Acid–Base Theories Compared, and When Each One Runs Out

The three theories and their limits

What Brønsted–Lowry adds

Defining acidity as proton donation removes the requirement for water and introduces conjugate pairs. Every acid has a conjugate base formed by losing a proton, and the strengths are inversely related — a strong acid has a weak conjugate base.

It also allows amphoteric behaviour to be described naturally: water donates a proton to a base and accepts one from an acid, so it is both, depending on the partner.

The levelling effect

In water, any acid stronger than the hydronium ion is fully deprotonated, so all such acids appear equally strong. Water levels them. Distinguishing their intrinsic strengths requires a less basic solvent — one that does not accept protons so readily.

The same applies at the other end: any base stronger than hydroxide is fully protonated in water. Explaining why several strong acids all appear equally strong in aqueous solution is a standard question, and the answer is the levelling effect.

What Lewis theory adds

Defining acidity as electron pair acceptance covers reactions with no proton involved at all. A metal ion accepting a lone pair from a ligand is a Lewis acid–base reaction, which makes the whole of coordination chemistry an application of the theory.

Boron trihalides accepting a lone pair, and metal ions forming complexes, are the standard examples. Neither can be described as an acid–base reaction under the earlier theories.

The cost of the generality is that Lewis acidity is not easily quantified on one scale, since it depends on the partner. That is why HSAB exists as a companion framework — it addresses which Lewis acid prefers which Lewis base, a question the definition alone leaves open.

Solvent system theory

For solvents that self-ionise, an acid is a species increasing the concentration of the solvent's characteristic cation, and a base one increasing the characteristic anion.

In water this reduces to the familiar picture. In liquid ammonia the corresponding species are the ammonium and amide ions, so ammonium salts behave as acids and amides as bases — which lets acid–base chemistry be done in a solvent where water-based definitions do not apply.

This is why liquid ammonia is used for reactions requiring bases far stronger than any that survive in water.

Choosing the right framework

Frequently asked questions

Is every Lewis acid also a Brønsted acid?

No. Boron trifluoride accepts an electron pair but has no proton to donate, so it is a Lewis acid only. Every Brønsted acid, however, is also a Lewis acid.

What is the levelling effect?

The inability of a solvent to distinguish between acids stronger than its own conjugate acid, because all of them are completely ionised. A less basic solvent is needed to compare them.

Why is Lewis theory not simply used for everything?

Because its breadth reduces its predictive content. For aqueous equilibria the Brønsted framework gives quantitative answers that the Lewis definition alone does not.

How does HSAB relate to Lewis theory?

It supplements it. Lewis theory says what an acid and base are; HSAB predicts which acid–base combinations are favoured, which the definition does not address.