjugate base of the weak HF acid. The F‒ ions react with water.
F‒(aq) + H2O(l) HF(aq) + OH‒(aq)
An anion that is a conjugate base of a weak acid reacts with water to produce OH‒ ions, making the solution basic.
H2O

38. A‒(aq) + H2O(l) HA(aq) + OH‒(aq)
16.10
Basic Salt Solutions
In general, an anion that is a conjugate base of a weak acid (HA) reacts with water to produce OH‒ ions, making the solution basic.
A‒(aq) + H2O(l) HA(aq) + OH‒(aq)
Examples of A‒ are:
fluoride ions (F‒)
acetate ions (CH3COO‒)
nitrite ions (NO2‒)
sulfite ions (SO32‒)
hydrogen carbonate ions (HCO3‒)
But why not chloride ions (Cl‒) , nitrate ions (NO3‒) or perchloride ions (ClO4‒) ?

39. Basic Salt Solutions Cl‒(aq) + H2O(l) HCl(aq) + OH‒(aq)
16.10
The Cl‒ ion is the conjugate base of the strong HCl acid, so it will not hydrolyze like the F‒ ion does. Thus, it doesn’t affect the pH of the solution as it doesn’t produce OH‒ ions.
Cl‒(aq) + H2O(l) HCl(aq) + OH‒(aq)
A‒(aq) + H2O(l) HA(aq) + OH‒(aq)
The A‒ ion is the conjugate base of the weak HA acid.

40. 16.20
Calculate the pH of a 0.10-M solution of sodium fluoride (NaF) at 25°C. (Ka of HF = 7.1 × 10-4 )
Setup
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41. Solution 16.20 Indeed, it’s a basic solution (pH = >7) 41
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42. 16.10
Acidic Salt Solutions
A cation that is a conjugate acid of a weak base reacts with water to produce hydronium ions, making the solution acidic.
NH4+(aq) + H2O(l) NH3 (aq) + H3O+ (aq)
The NH4+ ion is the conjugate acid of the weak NH3 base. You can determine its Ka from the Kb value of NH3 using the fact that: Kw = Ka  Kb .
NH4Cl (s) NH4+(aq) + Cl‒(aq)
H2O

43. Acid-Base Properties of Salt Solutions
16.10
Acid-Base Properties of Salt Solutions
Acidic Salt Solutions
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44. Calculate the pH of a 0.10-M solution of ammonium chloride (NH4Cl) at 25°C. (Kb of NH3 = 1.8 × 10-5 )
Setup
16.21
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45. Solution 16.21 Indeed, it’s an acidic solution (pH = <7) 45
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46. Identify the species that will react (hydrolyze) with water
Get help from Ka or Kb, if given

47. Acid-Base Properties of Salt Solutions
16.10
Acid-Base Properties of Salt Solutions
Acidic Salt Solutions
The metal ion in a dissolved salt can also react with water to produce an acidic solution.
The extent of hydrolysis is greatest for the small and highly charged metal cations such as Al3+, Cr3+, Fe3+, Bi3+, and Be2+.
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48. Acid-Base Properties of Salt Solutions
16.10
Acid-Base Properties of Salt Solutions
Acidic Salt Solutions
For example, when aluminum chloride dissolves in water, each Al3+ ion becomes associated with six water molecules.
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49. Acid-Base Properties of Salt Solutions
16.10
Acid-Base Properties of Salt Solutions
Acidic Salt Solutions
Al(OH)(H2O)52+ can undergo further ionization:
and so on.
It is generally sufficient, however, to take into account only the first stage of hydrolysis when determining the pH of a solution that contains metal ions.
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50. Acid-Base Properties of Salt Solutions
16.10
Acid-Base Properties of Salt Solutions
The extent of hydrolysis is greatest for the smallest and most highly charged metal ions because a compact, highly charged ion is more effective in polarizing the O—H bond and facilitating ionization.
This is why relatively large ions of low charge, including the metal cations of Groups 1A and 2A (the cations of the strong bases), do not undergo significant hydrolysis (Be2+ is an exception).
Thus, most metal cations of Groups 1A and 2A do not impact the pH of a solution.
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51. Acid-Base Properties of Salt Solutions
16.10
Acid-Base Properties of Salt Solutions
Neutral Salt Solutions
Similarly, anions that are conjugate bases of strong acids do not hydrolyze to any significant degree.
Consequently, a salt composed of the cation of a strong base and the anion of a strong acid, such as NaCl, produces a neutral solution.
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52. Acid-Base Properties of Salt Solutions
16.10
Acid-Base Properties of Salt Solutions
Neutral Salt Solutions
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53. Acid-Base Properties of Salt Solutions
16.10
Acid-Base Properties of Salt Solutions
Exercise:
Predict whether the a 0.10 M solution of each of the following salts will be basic, acidic or neutral.
(a) KNO2
(b) CH3NH3Br
(c) NaCl
(d) CsF
(e) AlCl3
(f) KCN
basic
acidic
neutral
basic
acidic
basic

54. Calculations Involving Salt Solutions
Exercise:
Predict whether the below salt would produce acidic, basic or neutral solution?
NH4NO2 (Ka = 5.6  10-10) and (Kb = 2.2  10-11)
Since Ka > Kb, sol. will be acidic

55. Exercise:
The Ka of hypochlorous acid (HCIO) is 3.0  10-8 at 25 °C. What is the % ionisation of hypochlorous acid in a M aqueous solution of HCIO at 25 °C?
0.015 +x
− x
0.015 − x x
HClO H+
ClO-
Ka =
[H+][ClO‒]
[HClO] =
(x)2
(0.015-x) =
(x)2
(0.015)
= 3.0  10-8
x = 2.12  10-5 = H+

56. [H+] % Ionization= [HClO] % Ionization= [0.015]  100 % [2.12  10-5 ]
= 0.14 %
THINK ABOUT IT: how to calculate [H+] and pH, if %ionization is given.

57. Calculations Involving Salt Solutions
Find the pH of a solution of M NaHCO3. (Ka of H2CO3 = 4.3 * 10-13)
pH = ~12.7