The Inorganic Carbon Complex
109
(0.001 mol H + /1). The "p" of pH refers to the potential (puissance) of the hydrogen ion
activity. Therefore, by definition, the average of pH values cannot be determined by
simple arithmetic, but must be calculated from the antilogarithms (concentration
units).
For very dilute solutions ( < 0.001 M), activity coefficients approximate unity, and
thus activity approximates concentration (see Exercise 7, pp. 101). Rigorously,
however, pH should be defined as:
pH=-log[H+]y
where y = the activity coefficient for hydrogen ion (H +).
Measurement of pH
Ostensibly, pH is one of the easiest measurements to make; in fact, it is extremely
difficult to obtain accurate measurements, particularly under field conditions. Measurement of pH in limnology can be made either by color reactions of organic
compounds (pH indicators) or by electrometric methods (pH meters). Since electrometric pH meters are available so commonly and permit much more reliable
measurements than do color indicators, electrometric measurements should be used
whenever possible.
Method 1: Colorimetric (Indicators). Certain organic compounds change color as the
pH of the surrounding medium varies. If a given indicator were added to a series of
buffer solutions of known pH values, a series of color standards would be formed for
each pH. The pH of an unknown sample would be estimated by adding the indicator
and comparing the color with the standards. Each ofthe commonly used indicators has
a limited pH range; therefore, several sets of standards must be prepared, each with a
different indicator. Details of common methods employed are given in Golterman and
Clymo (1969). Such methods are much less accurate and convenient than electrometric
measurements and, if more than a few measurements were to be made, would be more
expensive when all costs are considered. Colorimetric measures are not suitable for
waters of low buffer capacity nor for waters containing colored materials such as
dissolved organic compounds.
Various sets of liquid color standards (sealed in glass) or stained glass are available
for comparison to water samples treated with organic indicators. Various test papers,
impregnated with a series of colored organic dyes, permit estimation of pH by reaction
with a water sample and comparison to a standard color sequence.
Method 2: Electrometric (pH Meters). The pH of a sample should be measured by an
electrometer connected to two electrodes (or one electrode in which both are combined)
that are immersed in the solution. The electrode system consists of two components
(Fig. 8.1). The glass electrode is sealed; the sensitive part is on the bottom. Extreme care
must be exercised to avoid scratching this sensitive surface. Abrasions to this area cause
sluggish, variable responses. The glass electrode is in contact with a layer of water
adsorbed to its surface. The electrode potential is a function of the H + concentrations
in the water and in the electrolyte within the electrode. This potential varies by 59.16 mv
per pH unit (25°C). The electrical potential developed at this electrode depends on H +
ion concentration. Therefore, sufficient time must be allowed for ions of this layer to
come into equilibrium with those of the sample water. In dilute waters this equilibrium
takes a relatively long time to be reached, particularly if the electrode were stored in a
strong buffer solution prior to use.
109
(0.001 mol H + /1). The "p" of pH refers to the potential (puissance) of the hydrogen ion
activity. Therefore, by definition, the average of pH values cannot be determined by
simple arithmetic, but must be calculated from the antilogarithms (concentration
units).
For very dilute solutions ( < 0.001 M), activity coefficients approximate unity, and
thus activity approximates concentration (see Exercise 7, pp. 101). Rigorously,
however, pH should be defined as:
pH=-log[H+]y
where y = the activity coefficient for hydrogen ion (H +).
Measurement of pH
Ostensibly, pH is one of the easiest measurements to make; in fact, it is extremely
difficult to obtain accurate measurements, particularly under field conditions. Measurement of pH in limnology can be made either by color reactions of organic
compounds (pH indicators) or by electrometric methods (pH meters). Since electrometric pH meters are available so commonly and permit much more reliable
measurements than do color indicators, electrometric measurements should be used
whenever possible.
Method 1: Colorimetric (Indicators). Certain organic compounds change color as the
pH of the surrounding medium varies. If a given indicator were added to a series of
buffer solutions of known pH values, a series of color standards would be formed for
each pH. The pH of an unknown sample would be estimated by adding the indicator
and comparing the color with the standards. Each ofthe commonly used indicators has
a limited pH range; therefore, several sets of standards must be prepared, each with a
different indicator. Details of common methods employed are given in Golterman and
Clymo (1969). Such methods are much less accurate and convenient than electrometric
measurements and, if more than a few measurements were to be made, would be more
expensive when all costs are considered. Colorimetric measures are not suitable for
waters of low buffer capacity nor for waters containing colored materials such as
dissolved organic compounds.
Various sets of liquid color standards (sealed in glass) or stained glass are available
for comparison to water samples treated with organic indicators. Various test papers,
impregnated with a series of colored organic dyes, permit estimation of pH by reaction
with a water sample and comparison to a standard color sequence.
Method 2: Electrometric (pH Meters). The pH of a sample should be measured by an
electrometer connected to two electrodes (or one electrode in which both are combined)
that are immersed in the solution. The electrode system consists of two components
(Fig. 8.1). The glass electrode is sealed; the sensitive part is on the bottom. Extreme care
must be exercised to avoid scratching this sensitive surface. Abrasions to this area cause
sluggish, variable responses. The glass electrode is in contact with a layer of water
adsorbed to its surface. The electrode potential is a function of the H + concentrations
in the water and in the electrolyte within the electrode. This potential varies by 59.16 mv
per pH unit (25°C). The electrical potential developed at this electrode depends on H +
ion concentration. Therefore, sufficient time must be allowed for ions of this layer to
come into equilibrium with those of the sample water. In dilute waters this equilibrium
takes a relatively long time to be reached, particularly if the electrode were stored in a
strong buffer solution prior to use.