Acids and bases: strong and weak acids, pH and neutralisation

An acid is a substance that gives away hydrogen ions (protons), and a base is one that takes them up; when the two react they neutralise each other to form a salt. Acids and bases turn up everywhere in a lab. They set the pH of buffers, decide the charge on a peptide, and dissolve or precipitate a sample. This guide explains what acids and bases are, why some are "strong" and others "weak", how both relate to pH, and what happens when one neutralises the other.

For buffers, pKa and the Henderson–Hasselbalch equation, see buffers and pH. Terms are defined in the glossary.

What is an acid?

An acid is a substance that gives away hydrogen ions (H⁺) in solution. A hydrogen ion is just a proton, so chemists often say an acid is a proton donor. This is the Brønsted–Lowry definition, the one used in most lab work.

In water, the proton doesn't float about on its own. It attaches to a water molecule to form a hydronium ion (H₃O⁺). The more hydronium ions a solution holds, the more acidic it is, and the lower its pH.

Familiar laboratory acids include hydrochloric acid (HCl), acetic acid (the acid in vinegar), trifluoroacetic acid (TFA) and phosphoric acid.

What is a base in chemistry?

A base is the opposite: a substance that accepts hydrogen ions, a proton acceptor. Many bases work by supplying hydroxide ions (OH⁻), which combine with H⁺ to make water. Others, such as ammonia, take a proton directly.

A base that dissolves in water is called an alkali. Sodium hydroxide and ammonia solution are common examples. Solutions of bases have a pH above 7.

Some substances can act as either. Water is one: it can give a proton or take one. Amino acids are another, because they carry both an acidic carboxyl group and a basic amine group. At neutral pH they exist as a zwitterion, with one positive and one negative charge.

Three definitions: Arrhenius, Brønsted–Lowry and Lewis

Chemists use three definitions, each wider than the last:

The Arrhenius definition only works in water. Brønsted–Lowry covers bases such as ammonia that don't contain hydroxide, and it is the one used in most lab work (OpenStax Chemistry 2e, 14.1). The Lewis definition also covers reactions with no proton at all (OpenStax Chemistry 2e, 15.2).

Conjugate pairs

When an acid gives up its proton, what remains can take a proton back, so it is a base. It is called the acid's conjugate base. Acetic acid and acetate are a conjugate pair, as are ammonium and ammonia. The buffers guide has a table of common pairs.

Acids, bases and pH

pH is a measure of the hydrogen ion concentration. It is the negative logarithm of that concentration, so each step of one pH unit is a tenfold change.

Water itself splits very slightly into H⁺ and OH⁻. At 25 °C the two concentrations multiply to a fixed value, which is why pH and pOH (the same scale for hydroxide) always add up to 14. Neutral means the two ions are equal, which is pH 7 only at 25 °C. Warmer water is still neutral at a slightly lower pH.

To measure pH, see pH meters or, for a quick check, pH indicators.

Strong and weak acids

"Strong" and "weak" describe how completely an acid gives up its protons in water. They say nothing about how concentrated it is.

• A strong acid ionises completely. Every molecule gives up its proton. Hydrochloric, nitric and sulfuric acid (its first proton) are strong acids.
• A weak acid ionises only partly. Most molecules keep their proton, and the solution settles at a balance between the two forms. Acetic acid, phosphoric acid and the side chains of aspartic and glutamic acid are weak acids.

The pKa measures how weak an acid is. It is the pH at which half the acid has given up its proton. The lower the pKa, the stronger the acid.

The TFA figure is from the TFA guide, which explains why so much of it ends up in synthetic peptide powders.

The same concentration, a different pH

Because a weak acid ionises only partly, it gives a higher pH than a strong acid at the same concentration:

• 0.1 M hydrochloric acid: fully ionised, so [H⁺] is 0.1 M and the pH is 1.0.
• 0.1 M acetic acid: only about 1.3% ionised, so the pH is about 2.9.

That is why 0.6% acetic acid in water, which is about 0.1 M, is mild enough to use as a solvent for basic peptides. See acetic acid as a peptide solvent.

Strong vs concentrated

A dilute solution of a strong acid can be less acidic than a concentrated solution of a weak one. "Strong" is about the molecule. "Concentrated" is about how much of it is in the water. A label such as "0.1% TFA" tells you the concentration; the acid's identity tells you its strength.

Examples of common acids and bases

Strong and weak bases

Bases follow the same pattern.

• Strong bases ionise completely. The hydroxides of sodium, potassium and the heavier alkaline-earth metals are the common ones (Brown et al.).
• Weak bases take up protons only partly. Ammonia, Tris and the amine groups in peptides are weak bases.

For a weak base, the pKa quoted is usually that of its conjugate acid. Tris buffer, for example, is described by the pKa of protonated Tris, which is about 8.1.

Neutralisation

Neutralisation is the reaction between an acid and a base that forms a salt and, usually, water. The hydrogen ions and hydroxide ions combine to form water, and the other ions are left behind as the salt.

Acid + base → salt + water

For example, hydrochloric acid and sodium hydroxide give sodium chloride and water:

HCl + NaOH → NaCl + H₂O

Sulfuric acid has two protons to give, so it needs two sodium hydroxides; see balancing chemical equations for how to get the numbers right. Ammonia takes the proton directly, so that reaction makes a salt without water.

Neutralisation gives out heat. For a strong acid and a strong base, the heat released is close to the same for every pair, because the reaction is always H⁺ + OH⁻ → H₂O. See energy in reactions.

"Neutralised" doesn't always mean pH 7. A weak acid neutralised by a strong base leaves a solution that is slightly basic, because the conjugate base takes a few protons back from water. This matters when choosing an indicator for a titration; see pH indicators.

Salts

A salt is the ionic compound left after neutralisation. Not all salts are neutral in water:

Synthetic peptides are often supplied as salts too. A peptide with basic groups pairs with the anion of the acid used to purify it, such as trifluoroacetate or acetate. That is why a peptide supplied as a TFA salt gives a slightly acidic solution in water. See TFA counter-ions and salt exchange.

Acids and bases in peptide work

• Charge. Each acidic and basic group on a peptide gains or loses a proton depending on pH. That sets the peptide's net charge and its isoelectric point.
• Solubility. A peptide dissolves best away from its isoelectric point. Dilute acid helps basic peptides, and dilute base helps acidic ones. See peptide solubility.
• Stability. Strongly acidic or basic conditions speed up hydrolysis of peptide bonds and other breakdown reactions.
• Analysis. HPLC mobile phases for peptides usually contain about 0.1% TFA or formic acid. See HPLC and LC-MS solvents.

Handling acids and bases

Concentrated acids and bases are corrosive and carry the corrosive hazard symbol. Add acid to water, not water to acid, because the heat released can boil the water and spit acid. Work in a fume cupboard with volatile acids such as concentrated hydrochloric acid or TFA. Store acids and bases apart, and follow the safety data sheet for each. See COSHH and safety data sheets and storing lab chemicals safely.

Related guides

For what a solution is and how things dissolve, see solvents, solutes and solutions. For buffers that hold pH steady, see buffers and pH. For measuring pH accurately, see pH meters. For the colour-change dyes used in titrations, see pH indicators. Every lot has its own COA, published on the product page and on that lot's own page. See Quality.