Page 87 - Vogel's TEXTBOOK OF QUANTITATIVE CHEMICAL ANALYSIS

STABlLlTY CONSTANTS OF EDTA COMPLEXES 2.27

In equation (q) only the fully ionised form of EDTA, i.e. the ion Y4-, has
been taken into account, but at low pH values the species HY3-, HzYZ-, H3Y -
and even undissociated H4Y may well be present; in other words, only a part
of the EDTA uncombined with metal may be present as Y4-. Further, in
equation (q) the metal ion Mn+ is assumed to be uncomplexed, i.e. in aqueous
solution it is simply present as the hydrated ion. If, however, the solution also
contains substances other than EDTA which can complex with the metal ion,
then the whole of this ion uncombined with EDTA may no longer be present
as the simple hydrated ion. Thus, in practice, the stability of metal-EDTA
complexes may be altered (a) by variation in pH and (b) by the presence of
other complexing agents. The stability constant of the EDTA complex will then
be different from the value recorded for a specified pH in pure aqueous solution;
the value recorded for the new conditions is termed the 'apparent' or
'conditional' stability constant. It is clearly necessary to examine the effect of
these two factors in some detail.

(a) pH effect. The apparent stability constant at a given pH may be calculated
from the ratio Klcc, where cc is the ratio of the total uncombined EDTA (in al1
forms) to the form Y4-. Thus K,, the apparent stability constant for the
metal-EDTA complex at a given pH, can be calculated from the expression

log K, = log K - log cc (30)
The factor cc can be calculated from the known dissociation constants of EDTA,
and since the proportions of the various ionic species derived from EDTA will
be dependent upon the pH of the solution, cc will also Vary with pH; a plot of
log cc against pH shows a variation of log cc = 18 at pH = 1 to log cc = O at
pH = 12: such a curve is very useful for dealing with calculations of apparent
stability constants. Thus, for example, from Table 2.4, log K of the EDTA
complex of the PbZ+ ion is 18.0 and from a graph of log cc against pH, it is
found that at a pH of 5.0, log cc = 7. Hence from equation (30), at a pH of 5.0
the lead-EDTA complex has an apparent stability constant given by:
log K, = 18.0 - 7.0 = 11.0
Carrying out a similar calculation for the EDTA complex of the MgZ+ ion
(log K = 8.7), for the same pH (5.0), it is found:
log K,(Mg(II) - EDTA) = 8.7 - 7.0 = 1.7
These results imply that at the specified pH the magnesium complex is
appreciably dissociated, whereas the lead complex is stable, and clearly titration
of an Mg(I1) solution with EDTA at this pH will be unsatisfactory, but titration
of the lead solution under the same conditions will be quite feasible. In practice,
for a metal ion to be titrated with EDTA at a stipulated pH the value of log KH
should be greater than 8 when a metallochromic indicator is used.
As indicated by the data quoted in the previous section, the value of log cc
is small at high pH values, and it therefore follows that the larger values of
log KH are found with increasing pH. However, by increasing the pH of the
solution the tendency to form slightly soluble metallic hydroxides is enhanced
owing to the reaction: