POTENTIOMETRIC TITRATIONS

POTENTIOMETRIC TITRATIONS

Electrodes used in potentiometric titrations

The experimental setup for potentiometric measurement comprises a set of an indicating and reference electrodes or two identical indicating electrodes, which should be treated carefully. Do not place the electrodes anywhere except attached to the electrode holder. At the end of the experiment rinse the electrodes and place each one in its housing as required.

Glass electrode. Combined glass reference electrode consists of indicator and reference electrodes in the same body. Great care should be taken of it: never touch the glass part of the electrode with anything except soft tissue paper. While in use, the bulb of the glass electrode and the diaphragm of the reference electrode should be immersed in solution. For short-term storage the combined glass electrode should be immersed in solution of 2 M KCl. Buffer solutions of known pH are used for the pH calibration. The pH values of some buffers are temperature dependent. For high accuracy, calibration and measurements are to be performed at the same temperature.

Ion-selective electrodes are used for detection of specific ions in a mixture of ions. The sensor element, ion-selective membrane, has a construction similar to that of glass electrode. For calibration of ion-selective electrodes a standard addition method is often employed.

Silver indicating electrodes are silver wires with 1-2 mm diameter. When used in precipitation titration, the silver-salt precipitate should be occasionally removed from the electrode surface (mechanically with fine grade emery paper, or chemically immersing the electrode in NH₃ solution). It is simpler, however, to prevent the coating of the electrodes by addition of a surfactant as polyvinyl alcohol (1 drop 0.3% PVA to every 5 ml of solution).

Mercury-coated indicating electrodes are reported to be prepared by lightly amalgamating a gold wire. The disadvantage in use of gold is that it is consumed with time by the amalgam formation. Silver wire used instead of gold, however, can serve many years. The preparation of mercury-coated silver electrode is done by the instructor in a hood (mercury vapors are poisonous!). The silver wire (~1.5 mm diameter) is (a) rubbed with emery paper, rinsed with distilled water and dried with tissue; (b) dipped into mercury to form an amalgame; (c) the mercury is gently spread on the wire with soft tissue. This electrode may be used during several runs of titrations without any renewal. For renewal, step (a) may be omitted.

Platinum redox electrodes are used in redox potentiometric titrations. In excess of oxidant oxide films are formed on the platinum electrodes. The potential response of the electrode is distorted, and the film must be removed. Efficient pretreatment is achieved by cathodically polarizing the platinum electrode in 0.5 MH₂SO₄ at current density of 0.5 mA/cm² for 5 - 15 min. Platinum wire is recommended to use as an auxiliary electrode.

Gold redox electrodes are seldom used in potentiometric titrations. According to our recent experience, the gold electrodes are better behaved than platinum electrodes in view of rate of response and stability toward formation of oxides. These features are of high importance in continuous mode of titration. A good example is the use of gold electrodes in the titration of ascorbic acid with bromine in continuous mode, where the response of the platinum electrode is unsatisfactory.

Reference electrodes. Calomel and silver/silver-chloride electrodes are commonly used in potentiometric titration. In the case of possible interferences of chlorides (as in determination of halides), a mercurous sulfate electrode may be used. In the following series of experiments a home made Ag/AgCl/1 M KClreference electrode is used. Its potential is -19 mV vs SCE, at 25⁰C.

Acid - Base Indicators and Titrations

Acid - Base Indicators and Titrations

Acid - Base indicators (also known as pH indicators) are substances which change colour with pH. They are usually weak acids or bases, which when dissolved in water dissociate slightly and form ions.

Consider an indicator which is a weak acid, with the formula HIn. At equilibrium, the following equilibrium equation is established with its conjugate base:

colourless (Acid) pink (Base)Phenolphthalein is a colourless, weak acid which dissociates in water forming pink anions. Under acidic conditions, the equilibrium is to the left,and the concentration of the anions is too low for the pink colour to be observed. However, under alkaline conditions, the equilibrium is to the right, and the concentration of the anion becomes sufficient for the pink colour to be observed.

The acid and its conjugate base have different colours. At low pH values the concentration of H₃O⁺ is high and so the equilibrium position lies to the left. The equilibrium solution has the colour A. At high pH values, the concentration of H₃O⁺ is low - the equilibrium position thus lies to the right and the equilibrium solution has colour B.

Phenolphthalein is an example of an indicator which establishes this type of equilibrium in aqueous solution:

We can apply equilibrium law to indicator equilibria - in general for a weak acid indicator:

K_ln is known as the indicator dissociation constant. The colour of the indicator turns from colour A to colour B or vice versa at its turning point. At this point:

So from equation:

The pH of the solution at its turning point is called the pK_ln and is the pH at which half of the indicator is in its acid form and the other half in the form of its conjugate base.

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Indicator Range

At a low pH, a weak acid indicator is almost entirely in the HIn form, the colour of which predominates. As the pH increases - the intensity of the colour of HIn decreases and the equilibrium is pushed to the right. Therefore the intensity of the colour of In^- increases. An indicator is most effective if the colour change is distinct and over a low pH range. For most indicators the range is within ±1 of the pK_ln value: - please see the table below for examples, to the right is a model of the acid form of each indicator - with the colour of the solution at the turning point.

Indicator		Colour			pKln	pH range
		Acid	Base
Thymol Blue - 1st change		red	yellow		1.5	1.2 - 2.8
Methyl Orange		red	yellow		3.7	3.2 - 4.4
Bromocresol Green		yellow	blue		4.7	3.8 - 5.4
Methyl Red		yellow	red		5.1	4.8 - 6.0
Bromothymol Blue		yellow	blue		7.0	6.0 - 7.6
Phenol Red		yellow	red		7.9	6.8 - 8.4
Thymol Blue - 2nd change		yellow	blue		8.9	8.0 - 9.6
Phenolphthalein		colourless	pink		9.4	8.2 - 10.0

A Universal Indicator is a mixture of indicators which give a gradual change in colour over a wide pH range - the pH of a solution can be approximately identified when a few drops of universal indicator are mixed with the solution.

Indicators are used in titration solutions to signal the completion of the acid-base reaction.