Instant Notes: Analytical Chemistry

The practical (oroperational) definition of pH recognizes that it is deter-

mined using electrochemical cells having an electrode selective to hydrogen

ions. This has been discussed in Topics C2 and C3, but a typical cell is:

Reference electrode ||Solution X |glass electrode.

This gives an output e.m.f ,EX,

EX=E*+(RT/F) ln [a(H 3 O+)X]

The constantE* depends on the exact nature of the reference and glass elec-

trodes, and is best eliminated by calibration with a standard solution S which

has a pH that is accurately known.

ES=E*+(RT/F) ln [a(H 3 O+)S]

Subtracting these and converting the logarithms gives a practical definition of

pH:

pH(X) =pH(S) +(ES-EX)/(RT ln(10)/F)

Typical calibration buffers are discussed below.

The pH scale In all aqueous solutions, pH values may range between about 0 and 14 or more
as shown in Figure1. Molar solutions of strong mineral acids, such as HCl,
HNO 3 or H 2 SO 4 have pH values less than 1. Weak acids, such as ethanoic or
citric acid in decimolar solution have a pH of around 3.
A useful standard is 0.05 M potassium hydrogen phthalate which, at 15∞C has
a pH of 4.00. Although pure water is neutral and has a pH of 7.00, freshly
distilled water rapidly absorbs carbon dioxide from the air to form a very dilute
solution of carbonic acid, and therefore has a pH of around 6.

C4 – pH and its control 75

01234567891011121314

Very acid Acidic Neutral Alkaline Very alkaline

Fig. 1. The pH scale.

Another standard occasionally used is 0.05 M borax (sodium tetraborate,

Na 2 B 4 O 7 ), which has a pH of 9.18 at 25∞C.

Dilute alkalis such as ammonia or calcium hydroxide (lime water) have pH

values near to 12, and for molar caustic alkalis, such as NaOH, the pH is over 13.

Buffers As many reactions depend greatly upon the concentration of hydrogen ions in
the solutions being used, it is important to control the pH. This is usually
achieved by using a solution which has a pH that is accurately known and that
resists any change in pH as solvent for the experiment. Such solutions are called
buffers.
The equilibria that govern the reactions of weak acids or bases in aqueous
solution will resist attempts to change them. This is known as Le Chatelier’s
principle. For example, the dissociation of ethanoic acid obeys the equation: