Acid–Base Chemistry Without Water

Almost everything we learn about acids and bases at school happens in water. pH, Kw, “strong acids fully ionise”, neutral at 7: all of it is water chemistry. But water is just one solvent among many. Change the solvent and the rules shift: acids that were strong become weak, weak bases become strong, and the meaning of “neutral” changes completely.

Non-aqueous acid–base chemistry is used in industry, in analytical chemistry and in synthesis. It also deepens our understanding of what acidity really is.

Why water isn’t special

Water is amphiprotic: it can donate a proton and accept a proton. That lets it self-ionise:

2H₂O ⇌ H₃O⁺ + OH⁻

In water, H₃O⁺ is the characteristic acidic species and OH⁻ the characteristic basic species. A substance that increases [H₃O⁺] is an acid; one that increases [OH⁻] is a base.

Many other solvents also self-ionise, producing their own characteristic cation and anion.

The solvent-system definition

In the early 20th century, chemists including Edward Franklin, who studied liquid ammonia extensively, generalised the Arrhenius idea:

• An acid is a substance that increases the concentration of the solvent’s characteristic cation.
• A base is a substance that increases the concentration of the solvent’s characteristic anion.

Solvent	Self-ionisation	Acidic species	Basic species
Water	2H₂O ⇌ H₃O⁺ + OH⁻	H₃O⁺	OH⁻
Liquid ammonia	2NH₃ ⇌ NH₄⁺ + NH₂⁻	NH₄⁺	NH₂⁻ (amide)
Anhydrous sulfuric acid	2H₂SO₄ ⇌ H₃SO₄⁺ + HSO₄⁻	H₃SO₄⁺	HSO₄⁻
Glacial ethanoic acid	2CH₃COOH ⇌ CH₃COOH₂⁺ + CH₃COO⁻	CH₃COOH₂⁺	CH₃COO⁻
Liquid hydrogen fluoride	3HF ⇌ H₂F⁺ + HF₂⁻	H₂F⁺	HF₂⁻
Bromine trifluoride	2BrF₃ ⇌ BrF₂⁺ + BrF₄⁻	BrF₂⁺	BrF₄⁻

The last example has no hydrogen at all, which shows the solvent-system definition can go beyond protons.

Liquid ammonia

Ammonia is a gas at room temperature but a liquid below −33 °C. Its self-ionisation is far smaller than water’s (its ionic product is about 10⁻³³ at −50 °C, compared with water’s 10⁻¹⁴ at 25 °C).

In liquid ammonia:

• Ammonium salts such as NH₄Cl are acids: they supply NH₄⁺.
• Metal amides such as NaNH₂ are bases: they supply NH₂⁻.
• Neutralisation is NH₄⁺ + NH₂⁻ → 2NH₃, the equivalent of H₃O⁺ + OH⁻ → 2H₂O in water.

Ammonia is a much stronger base than water, so acids that are weak in water can be strong in ammonia. Ethanoic acid, only about 1% ionised at 0.1 mol/dm³ in water, is essentially completely ionised in liquid ammonia, because ammonia grabs its proton so readily.

Liquid ammonia is also famous for dissolving alkali metals to give deep blue solutions of solvated electrons, which are powerful reducing agents used in organic synthesis (the Birch reduction).

Glacial ethanoic acid

Pure ethanoic acid is a much weaker base than water; it holds on to protons less willingly. That changes how strong acids behave:

• In water, HClO₄, HBr, H₂SO₄, HCl and HNO₃ all appear equally strong, because water accepts every proton they offer (the levelling effect).
• In glacial ethanoic acid, they only partly ionise, and their differences become measurable: HClO₄ > HBr > H₂SO₄ > HCl > HNO₃.

A solvent like this, which reveals differences between acids, is called a differentiating solvent. Water is a levelling solvent for strong acids.

Glacial ethanoic acid is also used to titrate very weak bases that can’t be titrated in water, such as some amines and pharmaceutical compounds. Perchloric acid in ethanoic acid is an extremely strong titrant, able to protonate bases too weak to give a sharp end point in water. This is a standard method in pharmacopoeia assays of drug purity.

Anhydrous sulfuric acid

Pure sulfuric acid is a highly acidic solvent that self-ionises significantly:

2H₂SO₄ ⇌ H₃SO₄⁺ + HSO₄⁻

In it, most substances that we think of as acids actually behave as bases, because sulfuric acid is so willing to donate protons. Even nitric acid acts as a base, forming the nitronium ion:

HNO₃ + 2H₂SO₄ → NO₂⁺ + H₃O⁺ + 2HSO₄⁻

The nitronium ion is the reactive electrophile in the nitration of benzene, which is why concentrated sulfuric and nitric acids are mixed for nitrations. See superacids for acids that are even stronger than sulfuric acid.

Levelling and differentiating: a summary

Solvent	Levels strong acids?	Levels strong bases?	Useful for
Water	yes, to H₃O⁺	yes, to OH⁻	everyday chemistry
Liquid ammonia	yes, even many weak acids become strong	no; very strong bases like NH₂⁻ survive	reactions needing very strong bases, alkali metal reductions
Glacial ethanoic acid	no; differentiates strong acids	levels bases to ethanoate	titrating very weak bases
Sulfuric acid	no	most “acids” act as bases	generating electrophiles like NO₂⁺

The general rule:

• The strongest acid that can exist in a solvent is the solvent’s characteristic cation (H₃O⁺ in water, NH₄⁺ in ammonia).
• The strongest base is the solvent’s characteristic anion (OH⁻ in water, NH₂⁻ in ammonia).

Anything stronger simply reacts with the solvent. This is why chemists choose the solvent to suit the strength of acid or base a reaction needs, not the other way round.

Aprotic solvents

Some solvents, like dimethyl sulfoxide (DMSO), acetonitrile and tetrahydrofuran (THF), don’t self-ionise significantly and don’t donate protons. In these aprotic solvents, anions are poorly solvated and become much more reactive bases. That’s why pKa values measured in DMSO can be very different from values in water: phenol, for instance, has a pKa of about 10 in water but about 18 in DMSO. Organic chemists often use DMSO pKa tables to compare very weak acids like hydrocarbons.

Molten salts and Lux–Flood acids

At high temperatures, in molten oxides and salts (such as in glass-making and metallurgy), acid–base behaviour is often described by the Lux–Flood definition:

• A base is an oxide ion donor.
• An acid is an oxide ion acceptor.

For example, in a furnace:

CaO + SiO₂ → CaSiO₃

Calcium oxide donates O²⁻ (base); silicon dioxide accepts it (acid). This is exactly what happens when limestone is added to a blast furnace to remove silica as slag, and it’s consistent with the idea that metal oxides are basic and non-metal oxides are acidic. See acidic oxides vs basic oxides.

How the definitions fit together

Definition	Acid	Base	Works in
Arrhenius	increases [H₃O⁺]	increases [OH⁻]	water
Solvent system	increases solvent cation	increases solvent anion	any self-ionising solvent
Brønsted–Lowry	proton donor	proton acceptor	any protic system
Lux–Flood	oxide ion acceptor	oxide ion donor	melts, high temperature
Lewis	electron-pair acceptor	electron-pair donor	everything above and more

The Lewis definition is broad enough to cover all the others. See Lewis acids and bases.

Key takeaways

• Many solvents self-ionise like water, each with its own characteristic acidic cation and basic anion.
• The solvent-system definition extends Arrhenius’s ideas to any such solvent.
• Basic solvents like ammonia make acids stronger; weakly basic solvents like ethanoic acid differentiate strong acids.
• The strongest acid and base possible in a solvent are its own characteristic cation and anion.
• In molten oxides, the Lux–Flood definition treats oxide-ion donors as bases and acceptors as acids.