1. Recap – Last Lecture An acid is a proton donor A base is a proton acceptor A conjugate pair differ by H + Strong A/B is completely dissociated Weak A/B is in equilibrium The smaller pK a, the stronger the acid and the weaker its conjugate base. 1

2. 2 Predicting Reactions A reaction between a strong acid and a strong base will go to completion. A strong acid – weak base or weak acid – strong base reaction will go to completion. eg CH 3 COOH + OH – → CH 3 COO – + H 2 O For a reaction between a weak acid and a weak base, a comparison of pK a values enables us to determine whether the reaction will occur.

3. 3 Predicting Reactions Example: Will hydrogencarbonate (HCO 3 – ) react with phenol (C 6 H 5 OH)? C 6 H 5 OH + HCO 3 – ⇌ C 6 H 5 O – + H 2 CO 3 pK a =10.0pK a =6.35 The answer is ‘no’. The equilibrium lies to the left. But hydorgencarbonate will react with acetic acid: CH 3 COOH + HCO 3 – CH 3 COO – + H 2 CO 3 pK a =4.7pK a =6.35

4. 4 Buffer A buffer is a solution composed of moderate quantities of both members of a conjugate acid-base pair (e.g. CH 3 COOH and CH 3 COO – (Na + )). It maintains a solution at approximately constant pH even when small quantities of H + or OH – are added. Conjugate pairpKapKa Optimum pH of buffer CH 3 COOH / CH 3 COO - 4.76 H 2 PO 4 - / HPO 4 2- 7.20 NH 4 + / NH 3 9.26

5. 5 Relationship between pK a and pH The equilibrium between a conjugate acid – base pair is affected by pH. HA ⇌ H + + A – At high [H + ], (low pH), the equilibrium is towards the left and visa versa. Comparison of pH with pK a of the weak acid/base system indicates in which direction the equilibrium lies. If the pH is on the ‘acid side’ of the pK a, the conjugate acid will predominate. If the pH is on the ‘base side’ of the pK a, the conjugate base will predominate.

6. 6 Example Which of CH 3 COOH / CH 3 COO – will dominate at a physiological pH = 7.4 given CH 3 COOH pK a = 4.76? CH 3 COOH(aq) ⇌ CH 3 COO – (aq) + H + (aq) Answer: pH > pK a ie pH is on the ‘base side’ of the pK a so the conjugate base will dominate: CH 3 COO –

7. 7 Example A more complicated diagram results for a polyprotic acid, eg H 3 PO 4. Note pK a1 < pK a2 < pK a3 pK a (H 3 PO 4 ) = 2.2 pK a (H 2 PO 4 - ) = 7.2 pK a (HPO 4 2- ) = 12.3

8. 8 Amino Acids An amino acid contains (at least) two groups that are acid base active – a carboxylic acid and an amine. Each group has a pK a associated with it. e.g. glycine At physiological pH, this will exist in an ionic form:

9. 9 Applications – CO 2 in blood Buffers are an important part of living systems. Human blood has a normal pH range of 7.35-7.45. Any deviation from this can disrupt cell membranes, proteins and enzyme activity. The major buffer solution that controls blood pH is CO 2 (g) + H 2 O(l) ⇌ H 2 CO 3 (aq) ⇌ H + (aq) + HCO 3 – (aq) … (1) Notes: Although carbonic acid (H 2 CO 3 )is diprotic, only the first ionization is important. CO 2 is a gas which provides the body with a mechanism to adjust the equilibrium. Removal of CO 2 by exhalation shifts the equilibrium to the left, consuming H + ions. Buffer solution in blood plasma has [HCO 3 – ] ~ 0.024 M and [H 2 CO 3 ] ~ 0.0012 M (20/1 ratio). As a consequence the buffer solution has a high capacity to neutralize additional acid, but low capacity to neutralize additional base. Principle organs that regulate blood pH are the lungs and kidneys with input from brain sensors. When [CO 2 ] rises, equilibrium shifts to the right which increases [H + ] which, in turn, triggers receptors that increase breathing rate and elimination of CO 2 to restore pH. Kidneys absorb & release H + and HCO 3 – ; much excess acid leaves the body in urine (pH 5-7).

10. 10 Applications – O 2 in blood Regulation of pH in the blood relates directly to effectiveness of oxygen transport. The blood transports oxygen using haemoglobin (Hb) which reversibly binds H + and O 2 in a competitive equilibrium: HbH + + O 2 ⇌ HbO 2 + H + ………(2) Notes: When the blood reaches the tissue where the [O 2 ] is low, the equilibrium is shifted to the left to release O 2. An increase in [H + ] and an increase in temperature also shifts the equilibrium to the left (ΔH < 0, exothermic). During exertion several factors work to ensure delivery O 2 to active tissue. As O 2 is consumed the equilibrium 2 shifts to the left. Exertion raises body temp which also shifts equilibrium 2 to the left. Increased metabolic production of CO 2 shifts equilibrium 1 to the right increasing [H + ] which also shifts equilibrium 2 to the left. Other acids such as lactic acid are produced when tissue is starved of O 2 which further shifts equilibrium 2 to the left. (This extra production of acids may produce cramps.) Increased [H + ] stimulates increase rate of breathing, which provides more O 2 and eliminates CO 2.

11. By the end of this lecture, you should: −predict the outcome of an acid - base reaction. −understand the function and composition of a buffer. −recognise the relationship between pH, pK a and the form of the conjugate pair present. −be able to predict the charge state of an amino acid at a particular pH. −be able to complete the worksheet (if you haven’t already done so…) 11 Learning Outcomes: