186
C. ALBERS
temperature S is highest for pure water (chlorinity = 0) and decreases
if the salt content and hence the ionic strength increase. The dissociation
constants show the opposite effect.
In the same way the ionic strength increases the ionic product of
water Klv. If we pass from concentrations to activities we obtain from
Eq. (1)
where aIIt *aoEr- is called the thermodynamic dissociation product. Since
Kw and alr20 are affected by the ionic strength in an opposite direction,
the effect of the ionic strength on the thermodynamic dissociation
product is small. It seems reasonable, therefore, to derive the definition
of neutrality from the thermodynamic dissociation product rather than
from the ionic product K, .
2. EFFECT OF TEMPERATURE
There is no constant in the whole realm of physical chemistry which
does not depend on temperature. As a general rule the dissociation of a
molecule into ions is favored by an increase in temperature. As a consequence most values of K’ increase together with the temperature. This
can be seen from Table I1 which also shows that the effect of temperature on the second dissociation constant for carbonic acid is even greater
than that on the first.
Of particular interest is the dissociation of water. Since with increasing temperature K , increases, the p H indicating neutrality decreases. It is only a t 25°C that pure water has a p H of 7.0, whereas at
5°C it has a p H of 7.36 and at 50°C it has a pH of 6.65. That is to say,
the “meaning” of p H depends on the temperature. A pH of 7.0 “means”
a neutral solution at 25”C, a slightly acid solution at 5”C, and a slightly
alkaline solution at 50°C. To avoid the difficulty of interpreting a p H
value, Winterstein (1954) suggested the use of the ratio [ O H - ] / [ H I
which of course is unity for a neutral solution, smaller than unity for an
acid solution, and larger than unity for an alkaline solution. For similar
reasons Rahn (1967) introduced the term “relative alkalinity” which is
defined by [H’],/[H’], where [H+Ix is the hydrogen ion concentration
at neutrality and [ H’] the actual hydrogen ion concentration. Since
[ H + ] ~ / [ H + ]
=dK,\-/[H+] and [OH-]/[H+] = K,/[H+]2, the relative alkalinity of Rahn is the square root of Winterstein’s term [OH-]/[H+].
A line chart for both terms as a function of temperature and p H is shown
in Fig. 6.
C. ALBERS
temperature S is highest for pure water (chlorinity = 0) and decreases
if the salt content and hence the ionic strength increase. The dissociation
constants show the opposite effect.
In the same way the ionic strength increases the ionic product of
water Klv. If we pass from concentrations to activities we obtain from
Eq. (1)
where aIIt *aoEr- is called the thermodynamic dissociation product. Since
Kw and alr20 are affected by the ionic strength in an opposite direction,
the effect of the ionic strength on the thermodynamic dissociation
product is small. It seems reasonable, therefore, to derive the definition
of neutrality from the thermodynamic dissociation product rather than
from the ionic product K, .
2. EFFECT OF TEMPERATURE
There is no constant in the whole realm of physical chemistry which
does not depend on temperature. As a general rule the dissociation of a
molecule into ions is favored by an increase in temperature. As a consequence most values of K’ increase together with the temperature. This
can be seen from Table I1 which also shows that the effect of temperature on the second dissociation constant for carbonic acid is even greater
than that on the first.
Of particular interest is the dissociation of water. Since with increasing temperature K , increases, the p H indicating neutrality decreases. It is only a t 25°C that pure water has a p H of 7.0, whereas at
5°C it has a p H of 7.36 and at 50°C it has a pH of 6.65. That is to say,
the “meaning” of p H depends on the temperature. A pH of 7.0 “means”
a neutral solution at 25”C, a slightly acid solution at 5”C, and a slightly
alkaline solution at 50°C. To avoid the difficulty of interpreting a p H
value, Winterstein (1954) suggested the use of the ratio [ O H - ] / [ H I
which of course is unity for a neutral solution, smaller than unity for an
acid solution, and larger than unity for an alkaline solution. For similar
reasons Rahn (1967) introduced the term “relative alkalinity” which is
defined by [H’],/[H’], where [H+Ix is the hydrogen ion concentration
at neutrality and [ H’] the actual hydrogen ion concentration. Since
[ H + ] ~ / [ H + ]
=dK,\-/[H+] and [OH-]/[H+] = K,/[H+]2, the relative alkalinity of Rahn is the square root of Winterstein’s term [OH-]/[H+].
A line chart for both terms as a function of temperature and p H is shown
in Fig. 6.