Stomach acid can dissolve a steak; the inside of a cleaning bottle can dissolve grease. Opposite as they seem, both are the same story told from two sides: acids give away protons and bases take them.
This article covers what acids and bases are, the two classic definitions, how the pH scale works, why “strong” is not the same as “concentrated”, and what happens when an acid meets a base.
Two ways to define them
Arrhenius (1884): an acid releases H⁺ ions in water (HCl → H⁺ + Cl⁻) and a base releases OH⁻ ions (NaOH → Na⁺ + OH⁻). Simple and good for most lab work, but it only works in water.
Brønsted–Lowry (1923): an acid is a proton donor and a base is a proton acceptor. This explains ammonia, which has no OH⁻ yet is clearly basic: NH₃ + H₂O ⇌ NH₄⁺ + OH⁻. Every acid has a conjugate base (what is left after it gives away H⁺), and every base a conjugate acid.
The pH scale
Hydrogen-ion concentrations span many powers of ten, so chemists use a logarithm: pH = −log[H⁺]. Pure water at 25 °C has [H⁺] = 1 × 10⁻⁷ mol/L, so its pH is 7. Each step of one pH unit is a tenfold change: pH 3 is ten times more acidic than pH 4 and a thousand times more than pH 6.
The partner quantity is pOH = −log[OH⁻], and in water at 25 °C pH + pOH = 14. The pH calculator converts between all four quantities.
Strong versus weak
Strength is about how completely an acid ionizes, not how much of it there is. Hydrochloric, nitric and sulfuric acids are strong: in water, essentially every molecule gives up its proton. Acetic acid (vinegar), citric acid (lemons) and carbonic acid (soda) are weak: only a small fraction ionizes, and the rest stays as whole molecules in equilibrium.
That is why a dilute strong acid can have a higher pH than a concentrated weak acid. The weak-acid equilibrium is measured by Ka, or pKa = −log Ka: the lower the pKa, the stronger the acid.
Neutralization
When an acid and a base react, the H⁺ and OH⁻ combine into water and the leftover ions form a salt:
Neutralization releases heat and is the basis of titration: adding a base of known concentration drop by drop until an indicator changes colour tells you exactly how much acid was present. Antacids, soil liming and treating acid spills all use the same reaction.
Acids and bases around you
- Your stomach keeps hydrochloric acid at pH 1.5–3.5 to digest proteins.
- Blood is held between pH 7.35 and 7.45 by buffers; outside that range you become seriously ill.
- Ocean acidification: CO₂ dissolving into seawater forms carbonic acid and has lowered surface pH by about 0.1 since the industrial revolution — roughly 26% more H⁺.
- Soap is made by reacting fats with a strong base, a process called saponification.
Common misconceptions
“All acids are dangerous.”
Orange juice, yogurt and vitamin C are acidic. Danger depends on strength, concentration and quantity — and strong bases like drain cleaner can burn just as badly.
“pH can only go from 0 to 14.”
That range covers ordinary solutions. Very concentrated strong acids can have negative pH, and the neutral point itself shifts with temperature.
Frequently asked questions
What is the difference between an acid and a base?
An acid donates hydrogen ions (protons, H⁺) in a reaction; a base accepts them. In water, acids raise the H⁺ concentration and bases raise the OH⁻ concentration.
What does pH mean?
pH is the negative base-10 logarithm of the hydrogen ion concentration: pH = −log[H⁺]. At 25 °C, pH 7 is neutral, below 7 is acidic and above 7 is basic.
What is a strong acid?
A strong acid ionizes completely in water, like hydrochloric acid (HCl). A weak acid, like acetic acid in vinegar, ionizes only partly.
What happens when an acid reacts with a base?
They neutralize each other, forming water and a salt. For example, HCl + NaOH → NaCl + H₂O.
References and further reading
- Arrhenius, S. (1887) — On the dissociation of substances dissolved in water.
- Brønsted, J. N.; Lowry, T. M. (1923) — the proton-transfer definition of acids and bases.
- IUPAC — Gold Book: pH, acid, base, pKa.
- NOAA — Ocean acidification overview.