1. Buffers
AP Chemistry

2. Common Ion Effect
Adding a common ion to an equilibrium system will shift the equilibrium forward or reverse.
For a weak acid equilibrium this can affect the pH of the solution.
For example: Adding NaF to a solution of HF.
HF H+ + F-
Adding F- ions will shift the equilibrium to reactants and produce more HF and reduce [H+] ions in the process.
This makes the pH increase.

3. Buffers
Buffers are solutions in which the pH remains relatively constant, even when small amounts of acid or base are added
Contain a weak acid (HA) and its salt(NaA); H2CO3 & NaHCO3
Or a weak base (B) and its (BHCl); (NH3 and NH4Cl)

4. Buffers
A buffer system is better able to resist changes in pH than pure water
Since it is a pair of chemicals:
one chemical neutralizes any acid added, while the other chemical would neutralize any additional base
AND, they produce each other in the process!!!

5. How a Buffer Works Consider the following buffer system
HCO H H2CO3
If you add more H+ this buffer system it will react with the conjugate base HCO3- to produce more H2CO3.
If you add base, OH-, it will grab an H+ from H2CO3 to produce more HCO3- ion as follows
H2CO3 + OH H2O + HCO3-

6. Buffer Capacity
The buffer capacity is the amount of acid or base that can be added before a significant change in pH
This depends on the amounts of HA and A- present in the buffer
Most efficient buffer is when
[A-]
[HA]
= 1

7. Henderson-Hasselbach Equation
Derived from the equilibrium expression of a weak acid and the pH equation.
This equation allows you to determine the pH of a buffer solution
pKa = -log Ka
[A-]
[HA]
pH = pKa + log

8. The Common Buffer Problems
1. Compute pH of a buffer given the actual concentrations of the conjugate acid and conjugate base.
The pH is easily calculated with:
[A-]
[HA]
pH = pKa + log

9. Example: Determine the pH in which 1. 00 mole of H2CO3 (Ka = 4
Example: Determine the pH in which 1.00 mole of H2CO3 (Ka = 4.2 x 10-7) and 1.00 mole NaHCO3 dissolved in enough water to form 1.00 Liters of solution.
[A-]
[HA]
pH = pKa + log
pH = -log (4.2 x 10-7) + log (1.00M)/(1.00M)
pH = pKa = 6.4
** pKa of a weak acid can help determine pH of buffer you will make if it is mixed in a 1:1 mole ratio!

10. 2. Make a buffer problem [HCO3-] pH = pKa + log [H2CO3]
How many moles of NaHCO3 should be added to 1 liter of 0.100M H2CO3 (Ka = 4.2 x 10-7 ) to prepare a buffer with a pH of 7.00?
[HCO3-]
[H2CO3]
pH = pKa + log
7.00 = -log(4.2 x 10-7) + log [HCO3-]
(0.100)
0.60 = log[HCO3-]
0.001
[HCO3-] = moles should be added!
= [HCO3-]
0.100

11. 3. The pH shift problem:
If you add 2.00 mL of M HCl to 100 mL of a buffer consisting of M HA and M NaA, what will be the change in pH?
Ka of HA is 1.5x pKa = 4.82.
The initial pH is:

12. If we add H+ to the equilibrium system it will react with the strong conjugate base ‘A-, tor form more HA.
H+ + A-  HA
Set up an ICE chart to determine how the mole values will change using stoichiometry!
mmoles H+
mmoles A-
mmoles HA
Volume, mL
Initial
0.200
20.0
10.0
100+2 mL
Final
19.8
10.2
102 mL
Molarities
19.8/102
10.2/102