Buffer capacity refers to the amount of acid or base a buffer can "absorb" without a significant pH change. It is governed by the concentrations of the conjugate acid and base forms of the buffer. A 0.5 M buffer will require five times as much acid or base as a 0.1 M buffer for a given pH change. In this problem you begin with a buffer of known pH and concentration and calculate the new pH after a particular quantity of acid or base is added. In the laboratory you will carry out some stepwise additions of acid or base and measure the resulting pH values.
In calculating the pH change upon addition of acid or base, it is useful to keep track of what is happening by means of an ICE table (Initial, Change, Equilibrium) as follows. (Your textbook sets up ICE tables in terms of concentrations. It is less confusing when there are volume changes to express the change as moles, rather than molarity. It is also convenient to use mmol as the quantity when using volumes in mL because VmL×M = mmol.
Reaction: HPO42- + H3O+ ? H2PO4-
Initial mmolinitial X mmolinitial
Change -X -X +X
Equilibrium mmolinitial-X small mmolinitial+X
Starting with 45 mL of 0.50 M phosphate buffer, pH = 6.71, you add 4.9 mL of 1.00 M HCl. Using the Henderson Hasselbalch equation with a pK2 for phosphate of 6.64, calculate the following values to complete the ICE table. (Carry out your intermediate calculations to at least one more significant figure than needed. See "Phosphate Buffer Issues" to explain why you do not use the value of 7.21 for pK2 that is given in your textbook.)
What is the composition of the buffer to begin with, both in terms of the concentration and the molar quantity of the two major phosphate species? (Units required).