1. Chapter 16
Acids and Bases

2. Drill
Determine which strong acid and strong base the following salts were derived from:
LiCl
Ba3(PO4)2
CaSO4
Sr(NO3)2

3. Objectives
SWBAT
Distinguish between Arrhenius, Bronsted Lowry and Lewis acids and bases.

4. Arrhenius Definition Definitions:
Acids produce hydrogen ions in aqueous solution.
Bases produce hydroxide ions when dissolved in water.
Limited to aqueous solutions.
Only one kind of base (hydroxide).
NH3 ammonia could not be an Arrhenius base.

5. Bronsted-Lowry Definitions
Definition: An acid is a proton (H+) donor and a base is a proton acceptor.
Acids and bases always come in pairs.
HCl is an acid..
When it dissolves in water it gives its proton to water.
HCl(g) + H2O(l) H3O+ + Cl-
Water is a base that makes a hydronium ion

6. Remember Strong acids completely dissociate in water.
HCl + H2O  H3O Cl-1
This reaction goes to completion and there is no HCl left in the solution.
Use a single direction arrow.

7. Remember Weak acids only partially dissociate.
CH3COOH + NH3 ⇔ CH3COO-1 + NH4+1
This is an equilibrium reaction.
There are significant amounts of reactants and products in the solution.
Use a double headed arrow. ⇔

8. Remember Hydroxides (and some oxides) are strong bases.
All other common bases are weak.
Weak bases establish an equilibrium system like acids.

9. Acid Base Pairs General equation HA(aq) + H2O(l) H3O+(aq) + A-(aq)
Acid + Base Conjugate acid + Conjugate base
This is an equilibrium situation.
There is competition for H+ between H2O and A-
The stronger base controls direction of the rxn.
If H2O is a stronger base it takes the H+
Equilibrium would then move to right.

11. CH3COOH + NH3 ⇔ CH3COO-1 + NH4+1
Use the following reaction and the conjugate acid/base chart to determine which direction the equilibrium will lie.
CH3COOH + NH3 ⇔ CH3COO-1 + NH4+1
CH3COOH is a stronger acid than NH4+1
NH3 is a stronger base than CH3COO-1
The equilibrium will favor the side in which the weaker acid and base a present.
Equilibrium will lie to the right.

12. Acid Dissociation Constant Ka
HA(aq) + H2O(l) H3O+(aq) + A -1(aq)
Ka = [H3O+1][A-1] [HA]
H3O+1 is often written H+1 ignoring the water in equation (it is implied).
Since this is the equilibrium constant associated with weak acid dissociation, this particular Kc is most commonly called the acid dissociation constant Ka

13. Acid Dissociation Constant Ka
HA(aq) H+(aq) + A-(aq)
Ka = [H+][A-] [HA]
We can write the expression for any acid.
Strong acids dissociate completely.
Equilibrium lies far to right.
Conjugate base must be weak.

14. Back to Pairs Strong acids Ka is large [H+] is equal to [HA]
A-1 is a weaker base than water
Weak acids
Ka is small
[H+] <<< [HA]
A-1 is a stronger base than water

15. Types of Acids Monoprotic Acids have only one hydrogen.
Polyprotic Acids: more than 1 acidic hydrogen (diprotic, triprotic).
Oxyacids: Proton is attached to the oxygen of an ion.
Organic acids: contain the Carboxyl group -COOH with the H attached to O
Generally very weak.

16. Amphoteric
Amphoteric means that the substance can behave as both an acid and a base.
Water molecules interact with each other and ionize. At the same time, the ions in solution reform molecules of water as shown in the following reaction. (This means that water auto-ionizes)
2H2O(l) H3O+1(aq) + OH-1 (aq)
KW = [H3O+][OH-] = [H+][OH-]

17. In pure water the concentrations of H3O+1 and OH-1 will always be equal.
[H+] = [OH-] = 1.0 x 10-7
At 25ºC KW = 1.0 x10-14
Therefore:
Neutral solution [H+] = [OH-]= 1.0 x10-7
Acidic solution [H+] > [OH-]
Basic solution [H+] < [OH-]

18. pH
In 1909, Danish biochemist S. P. L Sorensen introduced the pH system.
pH representing power of hydrogen

19. pH pH= -log[H+] Used because [H+] is usually very small
As pH decreases, [H+] increases exponentially
Other equations:
pOH= -log[OH-]
pKa = -log K

20. Sig Figs for pH
Sig figs: the number of sig figs in the lead number is the number of decimal places for the pH value. (only the digits after the decimal place of a pH are significant)
[H+] = 1.0 x pH= sig figs

21. Relationships Derivation: KW = [H+][OH-] -log KW = -log([H+][OH-])
-log KW = -log[H+]+ -log[OH-]
pKW = pH + pOH
KW = 1.0 x10-14
14.00 = pH + pOH
[H+],[OH-],pH and pOH
Given any one of these we can find the other three.

22. 100
10-1
10-3
10-5
10-7
10-9
10-11
10-13
10-14
[H+] 1 3 5 7 9 11 13 14 pH
Acidic
Neutral
Basic 1 3 5 7 9 11 13 14
pOH
Basic
100
10-1
10-3
10-5
10-7
10-9
10-11
10-13
10-14
[OH-]

23. Strong Acids HBr, HI, HCl, HNO3, H2SO4, HClO4
These acids completely dissociate
Therefore, [H+] = [HA]
10-14 = [H+][OH-]

24. Weak Acids Ka will be small. ALWAYS WRITE THE MAJOR SPECIES.
It will be an equilibrium problem from the start.
Determine whether most of the H+ will come from the acid or the water.
Compare Ka or Kw
Rest is just like equilibrium chapter.

25. Example
Calculate the pH of 2.0 M acetic acid HC2H3O2 with a Ka 1.8 x10-5
Calculate pOH, [OH-], [H+]

26. A Mixture of Weak Acids The process is the same.
Determine the major species.
The stronger will predominate.
Bigger Ka if concentrations are comparable
Calculate the pH of a mixture 1.20 M HF (Ka = 7.2 x 10-4) and 3.4 M HOC6H5 (Ka = 1.6 x 10-10)

27. Percent Dissociation = amount dissociated x 100 initial concentration
For a weak acid percent dissociation increases as acid becomes more dilute.
Calculate the % dissociation of M and M Acetic acid (Ka = 1.8 x 10-5
As [HA]0 decreases [H+] decreases but % dissociation increases.
Le Chatelier

28. The Other Way
What is the Ka of a weak acid that is 8.1 % dissociated as M solution?

29. Bases The OH- is a strong base.
Hydroxides of the alkali metals are strong bases because they dissociate completely when dissolved.
The hydroxides of alkaline earths Ca(OH)2 etc. are strong dibasic bases, but they don’t dissolve well in water.
Used as antacids because [OH- ] can’t build up.

30. Bases without OH- Bases are proton acceptors. NH3 + H2O NH4+ + OH-
It is the lone pair on nitrogen that accepts the proton.
Many weak bases contain N
B(aq) + H2O(l) BH+(aq) + OH- (aq)
Kb = [BH+][OH- ] [B]

31. Strength of Bases Hydroxides are strong. Others are weak.
Smaller Kb weaker base.
Calculate the pH of a solution of 4.0 M pyridine (Kb = 1.7 x 10-9) N:

32. Polyprotic Acids Always dissociate stepwise.
The first H+ comes of much easier than the second.
Ka for the first step is much bigger than Ka for the second.
Denoted Ka1, Ka2, Ka3

33. Polyprotic Acids What does K stand for?
Is it easier to remove the first or second ionizable proton?
Is is easier to remove the first.
The K values become successively smaller as successive protons are removed.
You will need to do 2 or more ice boxes.

34. Polyprotic Acid H2CO3 H+ + HCO3-1 Ka1= 4.3 x 10-7
HCO H+ + CO3-2 Ka2= 4.3 x 10-10
Base in first step is acid in second.
In calculations we can normally ignore the second dissociation.

35. Calculate the Concentration
Of all the ions in a solution of 1.00 M Arsenic acid H3AsO4
Ka1 = 5.0 x 10-3
Ka2 = 8.0 x 10-8
Ka3 = 6.0 x 10-10

36. Sulfuric Acid is Special
In first step it is a strong acid.
Ka2 = 1.2 x 10-2
Calculate the concentrations in a 2.0 M solution of H2SO4
Calculate the concentrations in a 2.0 x 10-3 M solution of H2SO4

37. Salts
A salt is an ionic compound formed by the reaction between an acid and a base.
Salts are strong electrolytes that completely dissociate into ions in water.
Salts of the cation of strong bases and the anion of strong acids are neutral.
for example NaCl, KNO3

38. Basic Salts
If the anion of a salt is the conjugate base of a weak acid - basic solution.
In an aqueous solution of NaF
The major species are Na+, F-, and H2O
F- + H2O HF + OH-
Kb =[HF][OH-] [F- ]
but Ka = [H+][F-] [HF]

39. Basic Salts
Ka x Kb = [HF][OH-] x [H+][F-] [F- ] [HF]

40. Basic Salts Ka x Kb = [HF][OH-] x [H+][F-] [F- ] [HF]
Ka x Kb =[OH-] [H+]
Ka x Kb = KW

41. Ka tells us Kb The anion of a weak acid is a weak base.
Calculate the pH of a solution of 1.00 M NaCN. Ka of HCN is 6.2 x 10-10
The CN- ion competes with OH- for the H+

42. Acidic Salts
A salt with the cation of a weak base and the anion of a strong acid will be basic.
The same development as bases leads to
Ka x Kb = KW
Calculate the pH of a solution of 0.40 M NH4Cl (the Kb of NH3 1.8 x 10-5).
Other acidic salts are those of highly charged metal ions.

43. Anion of weak acid, cation of weak base
Ka > Kb acidic
Ka < Kb basic
Ka = Kb Neutral

44. Structure and Acid Base Properties
Any molecule with an H in it is a potential acid.
The stronger the X-H bond the less acidic (compare bond dissociation energies).
The more polar the X-H bond the stronger the acid (use electronegativities).
The more polar H-O-X bond -stronger acid.

45. Strength of Oxyacids
The more oxygen hooked to the central atom, the more acidic the hydrogen.
HClO4 > HClO3 > HClO2 > HClO
Remember that the H is attached to an oxygen atom.
The oxygens are electronegative
Pull electrons away from hydrogen

46. Strength of Oxyacids
Electron Density Cl O H

47. Strength of Oxyacids
Electron Density O Cl O H

48. Strength of Oxyacids
Electron Density O Cl O H O

49. Strength of Oxyacids
Electron Density O O Cl O H O

50. Hydrated Metals
Highly charged metal ions pull the electrons of surrounding water molecules toward them.
Make it easier for H+ to come off. H
Al+3 O H

51. Acid-Base Properties of Oxides
Non-metal oxides dissolved in water can make acids.
SO3 (g) + H2O(l) H2SO4(aq)
Ionic oxides dissolve in water to produce bases.
CaO(s) + H2O(l) Ca(OH)2(aq)

52. Lewis Acids and Bases :N F H B F H F H Most general definition.
Acids are electron pair acceptors.
Bases are electron pair donors. F H B F :N H F H

53. Lewis Acids and Bases :N F H B F H F H
Boron triflouride wants more electrons. F H B F :N H F H

54. Lewis Acids and Bases F H F B N H F H
Boron triflouride wants more electrons.
BF3 is Lewis base NH3 is a Lewis Acid. F H F B N H F H

55. Lewis Acids and Bases ( ) H
Al+3
+ 6 O H +3 ( ) 6 H Al O H