Determination of pH in precipitation

4.7  Determination of pH in precipitation

4.7.1  Potentiometric method

4.7.1.1  Principle

The method is based on the determination of the potential difference between an electrode pair consisting of a glass electrode sensitive to the difference in the hydrogen ion activity in the sample solution and the internal filling solution, and a reference electrode, which is supposed to have a constant potential independent of the immersing solution. The measured potential difference is compared with the potential obtained when both electrodes are immersed in a solution or buffer with known pH or hydrogen ion concentration. The pH is defined by the formula:
          pH_(sample) = pH_(reference) + (E_(sample) – E_(reference) F/RT1n10
where E are the electrode potentials, R is the universal gas constant, T the absolute temperature and F is the Faraday constant.
This is an operationally defined pH. Buffers of known pH are specified by National Bureau of Standards, now the National Institute of Standardized Technology (NIST). The primary standard and the most widely used buffer for pH-meter calibration is 0.05 M potassium hydrogen phthalate, which has a pH of 4.00 at 20° C, and a hydrogen ion activity of 10^-4 M. This latter hydrogen ion activity is based on theoretical calculations (the Bates-Guggenheim convention).
In precipitation samples, the ionic strength will typically be in the region 10^-3 to 10^-5. The activity coefficient for monovalent cations such as the hydrogen ion will therefore be in the range 0.95-0.99. This corresponds to <0.02 pH-units difference between pH and -log(H⁺). Much more critical is the assumption of a constant reference electrode potential when going from a relatively concentrated potassium hydrogen phthalate solution to extremely dilute precipitations samples. The problem arises because of the inherent possibility of building up a liquid junction potential between the internal solution of the reference electrode, and the sample solution. This liquid junction potential may be larger if the ionic strength difference between the two solutions is large. It is reduced by making the boundary between the concentrated filling solution and the sample as sharp as possible. Various designs of pH cells meeting this criterion have been proposed. Tests of commercial electrodes against dilute acid solutions and low ionic strength buffers with known pH or hydrogen ion concentrations have shown, however, that this problem has largely been overcome with modern pH instrumentation and electrode systems.
However, it is strongly recommended to check the electrode system at regular intervals, by measuring the “apparent pH” of a solution with low ionic strength with known pH or hydrogen ion concentration. The pH readings should be within 0.02 or 0.05 pH-units of the “theoretical” result. If this is not the case, or if the reading is unstable during stirring of the solution, the reference electrode should be replaced. New glass electrodes should be tested against at least two buffers to see that the response is Nernstian.
The reference electrode should preferably be stored in dilute potassium chloride solution (0.1M).

4.7.1.2  Instrumentation

pH-meter with the possibility of reading to the nearest 0.02 pH-units or preferably to the nearest 0.01 pH-unit.
A glass electrode and a reference electrode must be used with the pH-meter. The reference electrode should be suitable for measurement in low-ionic strength solutions and preferably be of the calomel type filled with saturated potassium chloride. Other reference electrodes or combination electrodes may be used, but all electrodes should be checked for acceptable performance.
Magnetic stirrer, with teflon coated stirring bar.
Beakers used for the test solution should be made of borosilicate glass or poly­ethylene.

4.7.1.3  Chemicals

Buffer solutions for the calibration of the pH-meter. Preferably the two buffer solutions given in Section 4.7.1.4, which are recommended as standards by the U.S. National Institute of Standards and Technology (NIST).

4.7.1.4  Reagents

National Bureau of Standards solutions with known pH.

1. 0.05 M potassium hydrogen phthalate (C₆H₄ (COOH) (COOK)
   pH = 4.00 at 20 °C
   pH = 4.01 at 25 °C

Dissolve 10.12 g potassium hydrogen phthalate, C₆H₄ (COOH) (COOK), dried at 120 °C, in 1000 ml distilled water.

2. 0.025 M potassium dihydrogen phosphate (KH₂PO₄) and 0.025 M disodium hydrogen phosphate (Na₂HPO₄)
   pH = 6.88 at 20 °C
   pH = 6.86 at 25 °C

Dissolve 3.39 g potassium dihydrogen phosphate, KH₄PO₄, and 3.53 g disodium hydrogen phosphate, Na₂HPO₄, dried at 120° C, in 1000 ml distil­led water.
Instead of the anhydrous disodium hydrogen phosphate, 4.43 g of undried dihydrate, Na₂HPO₄ · 2 H₂O, may be used.

Commercial available buffer solutions may also be used, but should be checked against the primary standard buffers described above. The buffers should be kept in the dark in well closed bottles of borosilicate or polyethylene.

4.7.1.5  Calibration

Calibrate the pH-meter according to the instruction manual for the instrument using one, or preferably two, buffer solutions. The temperature of the buffer solutions must be known. The calibration should be checked after each set of samples.

4.7.1.6  Analytical procedure

Measure the pH-value of the sample according to the instruction manual for the instrument. The solution may be stirred, but not vigorously. The temperature of the sample solution must be the same as the temperature of the buffer solution used for calibration.
Rinse the electrodes thoroughly with distilled water between each measurement, and wipe off the excess water with a soft paper.
Store the electrodes in 0.1 M KCl-solution or according to the manufacturers recommendations. The reference electrode should not be stored in distilled water!

4.7.1.7  Performance test of the electrode pair

As mentioned in Section 4.7.1.1 the behaviour of the reference electrode is the main source of errors in pH-measurements, especially in low ionic strength solutions. In order to check the performance of the reference electrode, control measurements should be made on solutions of dilute acids or dilute buffers to verify that correct values are obtained for solutions of lower ionic strengths. A solution which should give a pH ~4.00 could be used for the test. A 10^-4M HC1-solution should give a pH of 3.99 ± 0.05.
Electrode pairs should also show minimal differences between measurements made in stirred and unstirred low ionic strength solutions.
Usually the liquid junction between the solution and the saturated KCl-solution in the reference electrode is formed in a porous plut of ceramic fibre. Slow stirring removes the concentrated KCl-solution which slowly runs out through this capillary.
If the stirring is too vigorous, the ionic medium in the plug itself may be diluted. This will increase the liquid junction potential, and should be avoided. The liquid junction potential may also increase if the porous plug is clogged up by impurities.