1. Acid-Base Titrations

2. Introduction to Acids and Bases
Chapter 8

3. There are three (3) definitions of acids and bases: 1. Arrhenius, 2
There are three (3) definitions of acids and bases: 1. Arrhenius, 2. Bronsted-Lowry and 3. Lewis.
Name
Acids
Base
Definition
Example
Arrhenius
An H-containing compound that increases H+ concentration in water
HCl
HNO3
HBr HF
Acetic acid
An OH-containing compound that increases OH- concentration in water
KOH
NaOH
Bronsted-Lowry
Proton donor
Terminal alkynes
Proton acceptor
NH3
Lewis
Electron pair acceptor
Fe3+ and other transition metal cations
Electron pair donor

4. There are three (3) definitions of acids and bases: 1. Arrhenius, 2
There are three (3) definitions of acids and bases: 1. Arrhenius, 2. Bronsted-Lowry and 3. Lewis.
base
acid
acid
base
conjugate acid
conjugate base
base
acid
Conjugate acid-base pairs differ by only one proton.

5. There are three (3) definitions of acids and bases: 1. Arrhenius, 2
There are three (3) definitions of acids and bases: 1. Arrhenius, 2. Bronsted-Lowry and 3. Lewis.
Arrhenius Bases:
LiOH, Mg(OH)2
Arrhenius Acids:
HCl, H2SO4
Bronsted Acids:
H2O, HS-, H2PO4-
Bronsted Bases:
H2O, NH3, HS-
Lewis Acids:
Any species that accepts electron pairs. (H+ CO2, Mn+)
Lewis Bases:
Any species with a lone pair

6. Acids and Bases undergo neutralization reactions.
NaOH + HCl  NaCl + H2O
Na+ + OH– + H+ + Cl–  Na+ + Cl– + H2O
H+ + OH–  H2O

7. Acids and Base strengths are defined by how much H3O+ or OH- is produced by a given concentration
Strong Acid
Weak Acid

8. For a monoprotic acid HA
Acids and Base strengths are defined by how much H3O+ or OH- is produced by a given concentration
Ionized acid concentration at equilibrium
Initial concentration of acid
x 100%
percent ionization =
For a monoprotic acid HA
[H+]eq
Percent ionization =
x 100%
[HA]0 = initial concentration
[HA]0

9. Weak acids and bases are in equilibrium with their original species.
CH3COOH + H2O H3O+ +CH3COO–
Kc << 1

10. Weak acids and bases are in equilibrium with their original species.
NH3 + H2O NH4+ +OH–
Kc << 1

11. Ka and Kb tells us something about the relative acid/base strengths

12. Ka and Kb tells us something about the relative acid/base strengths

13. Water can act both as an acid and a base (AUTOIONIZATION)
The ion-product constant (Kw) is the product of the concentration of H3O+ and OH– ions at a particular temperature
At 250C
Kw = [H3O+][OH-] = 1.0 x 10-14

14. The p-scale conveniently handles a wide range of concentrations

15. Solving weak acid ionization problems:
Identify the major species that can affect the pH.
In most cases, you can ignore the autoionization of water. [H3O+] from water is negligible in comparison to [H3O+] from the weak acid.
Ignore [OH-] because it is determined by [H+].
Use ICE to express the equilibrium concentrations in terms of single unknown x.
Write Ka in terms of equilibrium concentrations. Solve for x by the approximation method. If approximation is not valid, solve for x exactly.
Calculate concentrations of all species and/or pH of the solution.

16. Exercises
What is the pH of a 0.5 M HF solution (at 25°C, Ka = 7.1 x 10-4)?
What is the pH of a 0.05 M HF solution?
What is the pH of a M monoprotic acid whose Ka is 5.7 x 10-4?

17. Parameter An NaOH solution A 0.30 M HNO3 pH 9.52 pOH [H3O+] [OH-]
Complete the following table:
Parameter
An NaOH solution
A 0.30 M HNO3 pH
9.52
pOH
[H3O+]
[OH-]

18. Parameter An NaOH solution A 0.30 M HNO3 pH 9.52 0.52 pOH 4.48 13.48
Complete the following table:
Parameter
An NaOH solution
A 0.30 M HNO3 pH
9.52
0.52
pOH
4.48
13.48
[H3O+]
3.0 x 10-10M
0.30 M
[OH-]
3.3 x 10-5 M
3.3 x M

19. Exercises What is the pH, [H3O+], [OH-] of 7.52 x 10-4 M CsOH?
What is the pOH, [H3O+], [OH-] of 1.59 x10-3 M HClO4?
What is the [H3O+], [OH-] and pOH in a solution with a pH of 2.77
What is the [H3O+], [OH-] and pH in a solution with a pOH of 11.27

20. Exercises
Chloroacetic acid has a pKa of What are [H3O+], pH, [ClCH2COOH], [ClCH2COO-] in 1.05 M [ClCH2COOH]
A M of weak acid is 12.5% dissociated. Calculate [H3O+], pH, [OH-], pOH of solution. Calculate Ka of acid

21. Buffers
Chapter 9

22. Relative Strengths and Reaction Direction
H2S + NH HS– + NH4+
Stronger
acid
Stronger
base
Weaker
base
Weaker
acid
Net direction
>
Ka of H2S = 9 x10-8
Ka of NH4+ = 5.68 x10-10
>
Kb of HS– = 1.11 x10-7
Kb of NH3 = 1.76 x10-5

23. HNO3 + H2O NO3– + H3O+ HF + H2O F– + H3O+ Net direction Net direction
Stronger
acid
Stronger
base
Weaker
base
Weaker
acid
Net direction
HF + H2O F– + H3O+
Weaker
acid
Weaker
base
Stronger
base
Stronger
acid
Net direction
Ka of HF = 6.8 x10-4
[Products] << [Reactants]
Ka << 1

24. A few observations
The conjugate base of a very strong acid is a very weak base
The conjugate base of a weak acid is a strong base
The conjugate acid of a strong base is a weak acid
The conjugate acid of a weak base is a strong acid

25. A few observations

26. Buffers solutions are solutions that resist changes in pH

27. HA/ A– NH4Cl/NH3 H3PO4/NaH2PO4 NH4SH/Na2S HCOOH/HCOOK HBr/KBr
Buffers contain appreciable amounts of a weak acid and its conjugate base
HA/ A–
NH4Cl/NH3
H3PO4/NaH2PO4
NH4SH/Na2S
HCOOH/HCOOK
HBr/KBr
H3IO3/Li2HIO3
NaOH/Na2O

28. HA / A– What do we mean by appreciable?
Buffers contain appreciable amounts of a weak acid and its conjugate base
What do we mean by appreciable?
A 50-mL solution of 0.25 M Acetic acid.
A 50-mL solution of 0.25 M Sodium Acetate
A solution containing M Acetic acid and M Acetate
HA / A–
** If we take the ratio of base to acid or acid to base, it should be within 10% of each other
+ base
+ acid

29. Buffers work because there are weak acids and weak bases present to counter-act small amounts of acid/bases.

30. Henderson-Hasselbach Equation:
The effectivity (and pH) of the buffer is dependent on the ratio between the weak acid and its conjugate base
Henderson-Hasselbach Equation:

31. BUFFER RANGE A 0.25 M Acetic acid buffer with pH 4.74
The effectivity (and pH) of the buffer is dependent on the ratio between the weak acid and its conjugate base
A 0.25 M Acetic acid buffer with pH 4.74
A 0.25 M Acetic acid buffer with pH 5.10
A 0.25 M Acetic acid buffer with pH 4.40
** If we take the ratio of base to acid or acid to base, it should be within 10% of each other
BUFFER RANGE

32. The closer the pH of the buffer to the pKa the better
The effectivity (and pH) of the buffer is dependent on the ratio between the weak acid and its conjugate base
A buffer for pH 10.00
A buffer for pH 4.00
A buffer for pH 7.00
The closer the pH of the buffer to the pKa the better

33. A 1.00 M Acetic acid buffer with pH 4.74
The effectivity (and pH) of the buffer is also dependent on the total amount of weak acid and conjugate base.
A 1.00 M Acetic acid buffer with pH 4.74
A 0.30 M Acetic acid buffer with pH 4.74
A 0.10 M Acetic acid buffer with pH 4.74
A M Acetic acid buffer with pH 4.74

34. + appreciable amounts (High Concentration)…
Buffer Capacity is the measure of the ability of a buffer to resist pH Changes
GOOD BUFFER RANGE
+ appreciable amounts (High Concentration)…

35. Buffers work through a phenomenon known as the common ion effect
What is the common ion effect?
This effect occurs when a reactant containing a given ion is added to an equilibrium mixture that already contains that ion, and the position of equilibrium shifts away from forming more of it

36. Preparation of Buffers
Choose the conjugate acid-base pair
Calculate the ratio of the buffer component concentrations
Determine the buffer concentration
Mix the solution and adjust pH

37. Buffers can be prepared using a weak acid and the salt of its conjugate base (or a weak base and the salt of its conjugate acid).
Example 1
Preparing a pH carbonate buffer. How many grams of Na2CO3 must one add to 1.5 L of freshly prepared 0.20 M NaHCO3 to make the buffer? Ka of HCO3- is 4.7 x 10-11
Example 2
Prepare a 50-mL of 0.12 M Acetic acid buffer with equal concentrations of acetic acid and acetate from 3.00 M acetic acid stock solution and sodium acetate salt

38. Buffers can also be prepared by using a weak acid (or weak base) then add a strong base (or acid) to desired pH.
5.0 g of CH3COONa is dissolved in mL of water. How many mL of 0.50 MHCl should be added to form a buffer with pH 4.90? Will diluting the final mixture to 500 mL affect the pH of the buffer? What is affected by dilution?

39. Acid-Base Titrations
Chapter 10

40. In an acid-base reaction, the key parameter that changes in the system is pH.
Example:
A mL sample of M HCl solution was titrated with M NaOH solution.
Calculate the pH when the following volume of the NaOH is added.
0.00
39.00
10.00
40.00
20.00
41.00
30.00
45.00
35.00
50.00

41. In an acid-base reaction, pH is monitored through colored indicators.

42. change occurs over ~2 pH units
Acid-base indicators are usually weak acids (HIn) which have different color than its conjugate base (In-).
basic
change occurs over ~2 pH units
acidic

43. Acid-base indicators are usually weak acids (HIn) which have different color than its conjugate base (In-).

44. TITRATION OF A STRONG ACID BY A STRONG BASE
40 mL of M HCl

45. TITRATION OF A WEAK ACID BY A STRONG BASE
REGION 1: Before addition (weak acid)
REGION 2: Before Equivalence point (buffer)
REGION 3: At Equivalence point (weak base)
REGION 4: After Equivalence point (strong base)

46. TITRATION OF A WEAK ACID BY A STRONG BASE
EXAMPLE in Book: Titration of mL of M MES (pKa = 6.27) with M NaOH
REGION 1: Before addition (weak acid)
REGION 2: Before Equivalence point (buffer)
REGION 3: At Equivalence point (weak base)
REGION 4: After Equivalence point (strong base)

47. TITRATION OF A WEAK BASE BY A STRONG ACID
REGION 1: Before addition (weak base)
REGION 2: Before Equivalence point (buffer)
REGION 3: At Equivalence point (weak acid)
REGION 4: After Equivalence point (strong acid)

48. TITRATION OF A WEAK BASE BY A STRONG ACID
Example in Book: Titration of mL of M pyridine (Kb = 1.6 x 10-9) with M HCl.
REGION 1: Before addition (weak base)
REGION 2: Before Equivalence point (buffer)
REGION 3: At Equivalence point (weak acid)
REGION 4: After Equivalence point (strong acid)

49. TITRATION OF A POLYPROTIC SYSTEMS

50. TITRATION OF A POLYPROTIC SYSTEMS