1. Acid Base Equilibria

2. Types of Acids H+ vs H3O+ Arrhenius Definition
Acids – Increase [H+] when dissolved in water
HCl
Bases – Increase [OH-] when dissolved in water
NaOH
Brønsted-Lowry Definition
Acids – Donates protons
Bases – Accepts protons
HCl(g) + H2O(l) → H3O+(aq) + Cl-(aq)
HCl(g) + NH3(g) → NH4Cl(s)
NH3(aq) + H2O(l) ↔ NH4+(aq) + OH-(aq)

3. Conjugate Acid Base Pairs
HX(aq) + H2O(l) ↔ X-(aq) + H3O+(aq)
HNO2(aq) + H2O(l) ↔ NO2-(aq) + H3O+(aq)
What is the conjugate base of each of the following acids?
HClO4, H2S, HCO3-, NH4+
What is the conjugate acid of each of the following bases?
CN-, SO42 -, H2O, HCO3-

4. Practice
There are two possible reactions that HSO4- can have with water. Write the reaction in which it acts as an acid and another where it acts as a base.
When lithium oxide is dissolved in water, the solution turns basic from the reaction of the oxide ion(O2-) with water. Write the reaction and identify the conjugate acid base pairs.

5. Acid Base Strength
Acid strength depends on how easily it gives up a proton
What are the strong acids?
Bases strength depends on how readily it accepts one.
What are the strong bases?
Strong Acids: HCl, HBr, HI, HClO3, HClO4, HNO3, H2SO4
Strong Bases: Group 1A metal hydroxides
Heavy Group 2A metal hydroxides
(from Ca down)

6. Acid Base Strength HX + H2O ↔ H3O+ + X-
Equilibrium lies on the side with weaker base
HCl + H2O ↔ H3O+ + Cl-
HC2H3O2 + H2O ↔ H3O+ + C2H3O2-

7. Practice
Identify whether equilibrium lies predominantly to the left or right.
HSO4- + CO32 - ↔ SO HCO3-
HPO H2O ↔ H2PO4- + OH-
NH4+ + OH- ↔ NH3 + H2O
Ans: R, L, R

8. Autoionization of Water
H2O + H2O ↔ H3O+ + OH-
What is the Ke q of this reaction
At 25°C Ke q = 1x101 4
What are the concentrations of each ion if each the concentration of [H3O+]=[OH-]?
How can we use this relationship to identify if a solution is acidic, basic, or neutral?

9. pH Scale pH scale goes from 0 to 14 pH = - log[H+] pOH = -log [OH-]
pH + pOH = 14

10. Strong Acids and Bases
For strong acids and bases no equilibrium is reached.
For monoprotic strong acids: [H+] = [HA]
What is the pH of a 0.040M solution of HClO4
pH = -log(0.040) = 1.40
For strong bases: [OH-] = to # of OH-'s x [Base]
What is the pH of a 0.028M solution of NaOH and a M solution of Ca(OH)2

11. Practice
An aqueous solution of HNO3 has a pH of What is the concentration of the acid?
What is the concentration of a solution of KOH for which the pH is 11.89; Ca(OH)2 for which the pH is 11.68?
Ans: # M, #2 7.8x10- 3M, 2.4x10- 3M

12. Weak Acids
Weak acids only partially ionize and therefore reach equilibrium.
Ke q can be used to tell what extent it ionizes, how?
HA(aq) + H2O(l) ↔ A-(aq) + H3O+(aq)
HA(aq) ↔ H+(aq) + A-(aq)
What is the equilibrium expression for this reaction?
Ka = acid-dissociation constant
How does Ka relate to acid strength?

13. Weak Acids

14. Using pH to find Ka
A Student prepared a 0.10M solution of formic acid (HCHO2) and measured its pH using a pH meter. The pH was found to be Calculate the Ka for formic acid at this temperature. What percentage of the acid is ionized?
Ans: 1.8x10- 4, 4.2%

15. Practice
Niacin, a type of B vitamin has the formula C5H4NCO2H. A 0.020M solution of the vitamin has a pH of What percentage of the acid is ionized in this solution? What is the acid dissociation constant Ka for niacin?
Ans: 2.7%, 1.6x10- 5

16. Using Ka to find pH
We need to know Ka and the initial concentration of the acid. Ka of acetic acid = 1.8x10- 5, [HC2H3O2] = 0.30M
Write the ionization equilibrium for the acid
HC2H3O2(aq) ↔ H+(aq) + C2H3O2-(aq)
Write the equilibrium constant expression and value of Ka
Ka = [H+][C2H3O2-]/[HC2H3O2]
Find equilibrium concentrations

17. Using Ka to find pH
4) Substitute the equilibrium concentrations into the equilibrium expression and solve for x.
** x may be disregarded as long as it is less than 5% of the initial [ ]
x = 2.3x10- 3 = [H+]
pH = 2.64

18. Practice
Calculate the pH of a 0.20M solution of HCN. Ka for HCN is 4.9x
Ans: pH = 5.00

19. Calculating Percent Ionization
Calculate the percentage of HF molecules ionized in (a) 0.10M HF solution; (b) 0.010M HF solution.
Find x then divide that by initial concentration.
Ans: a) 7.9%, 23%

20. Polyprotic Acids
Polyprotic acids have more than one ionizable H atoms and thus multiple Ka's.
Easier to remove 1s t proton than the next ones.
Finding pH for polyprotic acids
Compare size of Ka's
H2CO3(aq) ↔ H+(aq) + HCO3-(aq) Ka 1 = 4.3x10- 7
HCO3-(aq) ↔ H+(aq) + CO32 -(aq) Ka 2 = 5.6x
For most acids only Ka 1 is important

21. Polyprotic Acids

22. pH of Polyprotic Acids
What is the pH of a M solution of H2CO3? What is the [CO32 -] in the solution? Ka 1= 4.3x10-7, Ka 2 = 5.6x
Ans: pH = 4.40, [CO32 -] = 5.6x M

23. Practice
Calculate the pH and concentration of oxalate ion [C2O42 -], in a 0.020M solution of oxalic acid (H2C2O4). Ka 1= 5.9x10- 2, Ka 2= 6.4x10- 5
Ans: pH= 1.80, [C2O42 -]= 6.4x10- 5

24. Weak Bases B(aq) + H2O(l) ↔ HB+(aq) + OH-(aq)
NH3(aq) + H2O(l) ↔ NH4+(aq) + OH-(aq)
Kb = base-dissociation constant
What is the Kb expression for NH3?
Calculate the [OH-] in a 0.15M solution of NH3. Kb= 1.8x Ans: 1.6x10- 3M

25. Weak Bases

26. Ka and Kb Ka x Kb = Kw The larger Ka the lower the Kb
Calculate the Ka for HF if the Kb of F- is 1.5x

27. pH of Salt Solutions
An anion that is the conjugate base of a strong acid will not affect the pH of a solution. ex. Br-
An anion that is the conjugate base of a weak acid will cause an increase in pH. ex. CN-
A cation that is the conjugate acid of a weak base will cause a decrease in pH. ex. NH4+
With the exception of ions of group 1A and heavier members of group 2A, metal ions will cause a decrease in pH.
When a solution contains both the conjugate base of a weak acid and the conjugate acid of a weak base, the ion with the largest ionization constant will have the greatest affect on pH.

28. Practice
Predict whether the salt Na2HPO4 will form an acidic or basic solution on dissolving in water.
Predict whether the K2HC6H5O7 will form an acidic or basic solution in water. (look at citric acid)
Ans: basic, acidic

29. Acid Strength
Molecules only donate protons if the H—X bond is polar, where the anion X is more electronegative than the H.
As you move from left to right X becomes more electronegative and acid strength increases.
CH4 < NH3 << H2O < HF
Strong H—X bonds are harder to break than weak ones.
The strength of the H—X bond decreases as the size of X increases.
HF versus HCl, H2S versus H2O
The more stable the resulting anion, X-, the stronger the acid.

30. Binary Acid Strength

31. Oxyacid Strength Acids with OH groups and additional oxygen
atoms bound to a central
atom are called oxyacids.
H2SO4
For oxyacids that have the same number of OH groups and the same number of O atoms, acid strength increases with increasing electronegativity of the central atom.
For oxyacids that have the same central atom, acid strength increases as the number of oxygen atoms attached increases. ex. Chloric acids.

32. Lewis Acids and Bases Lewis acids – electron pair acceptor
Increases the types of compounds that we can consider acids
Lewis bases – electron pair donor
All bases that are Brønsted-Lowry bases are Lewis bases.
NH3 + BF3 → NH3BF3
Metal ions reacting with water act as Lewis acids

33. Homework