Page 49 - Vogel's TEXTBOOK OF QUANTITATIVE CHEMICAL ANALYSIS

ELECTROLYTIC DISSOCIATION 2.4

tertiary ionisations respectively. As already mentioned, these do not take place
to the same degree. ~he primary ionisation is always greater than the secondary,
and the secondary very much greater than the tertiary.
Acids of the type of acetic acid (CH,COOH) give an almost normal
freezing-point depression in aqueous solution; the extent of dissociation is
accordingly small. It is usual, therefore, to distinguish between acids which are
completely or almost completely ionised in solution and those which are only
slightly ionised. The former are termed strong acids (examples: hydrochloric,
hydrobromic, hydriodic, iodic(V), nitric and perchloric [chloric(VII)] acids,
primary ionisation of sulphuric acid), and the latter are called weak acids
(examples: nitrous acid, acetic acid, carbonic acid, boric acid, phosphorous
(phosphoric(II1)) acid, phosphoric(V) acid, hydrocyanic acid, and hydrogen
sulphide). There is, however, no sharp division between the two classes.
A base was originally defined as a substance which, when dissolved in water,
undergoes dissociation with the formation of hydroxide ions OH- as the only
negative ions. Thus sodium hydroxide, potassium hydroxide, and the hydroxides
of certain bivalent metals are almost completely dissociated in aqueous solution:
NaOH +Na+ +OH-

These are strong bases. Aqueous ammonia solution, however, is a weak base.
Only a small concentration of hydroxide ions is produced in aqueous solution:

General concept of acid and bases. The Bvmsted-Lowvy theovy. The simple
concept given in the preceding paragraphs suffices for many of the requirements
of quantitative inorganic analysis in aqueous solution. It is, however, desirable
to have some knowledge of the general theory of acids and bases proposed
independently by J. N. Br~nsted and by T. M. Lowry in 1923, since this is
applicable to al1 solvents. According to this theory, an acid is a species having
a tendency to lose a proton, and a base is a species having a tendency to add
on a proton. This may be represented as:
Acid = Proton + Conjugate base
A=H++B
It must be emphasised that the symbol H+ represents the proton and not the
'hydrogen ion' of variable nature existing in different solvents (OH:, NH;,
CH3C02H:, C2H50H:, etc.); the definition is therefore independent of
solvent. The above equation represents a hypothetical scheme for defining A
and B and not a reaction which can actually occur. Acids need not be neutral
molecules (e.g., HC1, H2 SO,, CH ,CO2 H), but may also be anions (e.g., HSO; ,
H2P0;, HOOC.CO0-) and cations (e.g. NH;, C6H5NHl, Fe(H20)B+).
The same is true of bases where the three classes can be illustrated by NH,,
C6H5NH2, H20; CH,COO-, OH-, HPOZ-, 0C2H5; Fe(H20)5(OH)2+.
Since the free proton cannot exist in solution in measurable concentration,
reaction does not take place unless a base is added to accept the proton from
the acid. By combining the equations A, =BI + H + and B2 + H +=A2, we
obtain