pH is the negative base-10 logarithm of the hydrogen ion concentration in moles per litre. The logarithm exists purely for convenience: hydrogen ion concentrations in ordinary solutions span roughly fourteen orders of magnitude, from about 1 mol/L in a strong acid down to about 10⁻¹⁴ mol/L in a strong base, and writing those numbers out directly would be unwieldy for everyday use. Taking a logarithm compresses that entire range into a manageable scale running roughly from 0 to 14.
Each whole unit is a factor of ten, not a fixed amount
Because pH is logarithmic, a solution at pH 3 has ten times the hydrogen ion concentration of one at pH 4, and a hundred times that of one at pH 5. This is why lemon juice at roughly pH 2 is not merely “a bit more acidic” than black coffee at roughly pH 5 — it is about a thousand times more concentrated in hydrogen ions, despite the two numbers looking close together on the scale. Misreading pH as a linear scale is one of the most common sources of confusion in introductory chemistry, and it leads people to badly underestimate how much stronger a low-pH solution actually is.
Where the number 14 actually comes from
Water self-ionises to a small degree, constantly splitting into H⁺ and OH⁻ ions and recombining. At 25 °C, the product of those two ion concentrations, [H⁺][OH⁻], is a constant equal to 1.0 × 10⁻¹⁴ — this constant is called Kw, the ion-product constant of water, and taking its negative logarithm gives exactly 14, which is where the familiar pH + pOH = 14 relationship comes from. Crucially, this is not a fixed law of nature; Kw itself changes with temperature, because the equilibrium of water’s self-ionisation shifts as temperature changes. At 50 °C, neutral water actually sits at pH 6.63 rather than 7.00 — still chemically neutral, since [H⁺] still equals [OH⁻] at that point, but numerically lower because Kw has increased. It is also possible for concentrated strong acids to have a calculated pH below 0, and concentrated strong bases to exceed pH 14 — the 0-to-14 scale is a convenient range for typical dilute solutions, not a hard mathematical boundary.
Strong acids and bases dissociate completely; weak ones do not
A strong acid like hydrochloric acid dissociates essentially completely in water, so a 0.1 mol/L solution of HCl gives [H⁺] = 0.1 mol/L directly, and therefore pH = 1 exactly. A weak acid behaves very differently: a 0.1 mol/L solution of acetic acid (the acid in vinegar) gives a pH around 2.9, not 1, because only roughly 1% of the acetic acid molecules actually ionise into H⁺ and acetate ions at that concentration — the rest remain as intact, neutral acetic acid molecules. Calculating the pH of a weak acid properly requires knowing its acid dissociation constant, Ka, and typically solving a quadratic equation derived from the equilibrium expression, rather than the simple direct calculation that works for strong acids.
Why pH matters beyond the chemistry classroom
pH governs an enormous range of practical processes: enzymes in the human body function only within narrow pH windows, which is why blood pH is tightly regulated to stay between about 7.35 and 7.45 — even a small deviation outside this range is a serious medical emergency. Soil pH determines which nutrients are available to plants and therefore what will grow well in a given location. Swimming pool pH has to be kept in a specific range both for swimmer comfort and for chlorine to work effectively as a disinfectant, since chlorine’s effectiveness itself depends heavily on the surrounding pH.