Buffer: Ph and Common Ion Essay

We judge ourselves by what we feel capable of doing, while others judge us by what we have already done. --Longfellow

Acid-base indicators are generally weak acids (hence we typically add them in small amounts) which are distinguished by the fact that the molecular form of the acid is a different color than the conjugate base. For the hypothetical indicator HIn we might write: HIn + H2O ' H3O+ + Incolor 1 color 2

The Common Ion Effect and Buffer Solutions
From our discussion in class you are already aware that acids and bases can be divided into two large groups: weak and strong. One characteristic which distinguishes, for example, HCl from CH3COOH is the extent to which each molecular substance dissociates into ions in solution. We make this distinction on paper when we write the expressions: HCl(g) + H2O(ℓ) → H3O+(aq) + Cl-(aq)
CH3COOH(aq) + H2O(ℓ) ' H3O+(aq) + CH3COO-(aq)
The first equation is a system which goes to completion while the second one represents an equilibrium system. So we are not surprised to find that a 1 Molar solution of HCl will have more H3O+(aq) present than a 1 Molar solution of CH3COOH.
Since the weak acid system is subject to the same factors which affect other systems at equilibrium, we might also expect some kind of reaction to occur when excess acetate ion (CH3COO-), the conjugate base of acetic acid, is added. In contrast, the HCl should not be affected by the addition of excess Cl-.
Mixed solutions involving weak acids or bases and their conjugates demonstrate an interesting application of the common ion effect: maintaining the pH within a narrow range. Such mixtures are called buffer solutions. Their behavior is based on establishing excesses of both the original acid or base, and the conjugate
(generally obtained by adding a salt). For the hypothetical pair HA (a weak acid) and
NaA (a salt containing the ion A-, the conjugate base of HA) this system of reactions is relevant in aqueous solution:
HA(aq) + H2O(ℓ) ' H3O+(aq) + A-(aq)
NaA(s) → Na+(aq) + A-(aq)
The presence of excess A- (the "common ion") causes a shift in the equilibrium of the first reaction and sets up the required condition for buffering behavior.
In this experiment you will have a chance to investigate these and other behaviors by using a universal indicator.

Changing the concentration of H3O+ (or the pH ) will thus change the color of the indicator. An indicator like that represented above will probably have three colors since at the 50% distribution point (when [HIn] = [In-]) colors 1 and 2 will blend to yield a transition color. This is true of indicators like bromthymol blue which is yellow in acidic solutions, blue in basic solutions, but green in neutral mixtures. Phenolphthalein is somewhat unusual in that the molecular form is colorless. Thus there is only the transition from colorless to pink as a solution becomes basic.
The universal indicator you will use in this experiment contains both of these indicators and others. Each indicator in the recipe is chosen so that the pH range from 1 to 14 is covered either by individual color changes or by overlapping colors of several indicators. The effect is like "liquid pH paper". Incidentally, the technique of choosing the proper indicator for a situation is not difficult. A table of indicator Ka values is all that is required. An indicator is chosen so that its pKa (-log Ka) is the same as the pH of interest. Preparing to experiment
You will be provided with the following materials:
1.
2.
3.
4.
5.

0.10 M HCl solution solid NaCl
0.10 M CH3COOH solution solid NaCH3COO a mixed solution that is 1.0 M in both
CH3COOH and NaCH3COO
6. a mixed solution that is 0.10 M in both
CH3COOH and NaCH3COO
7. 0.20 M NaOH
8. pH buffer solutions (1-6)
9. universal indicator solution (use 2 drops each time)
10. a 24/96-well combination micro-plate
11. a disposable beral pipet
12. a plastic stirrer

Design an experiment in which you:
--determine the pH of each acid-containing solution

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