In the Brønsted–Lowry model, an acid can donate a proton and a base can accept one. The same substance can sometimes behave differently depending on what it reacts with.
Where do we see it?
In water, pH gives a practical measure related to hydrogen-ion activity. A change of one pH unit represents a tenfold change in this activity, so the scale is logarithmic rather than linear.
Neutralisation does not always mean the final solution is exactly pH 7. The result depends on the acid, the base, their amounts and the chemistry of the products formed.
pH is not the whole story
Two solutions with the same pH can behave differently depending on the species present and their total concentration. pH is useful, but it is not a complete description.
Go one step further
A useful way to understand this phenomenon is to separate the driving force from the visible result. Temperature, pressure, concentration, surface area and the nature of the materials can each change the rate or the final state. In real systems, several of these factors often act at the same time.
That is why the same phenomenon can look different in a laboratory, in the kitchen or outdoors. The underlying chemistry or physics stays the same, but the conditions change the balance.
An Arrhenius acid increases hydronium in aqueous solution, while an Arrhenius base increases hydroxide. This is convenient for familiar laboratory solutions, but it is tied to water and does not describe every acid–base process.
Proton transfer creates conjugate pairs
Brønsted–Lowry chemistry tracks proton transfer: the donor is called the acid and the receiving species is the base. When HA gives H⁺ to B, the products A⁻ and BH⁺ are the conjugate base and conjugate acid. Writing both pairs makes reversible proton transfer easier to reason about.
Lewis acids and bases track electron pairs
A Lewis base donates an electron pair to form a bond; a Lewis acid accepts that pair. BF₃ reacting with NH₃ is a standard example: no proton moves, yet the donor–acceptor logic is central to coordination chemistry.
Strength, concentration and pH are different
Acid strength describes an equilibrium tendency in a given medium. Concentration describes amount per volume. pH describes hydrogen-ion activity on a logarithmic scale. Low-concentration strong acid and high-concentration weak acid solutions can therefore behave differently from simplistic labels.
Water can act on both sides
Water itself is amphiprotic: it can accept a proton from an acid or donate one to a sufficiently strong base. That is why acid–base labels belong to a reaction context rather than being permanent personality traits attached to a molecule.
Equilibrium constants compare strengths
For proton transfer HA + B ⇌ A⁻ + BH⁺, relative acid strengths help predict which side is favoured. The equilibrium tends toward the side containing the weaker acid–base pair, while quantitative work uses Ka, Kb or related constants.
Worked example
Why is NH₃ a Brønsted base in water?
Its nitrogen lone pair can accept a proton from water, producing NH₄⁺ while water becomes OH⁻.
Common traps to avoid
A strong acid is not the same thing as a concentrated acid.
The Lewis definition is broader than proton transfer, not a contradiction of it.