Once more into the world of hydronium and logarithm; but never mind the mole

By Associate Professor Allan Blackman

This article was orignally published in the Otago Daily Times on Wednesday 4 July 2012.

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What was hopefully obvious from last month’s column was that pH is far from a simple concept. For starters, it is a logarithmic function. In simple terms, this means that a change of 1 pH unit corresponds to a 10-fold change in the hydronium ion concentration – at the risk of upsetting the chemistry purists, one could say that a solution of pH 3 is 10 times as acidic as one of pH 4. To put this in a more understandable context, suppose we had 1 litre of a solution of pH 3 – if we added 9 litres of water to this (i.e. a 10-fold dilution) the final solution would have a pH of 4.

The useful pH scale ranges from 0 (a very acidic solution) to 14 (a very basic solution). Because of its logarithmic nature, this means that it spans a hydronium ion concentration range of 1 × 10¹⁴, or 100,000,000,000,000, between these pH values. To give some idea of the pH values of common substances, lemon juice, for example, has a pH around 2.3, orange juice, around 3.5, milk, around 6.7, seawater, around 8, household ammonia, around 11.5, and oven cleaner can be as high as 13, depending on its composition. Although it is supposedly common knowledge that pure water at 25 °C has a pH of 7.00, measurement of the pH of a sample of any water under all but the most stringently controlled conditions will yield a value somewhere between 5 and 6; this is because the water sample will contain dissolved carbon dioxide from the air, which renders the water very slightly acidic through formation of small amounts of ‘carbonic acid’.

So this is where we get to the importance of pH. Nature has evolved so that many of its important chemical reactions, particularly those that occur in living systems, are optimised to occur at particular pH values. If the pH of the system becomes too high or too low, then critical chemical reactions are impeded, and this can be fatal for the organism. For example, normal human blood has a pH between 7.35 and 7.45 – if our blood pH lowered to 7 or increased to 8, we would probably die. Nature has therefore developed a series of chemical species we call buffers, which ensure that the pH of blood does not change significantly.

Sadly, despite all I have written here, a true appreciation of exactly what pH means is contingent on understanding the mole, a chemical concept which is usually first introduced in 6th form (Year 12) Chemistry and is not necessarily understood by all even when University rolls around. My explanation of pH has only scratched the surface and is extremely simplistic – but hopefully it had given you some idea of what pH is all about.

Of course, the fact that pH is conceptually difficult doesn’t stop advertisers telling us that their clients’ products are ‘pH balanced’ ‘pH neutralising’, and other such meaningless terms. Treat all such claims with caution.