1. Buffers

2. What Are They? Solutions that resist changes in pH with addition of small amounts of acid or base Require two species: an acid to react with OH - and a base to react with H + It is necessary that the acidic and basic species not consume one another through a neutralization reaction Usually acid-base conjugate pair Can be prepared by mixing a weak acid or base with a salt of that acid or base

3. How Does It Work? Buffer Solution: HC 2 H 3 O 2 and NaC 2 H 3 O 2 HC 2 H 3 O 2  H + + C 2 H 3 O 2 - Small amount of acid added  shifts left  limits H +  thus limits pH If small amount of base added  H + reacts with OH -  shifts equilibrium right to make more H +  thus resists pH change

5. General Form HX  H + + X - K a = [H + ][X - ] [HX] [H + ] = K a [HX] [X - ]

6. pH is determined by two factors: value of K a and ratio of concentrations of conjugate acid-base pair If the amounts of HX and X - are large compared to the amount of acid or base added, the pH doesn’t change much Most effective where [HX] is about equal to the [X - ] ; pH is about equal to the pK a

7. Equation Henderson-Hasselbalch Equation pH = pK a + log [base] [acid]

8. Example Calculate the pH of a solution formed from.10M formic acid and.20M potassium formate.

9. Example What must be the concentration of NH 4 Cl in a.10M solution of NH 3 if the pH is to be 9.00?

10. Example Calculate the concentration of sodium formate that must be present in a.10M solution of formic acid to produce a pH of 3.80.