1. Part 2 – Acids and Bases
Topics we still need to remember and review and WILL BE ON THIS TOPIC 18 EXAM:
pH and pOH
[H+] and [OH-]
Strong acids and bases
Conjugate pairs

2. 18.1 Lewis acids and bases
A Lewis acid is a lone pair acceptor and a Lewis base is a lone pair donor.
When a Lewis base reacts with a Lewis acid a coordinate bond is formed.
A nucleophile is a Lewis base and an electrophile is a Lewis acid

3. Lewis Acid/Base Chemistry
Lewis acids are electron pair acceptors.
Lewis bases are electron pair donors.
All Brønsted–Lowry acids and bases are also called Lewis acids and bases.
There are compounds which do not meet the Brønsted–Lowry definition which meet the Lewis definition.

4. Lewis acids + bases
Coordinate covalent bond- formed between a Lewis acid and Lewis base when both electrons in a covalent bond are from the Lewis base

5. More examples

6. Comparing Ammonia’s Reaction with H+ and BF3

7. Nucleophiles
Nucleophile (“likes nucleus”) = electron-rich species that DONATES a lone pair to form a new (coordinate) covalent bond (Lewis BASE)

8. Electrophile
Electrophile (“likes electrons”) is an electron-deficient species that ACCEPTS a lone pair from another reactant to form a new (coordinate) covalent bond (Lewis ACID)

9. 18.2 Calculations involving acids and bases
The expression for the dissociation constant of a weak acid (Ka) and a weak base (Kb)
For a conjugate acid-base pair, Ka x Kb = Kw
The relationship between Ka and pKa is pKa = -logKa and between Kb and pKb is pKb = -logKb.

10. Kw is temperature dependent
Write the equation for the auto-ionization of water, then write the expression for the equilibrium constant

11. What’s up with water?

12. pH and pOH scales are inter-related

13. Summary

16. Dissociation constants
Weak acids and bases DO NOT fully dissociate in solution
You have to write the equilibrium expression and solve for the [H+] or [OH-]
Ka = acid dissociation constant
Kb = base dissociation constant

17. Generic Examples

18. Important Note
The values of Ka and Kb are constant for a specific acid and base respectively, at a specified temperature. Their values give a measure of the strength of the acid or base.
The higher the value of Ka or Kb at a particular temperature, the greater the ionization and so the acid or base is stronger.

20. Calculations with Ka and Kb

21. Write expression and solve
RICE box time
Write expression and solve

22. Write expression and solve
pH + pOH = 14
pOH = 2.20
[OH-] = = 6.3x10-3
RICE box time
Write expression and solve

23. Write expression and solve
RICE box time
Write expression and solve

24. Write expression and solve
RICE box time
Write expression and solve

25. pKa and pKb
To make the numbers easier to deal with, we convert the Ka and Kb values into their negative logarithms to the base 10, pKa and pKb.

26. Notes about pKa and pKb
pKa and pKb numbers are usually positive and have no units
The relationship beween Ka and pKa and between Kb and pKb is inverse
A change of one unit in pKa or pKa represents a ten-fold change in the value of Ka or Kb
pKa and pKb must be quoted at a specific temperature.
LIKE pH!!!!!!

27. Conjugate Acid/Base Pairs
The weaker an acid is, the lower its Ka value is, which means its Kb value is higher and its conjugate base is stronger.
Weak acid = strong conjugate base
Similarly, the weaker a base is, the stronger its conjugate acid. The stronger a base is, the weaker its conjugate acid.

28. Conjugate Acid/Base Pairs

29. 18.3 pH curves
The characteristics of the pH curves produced by the different combinations of strong and weak acids and bases.
An acid-base indicator is a weak acid or a weak base where the components of the conjugate acid-base pair have different colours.
The relationship between the pH range of an acid-base indicator, which is a weak acid, and its pKa value.
The buffer region on the pH curve represents the region where small additions of acid or base result in little or no change in pH.
The composition and action of a buffer solution.

30. Buffer solutions
A buffer solution is resistant to changes in pH on the addition of small amounts of acid or alkali
(non buffer solution below)

31. Buffers
Solutions of a weak conjugate acid–base pair that resist drastic changes in pH are called buffers.
These solutions contain relatively high concentrations (10–3 M or more) of both the acid and base. Their concentrations are approximately equal.

32. How do buffers work? Two main types
Acidic buffers maintain the pH below 7
Basic buffers maintain the pH above 7
Mixture of two solutions that contain both species in a conjugate acid-base pair so that they can react and neutralize either acids or bases added to the solutions
For example….

33. Acidic buffers
Made by mixing an aqueous solution of a weak acid with a solution of its salt of a strong alkali
Therefore, the mixture has a relatively high concentrations of both CH3COOH and CH3COO-, an acid and its conjugate base, which are ready to react and neutralize added OH- and H+

34. Basic buffers
Made by mixing an aqueous solution of a weak base with a solution of its salt of a strong acid
Therefore, the mixture has a relatively high concentrations of both CH3COOH and CH3COO-, an acid and its conjugate base, which are ready to react and neutralize added OH- and H+

35. Can you write equations to show how these species would neutralize any added H+ or OH-?

36. Making buffer solutions
The pH of a buffer is determined by the interactions of its components. Specifically, the pKa or pKa of the acid/base and the ratio of the initial concentrations of acid and salt, or base and salt, used in its preparation.
Bison

37. Ways to Make a Buffer
Mix a weak acid and a salt of its conjugate base or a weak base and a salt of its conjugate acid.
Add strong acid and partially neutralize a weak base or add strong base and partially neutralize a weak acid.

38. How a Buffer Works
Adding a small amount of acid or base only slightly neutralizes one component of the buffer, so the pH doesn’t change much.

39. Calculating the pH of a Buffer
For a weak acid: Ka = [H+][A–]/[HA]
Take –log of both sides:
–log Ka = –log[H+] + –log([A–]/[HA])
Rearrange:
–log[H+] = –log Ka +log([A–]/[HA])
Which is:
pH = pKa + log([A–]/[HA])
This equation is known as the Henderson– Hasselbalch equation. This applies only to buffers.

40. Henderson–Hasselbalch Equation
What is the pH of a buffer that is 0.12 M in lactic acid, CH3CH(OH)COOH, and 0.10 M in sodium lactate? Ka for lactic acid is 1.4 × 104.
pH = pKa + log([A–]/[HA])
= –log(1.4 × 10–4) + log[(0.10 M)/(0.12 M)]
= (–0.08) = 3.77

43. Buffer Capacity
The amount of acid or base the buffer can neutralize before the pH begins to change to an appreciable degree
Using the Henderson–Hasselbalch equation, pH will be the same for a conjugate acid–base pair of 1 M each or 0.1 M each; however, the buffer which is 1 M can neutralize more acid or base before the pH changes.
2 Days Older

44. pH Range
The range of pH values over which a buffer system works effectively
Optimal pH: where pH = pKa([HA] = [A–])
If one concentration is more than 10 times the other, the buffering action is poor; this means that the pH range of a buffer is usually ±1 pH unit from pKa.

45. Addition of a Strong Acid or a Strong Base to a Buffer
Adding of the strong acid or base is a neutralization reaction; calculate like stoichiometry problem to find [HA] and [A–] when all of the added acid or base reacts.
Use the Henderson–Hasselbalch equation to find pH.

46. Salt hydrolysis Salt of a strong acid + strong base = no hydrolysis
Salt of a weak acid + strong base = anion hydrolysis
Salt of a strong acid + weak base = cation hydrolysis
Salt of a weak acid + weak base = depends on Ka and Kb of species involved

47. Salt hydrolysis summary

48. Titration
In this technique, an acid (or base) solution of known concentration is slowly added to a base (or acid) solution of unknown concentration.
A pH meter or indicators are used to determine when the solution has reached the equivalence point: The amount of acid equals that exactly neutralizes
the base.

49. Titration of a Strong Acid with a Strong Base
From the start of the titration to near the equivalence point, the pH goes up slowly.
Just before (and after) the equivalence point, the pH rises rapidly.
At the equivalence point, pH = 7.
As more base is added, the pH again levels off.

50. Titration of a Strong Base with a Strong Acid
It looks like you “flipped over” the strong acid being titrated by a strong base.
Start with a high pH (basic solution); the pH = 7 at the equivalence point; low pH to end.

51. Titration of a Weak Acid with a Strong Base
Use Ka to find initial pH.
Find the pH in the “buffer region” using stoichiometry followed by the Henderson–Hasselbalch equation.
At the equivalence point the pH is >7. Use the conjugate base of the weak acid to determine the pH.
As more base is added, the pH levels off. This is exactly the same as for strong acids.

52. Ways That a Weak Acid Titration Differs from a Strong Acid Titration
A solution of weak acid has a higher initial pH than a strong acid.
The pH change near the equivalence point is smaller for a weak acid. (This is at least partly due to the buffer region.)
The pH at the equivalence point is greater than 7 for a weak acid.

53. Use of Indicators
Indicators are weak acids that have a different color than their conjugate base form.
Each indicator has its own pH range over which it changes color.
An indicator can be used to find the equivalence point in a titration as long as it changes color in the small volume change region where the pH rapidly changes.

54. Indicator Choice Can Be Critical!

55. Indicators summary

56. 1. strong acid and strong base

57. 2. Weak acid and strong base

58. 3. Strong acid and weak base

59. 4. Weak acid and weak base

60. 8.5 Acid Deposition
Rain is naturally acidic because of dissolved CO2 and has a pH of Acid deposition has a pH below 5.6.
Acid deposition is formed when nitrogen or sulfur oxides dissolve in water to form HNO3, HNO2, H2SO4, and H2SO3.
Sources of the oxides of sulfur and nitrogen and the effects of acid deposition should be covered.

61. Causes of acid deposition
All rain water is naturally acidic due to the presence of dissolved carbon dioxide (carbonic acid).
Acid rain refers to rain water with a pH below 5.6 due to the additional presence of acidic pollutants.
Acid deposition = all of the precipitation
Wet acid deposition- rain, snow, sleet, hail, fog, mist, dew fall to ground as aqueous precipitates
Dry acid deposition = acidifying particles or gases fall to the ground as dust or smoke and later dissolve in water to form acids.

62. Sulfur oxides
Produced from burning fossil fuels, particularly coal and heavy oil in power plants. Also smelting metals.

63. Nitrogen oxides
Produced mainly from internal combustion engines. Burning the fuel provides the heat energy necessary for nitrogen and oxygen in air to combine.

65. Impact on materials
(write some equations, list the effects, and the impact)

66. What happens to plants?

67. What happens to water?

68. What happens to human health?

69. List some things that are being done to help reduce these problems.