Acid–base balance: the basics

Published on 09/04/2015 by admin

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Last modified 09/04/2015

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1.4 Acid–base balance

The basics

The terms acidity and alkalinity simply refer to the concentration of free hydrogen ions (H⁺) in a solution. H⁺ concentration can be expressed directly in nanomoles per litre (nmol/L) or as pH (see over).

Solutions with high H⁺ (low pH) are acidic and those with low H⁺ (high pH) are alkaline. The point at which a substance changes from alkali to acid is the neutral point (pH = 7, H⁺ = 100 nmol/L).

An acid is a substance that releases H⁺ when it is dissolved in solution. Acids therefore increase the H⁺ concentration of the solution (lower the pH). A base is a substance that accepts H⁺ when dissolved in solution. Bases therefore lower the H⁺ concentration of a solution (raise the pH). A buffer is a substance that can either accept or release H⁺ depending on the surrounding H⁺ concentration. Buffers therefore resist big changes in H⁺ concentration.

Human blood normally has a pH of 7.35–7.45 (H⁺ = 35–45 nmol/L) so is slightly alkaline. If blood pH is below the normal range (< 7.35), there is an acidaemia. If it is above the normal range (> 7.45), there is an alkalaemia.

An acidosis is any process that lowers blood pH whereas an alkalosis is any process that raises blood pH.

What is pH?

The pH (power of hydrogen) scale is a simplified way of expressing large changes in H⁺ concentration, though if you’ve not come across it before you might think it was designed just to confuse you!

It is a negative logarithmic scale (Figure 10). The ‘negative’ means that pH values get lower as the H⁺ concentration increases (so a pH of 7.1 is more acidic than 7.2). The ‘logarithmic’ means that a shift in pH by one number represents a 10-fold change in H⁺ concentration (so 7 is 10 times more acidic than 8).