The doctrine about solution. Buffer solution KARAGANDA STATE MEDICAL UNIVERSITY Karaganda 2014y.

Let us consider the buffer pain of acetic acid and sodium acetate. Acetic acid, being a weak acid, feebly ionizes. On the other hand, sodium acetate ionizes to a large extent. CH 3 COOH ↔ CH 3 COO - + H + CH 3 COONa ↔ CH 3 COO - + Na + When an acid (say HCl) is added, the acetate ion of the byffer bind with H + ions (of HCI) to form acetic acid which is weakly ionizing. There fore, the pH change due to acid is resisted by the byffer. H + + CH 3 COO - → CH 3 COOH When a base (say NaOH) is added the H + ions of the buffer (acetic acid) combine with OH - ions to form wader, which is weakly dissociated. Thus, pH change due to base addition is also prevented by the buffer. OH - + H + → H 2 O

Buffering capacity: The efficiency of a buffer in maintaining a constant pH on the addition of acid or base is referred to as buffering capacity. In mostly depends on the concentration of the buffer components. The maximum buffering capacity is usually achieved by keeping the same concentration of the salt as well as the acid. But how does a buffer work? Suppose a buffered solution contains relatively large quantities of a weak acid HA and its conjugate base A -. When hydroxide ions are added to the solution, since the weak acid represents the best source of protons, the following reaction occurs: ОH - + HA  A - + H 2 O The net result is that OH - ions are not allowed to accumulate but are replaced by A - ions.

The stability of the pH under these conditions can be understood by examining the equilibrium expression for the dissociation of HA: or, rearranging,

In other words, the equilibrium concentration of H +, and thus the pH, is determined by the ratio When OH - ions are added, HA is converted to A -, and the ratio [HA]/[A - ] decreases. However, if the amounts of HA and A - originally present are very large compared with the amount of OH - added, the change in the [HA]/[A - ] ratio will be small. In Sample Exercises 1 and 2, Initially After adding 0.01 mol/l. OH - The change in the ratio [HA]/[A - ] is very small. Thus the [H + ] and the pH remain essentially constant.

Similar reasoning applies when protons are added to a buffered solution of a weak acid and a salt of its conjugate base. Because the A - ion has a high affinity for H +, the added H + ions react with A - to form the weak acid: H + + A -  HA and free H + ions do not accumulate. In this case there will be a net change of A - to HA. However, if [A - ] and [HA] are large compared with the [H + ] added, little change in the pH will occur.

The form of the acid dissociation equilibrium expression is often useful for calculating [H + ] in a buffered solution, since [HA] and [A - ] are known. For example, to calculate [H + ] in a buffered solution containing 0.10 M HF (К a = 7.2  ) and 0.30 M NaF, we simply substitute into Equation (1): Another useful form of Equation (1) can be obtained by taking the negative log of both sides:

That is, or, where inverting the log term reverses the sign: This log form of the expression for К a is called the Henderson-Hasselbalche equation and is useful for calculating the pH of solutions when the ratio [HA]/  A -  known. For a particular buffering system (acid-conjugate base pair), all solutions that have the same ratio [A - ]/[HA] will have the same pH. For example, a buffered solution containing 5.0 M HC 2 H 3 O 2 and 3.0 M NaC 2 H 3 O 2 will have the same pH as one containing M HC 2 H 3 O 2 and M NaC 2 H 3 O 2. This can be shown an follows:

System [HA]/  A -  0.50 M HC 2 H 3 O 2 and 3.0 M NaC 2 H 3 O M HC 2 H 3 O 2 and 3.0 NaC 2 H 3 O 2 Therefore, Note that in using this equation we have assumed that the equilibrium concentrations of A - and HA are equal to the initial concentrations. That is, we are assuming the validity of the approximations and where x is the amount of acid that dissociates. Since the initial concentrations of HA and A - are relatively large in a buffered solution, this assumption is generally acceptable.