A2-Level Acids and Bases
Acids are species that donate protons and bases are species that accept protons.
This is the Bronsted-Lowry definition of acids and bases.
In solution, acids actually release protons to water molecules, forming hydroxonium ions (H O ).
Hydroxonium ions then react with bases, giving the extra proton to the base and reforming water.
In equations, H (aq) is used to represent a proton, assuming it is understood that the proton actually exists in solution with water as hydroxonium ions.
Acids lose protons to form their conjugate base.
Bases gain protons to form their conjugate acid.
Acids and Bases
There are different theories to describe how acids and bases work (and even what acids and bases actually are!), but in A-level chemistry the Bronsted-Lowry theory is the one most commonly referred to.
In simple terms, this theory proposes that acids donate protons (H⁺ ions) and bases accept protons.
For example, in the following:
Hydrochloric Acid + Ammonia → Ammonium Chloride + Hydrogen
Here the HCl has ‘given’ its proton to the ammonia, leaving behind a negatively charged chloride ion (Cl⁻). The ammonia has gained a proton and become a positively charged ammonium ion (NH⁺). Chloride ions and ammonium ions are attracted to each other ionically, so ammonium chloride is formed.
The HCl acts as an acid (giving away a proton) and the NH acts as a base (accepting the proton).
This seems straightforward and easy to follow, but the actual process that happens is a little more involved.
Remember everything here is aqueous (in water). Acids actually ‘give’ their protons to water molecules, not directly to a base. This forms hydroxonium ions (H O⁺). It is these hydroxonium ions that then go on to react with bases.
Let’s take hydrochloric acid and sodium hydroxide:
Remember the acid is actually giving its proton to a water molecule.
This hydroxonium ion then reacts with the hydroxide ion (OH⁻) in the solution from the sodium hydroxide (NaOH).
If we combine these equations together, the water that reacts to form the hydroxonium ion reforms at the end, so it can be removed from the equation and we are left with:
Conjugate Acids and Conjugate Bases
As defined above, the Bronsted-Lowry theory proposes that acids donate protons (H⁺ ions) and bases accept protons. In a reaction these are referred to as conjugate pairs: the acid dissociates, forming its conjugate base. In reverse, the base can accept a proton (H⁺ ion) and form its conjugate acid.