Lecture 3: Acids, Bases and pH

1. An acid is defined as a proton donor and a base as a proton acceptor2. The following equation describes the solution of acid in water:

HA(aq) + H2O(l) ⇋ H3O+(aq) + A-(aq)

3. However, a shorter way to write the above is:

HA ⇋ H+ + A-

4. HA is the acid and A- is the conjugate base

5. The equations in 2 and 3 describe the dissociation of an acid in solution

NB- you must be able to write equations for the dissociation of any given acid (see Table 2.6 in your text book) and to identify the acid and conjuate base.

6. This process is also a proton transfer reaction, since a proton (H+) is transferred from the acid to H2O (see equation in 2)

7. The equilibrium constant for the equation in 3 is called the acid dissociation constant (Ka)

Ka = [H + ] [A - ] [HA]

8. Ka is an indication of the extent of dissociation of an acid at a certain temperature. The higher the Ka value, the more complete the dissociation

9. Strong acids are more dissociated in solution than weak acids, therefore Ka gives an indication of the strength of an acid

10. The Ka of a strong acid (e.g HCl or H2SO4) approaches infinity. In Biochemistry, we work mostly with weak acids (1.8 x 10-16 < Ka < 55.5)

11. Water also dissociates in solution. This process is referred to as the self-ionization of water

H2O ⇋ H+ + OH-

12. The acid dissiciation constant for H2O is given by

Ka = [H + ] [OH - ] [H2O]

Ka and [H2O] are constants, therefore [H+][OH-] must be constant too.

we define Kw = [H+][OH-]

Kw is called the ion product of water

Its value is 1 x 10-14 at room temperature, regardless of whether the solution is neutral, acidic or basic

13. pH is a convenient way to express [H+]

pH = -log[H+]

14. Similarly, pKa is a convenient way to express Ka values.

pKa = -log Ka

15. A monoprotic acid can donate 1 proton (e.g. HCl) A diprotic acid can donate 2 protons (e.g. HOOC-COOH) A triprotic acid can donate 3 protons (e.g. H3PO4) A polyprotic acid can donate 3 or more protons

16. Each dissociation step for a diprotic (or polyprotic) acid has a unique Ka value

HOOC-COOH ⇋ H+ + HOOC-COO- Ka1 = [H + ] [HOOC-COO - ] [HOOC-COOH]

HOOC-COO- ⇋ H+ + -OOC-COO-

Ka2 = [H + ] [ - OOC-COO - ] [HOOC-COO-]

17. The Henderson-Hasselbalch equation can be derived as follows (NB: know how to derive!!):

18. The Henderson-Hasselbalch equation:

a. relates the pH of a solution of a weak acid to the relative amounts of acid and conjugate base in the solution (i.e. the extent of dissociation)

b. can be used to determine the pH value of a weak acid solution given the ratio of acid and conjugate base components

c. can be used to determine the ratio of acid and base components that will yield a solution of desired pH.

19. Titration: an experiment in which measured amounts of acid (or base) are added to measured amounts of base (or acid)

NB – When you are presented with a titration curve of a weak acid, you must be able to label the axes and identify the inflection point and equivalence point (see Fig. 2.13 in text book; Lecture slide 23; class notes points 21 and 22)

20. Equivalent: The quantity of base that reacts with 1 mol of [H+] or the amount of acid that yields 1 mol of [H+]

21. Equivalence point: the point in an acid-base titration at which enough base has been added to exactly neutralize the acid (or vice versa)

a. The equivalence point is reached after the addition of 1 equavalent (eq) of base for every equivalent of acid

b. also called the titration end-point

22. Inflection point: the point in an acid-base titration at which enough base has been added to neutralize half of the acid (or vice versa)

a. The inflection point is reached after the addition of 0.5 equivalent (eq) of base for every equivalent of acid.

b. also called the titration mid-pointc. at inflection point: pKa = pH

23. From the Henderson-Hasselbalch equation, it follows that:

a. If pH < pKa, then substance will be protonated (H+ on)b. If pH > pKa, then substance will be deprotonated (H+ off)