1. Acids and Bases
Chapter 13

2. Some Properties of Acids
Produce H+ (as H3O+) ions in water (the hydronium ion is a hydrogen ion attached to a water molecule)
Taste sour
Corrode metals
Electrolytes
React with bases to form a salt and water
pH is less than 7
Turns blue litmus paper to red “Blue to Red A-CID”

3. Some Properties of Bases
Produce OH- ions in water
Taste bitter, chalky
Are electrolytes
Feel soapy, slippery
React with acids to form salts and water
pH greater than 7
Turns red litmus paper to blue “Basic Blue”

4. Acid Nomenclature Review
No Oxygen
w/Oxygen
An easy way to remember which goes with which…
“In the cafeteria, you ATE something ICky”

5. Acid/Base definitions Definition 1: Arrhenius
Arrhenius acid is a substance that produces H+ (H3O+) in water
Arrhenius base is a substance that produces OH- in water

6. Acid/Base Definitions
Definition #2: Brønsted – Lowry
Acids – proton donor
Bases – proton acceptor
A “proton” is really just a hydrogen atom that has lost it’s electron!

7. A Brønsted-Lowry acid is a proton donor
A Brønsted-Lowry base is a proton acceptor
conjugate acid
conjugate base
base
acid

8. ACID-BASE THEORIES
The Brønsted definition means NH3 is a BASE in water — and water is itself an ACID

9. Conjugate Pairs

10. Learning Check!
Label the acid, base, conjugate acid, and conjugate base in each reaction:
HCl + OH-    Cl- + H2O
Acid
Base
Conj.Base
Conj.Acid
H2O + H2SO4    HSO4- + H3O+
Conj.Base
Conj.Acid
Base
Acid

11. Acids & Base Definitions
Definition #3 – Lewis
Lewis acid - a substance that accepts an electron pair
Lewis base - a substance that donates an electron pair

12. Lewis Acids & Bases
Formation of hydronium ion is also an excellent example.
Electron pair of the new O-H bond originates on the Lewis base.

13. Lewis Acid/Base Reaction

14. The pH scale is a way of expressing the strength of acids and bases
The pH scale is a way of expressing the strength of acids and bases. Instead of using very small numbers, we just use the NEGATIVE power of 10 on the Molarity of the H+ (or OH-) ion. Under 7 = acid 7 = neutral Over 7 = base

15. (Remember that the [ ] mean Molarity)
Calculating the pH
pH = - log [H+]
(Remember that the [ ] mean Molarity)
Example: If [H+] = 1 X pH = - log 1 X 10-10
pH = - (- 10)
pH = 10
Example: If [H+] = 1.8 X 10-5 pH = - log 1.8 X 10-5
pH = - (- 4.74)
pH = 4.74

16. Try These! pH = - log [H+] pH = - log 0.15 pH = - (- 0.82) pH = 0.82
pH = - log 3 X 10-7
pH = - (- 6.52)
pH = 6.52
Find the pH of these:
A 0.15 M solution of Hydrochloric acid
2) A 3.00 X 10-7 M solution of Nitric acid

17. pH calculations – Solving for H+
If the pH of Coke is 3.12, [H+] = ???
Because pH = - log [H+] then
- pH = log [H+]
Take antilog (10x) of both sides and get
10-pH = [H+]
[H+] = = 7.6 x 10-4 M
*** to find antilog on your calculator, look for “Shift” or “2nd function” and then the log button

18. More About Water Equilibrium constant for water = Kw
H2O can function as both an ACID and a BASE.
In pure water there can be AUTOIONIZATION
Equilibrium constant for water = Kw
Kw = [H3O+] [OH-] = x at 25 oC

19. More About Water and so [H3O+] = [OH-] = 1.00 x 10-7 M Autoionization
Kw = [H3O+] [OH-] = x at 25 oC
In a neutral solution [H3O+] = [OH-]
and so [H3O+] = [OH-] = 1.00 x 10-7 M

20. pOH Since acids and bases are opposites, pH and pOH are opposites!
pOH does not really exist, but it is useful for changing bases to pH.
pOH looks at the perspective of a base
pOH = - log [OH-]
Since pH and pOH are on opposite ends,
pH + pOH = 14

21. pH
[H+]
[OH-]
pOH

22. [H3O+], [OH-] and pH What is the pH of the 0.0010 M NaOH solution?
[OH-] = (or 1.0 X 10-3 M)
pOH = - log
pOH = 3
pH = 14 – 3 = 11
OR Kw = [H3O+] [OH-]
[H3O+] = 1.0 x M
pH = - log (1.0 x 10-11) = 11.00

23. 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

24. Strong and Weak Acids/Bases
The strength of an acid (or base) is determined by the amount of IONIZATION.
HNO3, HCl, HBr, HI, H2SO4 and HClO4 are the strong acids.

25. Strong and Weak Acids/Bases
Generally divide acids and bases into STRONG or WEAK ones.
STRONG ACID: HNO3 (aq) + H2O (l)  H3O+ (aq) NO3- (aq)
HNO3 is about 100% dissociated in water.

26. Strong and Weak Acids/Bases
Weak acids are much less than 100% ionized in water.
*One of the best known is acetic acid = CH3CO2H

27. Strong and Weak Acids/Bases
Strong Base: 100% dissociated in water.
NaOH (aq)  Na+ (aq) + OH- (aq)
Other common strong bases include KOH and Ca(OH)2.
CaO (lime) + H2O -->
Ca(OH)2 (slaked lime)
CaO
Strong bases are the group I hydroxides
Calcium, strontium, and barium hydroxides are strong, but only soluble in water to 0.01 M

28. Strong and Weak Acids/Bases
Weak base: less than 100% ionized in water
One of the best known weak bases is ammonia
NH3 (aq) + H2O (l) ↔ NH4+ (aq) + OH- (aq)

29. Weak Bases

30. Equilibria Involving Weak Acids and Bases
Consider acetic acid, HC2H3O2 (HOAc)
HC2H3O2 + H2O ↔ H3O C2H3O2 -
Acid Conj. base
(K is designated Ka for ACID)
K gives the ratio of ions (split up) to molecules (don’t split up)

31. Ionization Constants for Acids/Bases
Conjugate
Bases
Increase strength
Increase strength

32. Equilibrium Constants for Weak Acids
Weak acid has Ka < 1
Leads to small [H3O+] and a pH of 2 - 7

33. Equilibria Involving A Weak Acid
You have 1.00 M HOAc. Calc. the equilibrium concs. of HOAc, H3O+, OAc-, and the pH.
Step 1. Define equilibrium concs. in ICE table.
[HOAc] [H3O+] [OAc-]
initial
change
equilib
-x +x +x
1.00-x x x

34. Equilibria Involving A Weak Acid
You have 1.00 M HOAc. Calc. the equilibrium concs. of HOAc, H3O+, OAc-, and the pH.
Step 2. Write Ka expression
This is a quadratic. Solve using quadratic formula.
or you can make an approximation if x is very small! (Rule of thumb: 10-5 or smaller is ok)

35. Equilibria Involving A Weak Acid
You have 1.00 M HOAc. Calc. the equilibrium concs. of HOAc, H3O+, OAc-, and the pH.
Step 3. Solve Ka expression
First assume x is very small because Ka is so small.
Now we can more easily solve this approximate expression.

36. Equilibria Involving A Weak Acid
You have 1.00 M HOAc. Calc. the equilibrium concs. of HOAc, H3O+, OAc-, and the pH.
Step 3. Solve Ka approximate expression
x = [H3O+] = [OAc-] = 4.2 x 10-3 M
pH = - log [H3O+] = -log (4.2 x 10-3) = 2.37

37. Equilibria Involving A Weak Acid
Calculate the pH of a M solution of formic acid, HCO2H.
HCO2H H2O ↔ HCO H3O+
Ka = 1.8 x 10-4
Approximate solution
[H3O+] = 4.2 x 10-4 M, pH = 3.37
Exact Solution
[H3O+] = [HCO2-] = 3.4 x 10-4 M
[HCO2H] = x 10-4 = M
pH = 3.47

38. Equilibrium Constants for Weak Bases
Weak base has Kb < 1
Leads to small [OH-] and a pH of

39. Relation of Ka, Kb, [H3O+] and pH

40. Equilibria Involving A Weak Base
You have M NH3. Calc. the pH.
NH3 + H2O ↔ NH OH-
Kb = 1.8 x 10-5
Step 1. Define equilibrium concs. in ICE table
[NH3] [NH4+] [OH-]
initial
change
equilib
-x +x +x
x x x

41. Equilibria Involving A Weak Base
You have M NH3. Calc. the pH.
NH3 + H2O  NH OH-
Kb = 1.8 x 10-5
Step 2. Solve the equilibrium expression
Assume x is small, so
x = [OH-] = [NH4+] = 4.2 x 10-4 M
and [NH3] = x 10-4 ≈ M
The approximation is valid !

42. Equilibria Involving A Weak Base
You have M NH3. Calc. the pH.
NH3 + H2O  NH OH-
Kb = 1.8 x 10-5
Step 3. Calculate pH
[OH-] = 4.2 x 10-4 M
so pOH = - log [OH-] = 3.37
Because pH + pOH = 14,
pH = 10.63

43. Types of Acid/Base Reactions: Summary

44. Weak Bases are weak electrolytes
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.

46. Strong Acid
Weak Acid

47. 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

48. 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

49. Molecular Structure and Acid Strength
H X
H+ + X-
Bond strength
Polarity
The stronger the bond
The weaker the acid
HF << HCl < HBr < HI

50. Molecular Structure and Acid StrengthZ O H O-
+ H+ d- d+
The O-H bond will be more polar and easier to break if:
Z is very electronegative or
Z is in a high oxidation state

51. Molecular Structure and Acid Strength
1. Oxoacids having different central atoms (Z) that are from the same group and that have the same oxidation number.
Acid strength increases with increasing electronegativity of Z
H O Cl O O •
H O Br O O •
Cl is more electronegative than Br
HClO3 > HBrO3
15.9

52. Molecular Structure and Acid Strength
2. Oxoacids having the same central atom (Z) but different numbers of attached groups.
Acid strength increases as the oxidation number of Z increases.
HClO4 > HClO3 > HClO2 > HClO

53. Acid-Base Properties of Salts
Neutral Solutions:
Salts containing an alkali metal or alkaline earth metal ion (except Be2+) and the conjugate base of a strong acid (e.g. Cl-, Br-, and NO3-).
NaCl (s) Na+ (aq) + Cl- (aq)
H2O
Basic Solutions:
Salts derived from a strong base and a weak acid.
NaCH3COO (s) Na+ (aq) + CH3COO- (aq)
H2O
CH3COO- (aq) + H2O (l) CH3COOH (aq) + OH- (aq)

54. Acid-Base Properties of Salts
Acid Solutions:
Salts derived from a strong acid and a weak base.
NH4Cl (s) NH4+ (aq) + Cl- (aq)
H2O
NH4+ (aq) NH3 (aq) + H+ (aq)
Salts with small, highly charged metal cations (e.g. Al3+, Cr3+, and Be2+) and the conjugate base of a strong acid.
Al(H2O)6 (aq) Al(OH)(H2O)5 (aq) + H+ (aq) 3+ 2+

55. Acid-Base Properties of Salts
Solutions in which both the cation and the anion hydrolyze:
Kb for the anion > Ka for the cation, solution will be basic
Kb for the anion < Ka for the cation, solution will be acidic
Kb for the anion  Ka for the cation, solution will be neutral