This chapter is barely relevant to Section J2(ii) of the 2017 CICM Primary Syllabus, which expects the exam candidates to "describe the methods of measurement of pH in blood", in the sense that it mentions the concept of pH. In truth it can only really be categorised into the extremely vague Section J1(i) , "explain the principles underlying acid-base chemistry". This concept is definitely an underlying principle, as one is constantly coming across these terms in professional papers, and it is essential to understand what is meant by "strong" and "weak" acids, or "strong" and "weak" bases when one is trying to comprehend the concept of pKa for the purpose of learning pharmacokinetics.
There are multiple different ways to explain this to somebody, and some are better than others. What follows is an earnest attempt by the author to explain these concepts to himself. The interested reader is directed to the football analogy of pKa, which, according to Silverstein (2000), "has the rather formidable disadvantage of being based upon a sport about which more than half the population of the United States, and an even higher proportion of the world population, knows relatively little". From this, it naturally follows that one should be able to represent the rules of sport by using the Brønsted–Lowry definition of acids and bases as an analogy, and to use this conceptual framework for teaching football to nerds.
The Brønsted–Lowry definition of an acid as a proton donor and a base as a proton acceptor represents the presence of these substances in solution in terms of pairs, i.e. each acid contains a proton or two, which it donates, and a conjugate base which is the molecule that remains when the protons are gone. This is represented by the equation,
HA + H2O → H3O+ + A-
where the A- is the conjugate base. One can therefore conceive of a situation where the solution of an acid or base is at some kind of an equilibrium, where not all of the acid has become H3O+ + A-, and all three versions are present in the solution (the original substance, the conjugate base, and the protons).
As the activity of protons is at the base of the definition for pH, one should intuitively grasp that any acid that remains undissociated in solution is not going to give up its protons, and therefore will not add to the activity of protons in the solvent, and is therefore not any sort of acid at all; whereas an acid that dissociates readily and completely will change the pH profoundly.
On this basis, we can define the strength of acids and bases by the extent to which they dissociate. This is represented as Ka, the acid dissociation constant, which is defined as
Ka = (A-) × (H+ ) / HA
Stronger acids have a larger Ka and dissociate completely, i.e. there should be no undissociated H2SO4 in your beaker, and only 2H+ and SO42- (the threshold for calling something "strong" is 100% dissociation). In contrast, acetic acid is a "weak" acid, and there is a lot of CH3COOH in the water. This is obviously also dependent on the solvent, as one acid may be strong in water but weak in another acid (though for the majority of the time in medical chemistry this does not matter, as we discuss ions in aquous solution, and ignore anything that might be happening in molten sulfur or liquid methane).
The strength or weakness generally depends on how easy it is for the proton to leave the conjugate base, which is in turn determined by the polarity of the H-A bond and the size of the atom to which the protons are bound. In this way one can imagine that the strongest acids will be those that have the largest A ions. This is correct; strong acids include:
The concept of Ka is obviously matched with the concept of Kb, or the dissociation constant for a base, which is the same idea, except comparing A+ and OH-.
pKa is the negative log(10) of Ka, which means a lower pKa means a stronger acid.
It is a useful concept because it matches the definition of pH, and can be used operationally to describe what the acid will be doing in a solution of a given acidity. Specifically, an acid of base will be 50% ionised when the pH of its solution is the same as its pKa. The impact of this on ionisation can be represented as a handy alliterative mnemonic device:
Acids are more ionised Above their pKa
Bases are more ionised Below their pKa