- •8.1 Activity Coefficients
- •8.2 Equilibrium Constants
- •Table 8.5 Ionic Product Constant of Water
- •8.2.1 Proton-Transfer Reactions
- •8.2.2 Formation Constants of Metal Complexes
- •8.3 Buffer Solutions
- •8.3.1 Standard Reference pH Buffer Solutions
- •8.3.2 Standards for pH Measurement of Blood and Biological Media
- •8.3.3 Buffer Solutions Other Than Standards
- •Table 8.20 Potentials of Reference Electrodes in Volts as a Function of Temperature
- •8.4.1 Electrometric Measurement of pH
- •8.5 Indicators
- •Table 8.24 Mixed Indicators
- •8.6 Electrode Potentials
- •8.7 Conductance
ELECTROLYTES, EMF, AND CHEMICAL EQUILIBRIUM |
8.115 |
8.4.1 Electrometric Measurement of pH
The pH value is defined for an aqueous solution in an operational (arbitrary but reproducible) manner according to the Bates-Guggenheim convention:
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pH x pH s |
E x E s |
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2.3026 RT /F |
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where |
R is the gas constant per mole, |
T is the temperature on the absolute scale, and |
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F is the faraday. |
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The pH |
x of the unknown medium is calculated |
from that of an |
accepted standard |
(pH |
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s ) and the |
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measured difference in the emf ( |
E ) of the electrode |
combination when the standard solution is |
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removed from the cell and replaced by the unknown. The double vertical line marks a liquid junction. |
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Electrodes as fabricated exhibit variations in the reproducibility of the reference electrode, in the |
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liquid-junction potential, and, with glass electrodes, in the asymmetry potential. These differences |
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are all eliminated in the standardizing procedure with standard reference pH buffers. (See R. G. |
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Bates, |
Determination of pH, Theory and Practice, |
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Wiley, New York, 1964.) |
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Electrode reversible |
Standard reference |
Salt bridge |
Reference |
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to hydrogen ions |
buffer or unknown |
(KCl. 3.5 |
M |
electrode |
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solution |
or saturated) |
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An electrometric pH-measurement system consists of (1) pH-responsive electrode, (2) reference |
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electrode, and (3) potential-measuring device— some form of high-impedance electronic voltmeter |
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for glass-electrode combinations and this or a potentiometer arrangement for other pH-responsive |
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electrodes. Electronic pH meters are simply voltmeters with scale divisions in pH units which are |
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equivalent to the values of 2.3026 |
RT /F (in mV) per |
pH unit. Values of this function at several |
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temperatures are given in Table 8.22. There is no compensation incorporated in the meter for the |
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changes in pH of the test solution as a function of temperature. Reliability of an indicator– reference |
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electrode combination must be ascertained by standardization of the pH meter with one standard |
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buffer and checking the pH response by immersing the combination in a second and different ref- |
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erence buffer. |
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The temperature compensator on a pH meter varies the instrument definition of a pH unit from |
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54.20 mV at 0 |
C to perhaps 66.10 mV at 60 |
C. This permits one to measure the pH of the sample |
(and reference buffer standard) at its actual temperature and thus avoid error due to dissociation equilibria and to junction potentials which have significant temperature coefficients.
TABLE 8.22 Values of 2.3026 |
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RT /F at Several Temperatures |
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In millivolts. |
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t C |
Value |
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t C |
Value |
t C |
Value |
t C |
Value |
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0 |
54.197 |
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25 |
59.157 |
50 |
64.118 |
80 |
70.070 |
5 |
55.189 |
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30 |
60.149 |
55 |
65.110 |
85 |
71.062 |
10 |
56.181 |
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35 |
61.141 |
60 |
66.102 |
90 |
72.054 |
15 |
57.173 |
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38 |
61.737 |
65 |
67.094 |
95 |
73.046 |
18 |
57.767 |
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40 |
62.133 |
70 |
68.086 |
100 |
74.038 |
20 |
58.165 |
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45 |
63.126 |
75 |
69.078 |
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Report of the National Academy of Sciences: National Research Council Committee of Fundamental Constants, 1963.