Quimica

A Practical Introduction to Ion Selective Electrodes-Part I: Theory
by Andrew Holmes
Consultant, Cardiff, U.K.

any of those who start off in finishing will eventually come to the question of solution control. It is usually the case that, by one means or another, a solution component can have its concentration determined by titration. But how accurate a methodology is this? For many solutesthe actual sample size may be very small, and the larger the tank containing the bulk solution the greater the opportunity for measurement error. The use of chemical indicators can also be problematic because the observed color change may take place over, for example, as much as two whole pH units in an acid-alkali titration and additionally may take place gradually so that the actual endpoint ispoorly defined. Clearly, if there were a way of substituting an alternative detector for a chemical indicator that could give a sharp and definite reading and did not entail excessive cost, this would be preferable to the type of errors that could arise from a traditional wet determination. One answer to these problems is known as potentiometry. This makes use of the fact that the activity of acharged solute corresponds to a fixed voltage, which can be detected using a suitable sensor that, in most cases, is designed to detect a single ionic species in a selective manner. Such sensors are usually referred to as ion selective electrodes (lSEs), and potentiometry using these devices is now a mature and well-characterized field of analysis. To put the use of potentiometry into some kind ofperspective, let's suppose one has an acidic solution, say a bright nickel or zinc. These solutions have to be maintained within a comparatively narrow pH range, and at any moment one could test this using a pH paper; however, the previous remarks regarding indicator color change apply, and some brightener systems actually affect the indicator dyes so that their color changes become meaningless. Ifa glass-bulb pH meter is used instead, which is in fact an ion electrode selective for hydrogen ions, then the result is more accurate, and different temperatures between solutions can be compensated for using a temperature reference electrode, again something a simple test paper cannot handle. The average life of a pH probe is roughly 6 months, but because it is
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reusable one couldhave used single-shot test papers equivalent to several times the replacement cost of a single probe. Add to this the fact that the meter unit itself should never need much more attention than the replacement of batteries, plus the improved solution control resulting from its use, and suddenly the cost benefits of using the most common type of ISE in regular use come sharply into focus. In fact thepH meter is only the most well known of a whole range of detectors whose economy, reliability, and reusability commend them to anyone requiring accurate results in chemical process control. What, then, are the principles involved?
HOW IT WORKS

As stated above, a charged solute gives rise to an electrochemical potential that is proportional to its activity. It is important to realize that theISE measures activity rather than concentration because of the involvement of physical factors, particularly temperature. This is explained by the so-called Nernst equation, where the relationship between the electrochemical potential of a solute in a half-cell and these physical parameters is established:
E = ED - RTfNFlloge [a(reduced form)/a(oxidized form)]

where E = standard electrodepotential for reference electrode used; R = gas constant, 8.314 JIK/mole; T = temperature (Kelvin); n = valency of oxidized species; and F = Faraday constant, 9.6487 x 104 Coulombs. We can see immediately that there are in fact very few variables in this equation; temperature may vary, but the place in which the measurement is made should be stable in that respect; Rand Fare constants, and for a...