Illustrated Guide to Home Chemistry Experiments

200 DIY Science: Illustrated Guide to Home Chemistry Experiments

LABORATORY 11.3:

oBSERvE THE CHARACTERISTICS of A BUffER SoLUTIoN

A buffer solution is usually a solution of a weak

acid and its conjugate base or, less commonly,

a solution of a weak base and its conjugate

acid. A buffer solution resists changes in the

concentrations of the hydronium ion and

hydroxide ion (and therefore pH) when the

solution is diluted or when small amounts of an

acid or base are added to it. The resistance of a

buffer solution to pH change is based upon Le

Chatelier’s Principle and the common ion effect.

RIREEqU d EqUIpmENT ANd SUppLIES

£ goggles, gloves, and protective clothing

£ beaker, 150 mL (2)

£ beaker, 250 mL

£ graduated cylinder, 100 mL

£ graduated cylinder, 10 mL

£ pipette, 1.00 mL

£ stirring rod

£ pH meter

£ acetic acid, 1.0 m (100 mL)

£ sodium acetate, 1.0 m (100 mL)

£ hydrochloric acid, 6.0 m (20 mL)

£ sodium hydroxide, 6.0 m (20 mL)

£ distilled or deionized water (boil and cool before use)

One common example of a buffer solution is a solution of acetic
acid (the weak acid) and sodium acetate (its conjugate base).
In solution, acetic acid reaches an equilibrium illustrated by the
following equation.

CH 3 CooH(aq) + H 2 o(l) ⇔ CH 3 Coo– + H 3 o+

As we learned earlier in this chapter, acetic acid does not
dissociate completely in solution. For example, in a 1 M solution
of acetic acid, only about 0.4% of the acetic acid molecules
dissociate into hydronium and acetate ions, leaving most of the
acetic acid in molecular form. Dissolving sodium acetate in the
acetic acid solution forces the equilibrium to the left, reducing the
hydronium ion concentration and therefore increasing the pH of
the solution.

Consider what happens if you add a small amount of a strong
acid or strong base to this buffer solution. Ordinarily, you
would expect adding a small amount of a strong acid or base
to cause a large change in the pH of a solution. But if you add
hydrochloric acid (a strong acid) to the acetate/acetic acid buffer
solution, the hydronium ions produced by the nearly complete
dissociation of the hydrochloric acid react with the acetate ions
to form molecular (non-dissociated) acetic acid. According to Le
Chatelier’s Principle, the equilibrium is forced to the left, reducing
the concentration of hydronium and acetate ions, and increasing
the concentration of the molecular acetic acid in the solution.

The acid dissociation constant for this buffer is:

ka = [H 3 o+] · [CH 3 Coo-]/[CH 3 CooH]

If the buffer solution contains equal amounts of acetic acid and
sodium acetate, we can assume that the sodium acetate is fully
dissociated and that dissociation of the acetic acid is negligible

(because the high concentration of acetate ions from the sodium

acetate drives the dissociation equilibrium for acetic acid far to

the left). We can therefore assume that the concentrations of

CH 3 COO– and CH 3 COOH are essentially identical, and simplify

the equilibrium equation to:

ka = [H 3 o+]

which means that the pH of this buffer solution is equal to the pKa.

If we add hydrochloric acid to the buffer solution, the HCl

ionizes completely in solution, yielding hydronium ions and

chloride ions. According to Le Chatelier’s Principle, the increase

in hydronium ions forces the acetic acid equilibrium to the left,

decreasing the concentration of hydronium ions and increasing

the concentration of molecular acetic acid. This equilibrium shift

changes the effective number of moles of acetic acid and acetate

ions, which can be calculated as follows:

final CH 3 Coo – moles = initial CH 3 Co moles – initial HCl moleso H

final CH 3 C moles = initial CHoo H 3 Coo – moles + initial HCl moles