2.1.4 - Acids and Bases


An acid is a chemical that releases protons, or hydrogen ions (H$\small^{+}$) into solution when dissolved in water. They have a pH less than 7.0 at 298K and can be neutralised using a base or alkali.

Commons compounds that become acids when mixed with water:

  • Hydrogen chloride:
    • Dissolves to become hydrochloric acid, HCl
    • Monoprotic, meaning it releases only 1 proton
    • Strong acid
  • Hydrogen sulphate:
    • Dissolves to become sulfuric acid, H$\small_{2}$SO$\small_{4}$
    • Diprotic (2 protons)
    • Strong acid
  • Hydrogen nitrate:
    • Dissolves to become nitric acid, HNO$\small_{3}$
    • Monoprotic
    • Weak acid
  • Hydrogen phosphate:
    • Dissolves to become phosphoric acid, H$\small_{3}$PO$\small_{4}$
    • Triprotic (3 protons)
    • Weak acid
  • Ethanoic anhydride:
    • Reacts with water to produce ethanoic acid, CH$\small_3$COOH
    • Monoprotic
    • Weak acid


A base is a chemical that accepts protons. A base that is dissolved in water is known as an alkali, which release hydroxide ions (OH$\small^{-}$) when dissolved in water. They have a pH of greater than 7.0 at 298K and can be neutralised using an acid.

Common bases include:

  • Ammonia, NH$\small_3$:
    • Has a lone pair on the nitrogen which can be used to form a covalent bond with a proton, thus ‘accepting’ it
    • Can dissolve in water to from ammonium ions and hydroxide ions (ammonium hydroxide), NH$\small_4$OH
    • Weak base
  • Sodium hydroxide, NaOH, a strong base
  • Potassium hydroxide, KOH, a strong base

Strong and Weak Acids or Bases

A strong acid or base is defined as one that fully dissociates into its ions when dissolved in water, whereas a weak acid or base is defined as one that only partially dissociates into its ions when dissolved in water.

The ionic equations of a strong acid and strong base are shown below as complete equations with only the right hand side produced:

$$ \text{H$_2$SO$_4$} _\text{(aq)} \rightarrow \text{2H$^+$} _\text{(aq)} + \text{SO$_4^{2-}$} _\text{(aq)} $$

$$ \text{NaOH} _\text{(aq)} \rightarrow \text{Na$^+$} _\text{(aq)} + \text{OH$^-$} _\text{(aq)} $$

On the other hand, the ionic equations for weak acids or bases are not complete, and instead are shown as equilibria, as both the left and right hand sides are produced:

$$ \text{CH$_3$COOH} _\text{(aq)} \leftrightharpoons \text{H$^+$} _\text{(aq)} + \text{CH$_3$COO$^-$} _\text{(aq)} $$

$$ \text{NH$_4$OH} _\text{(aq)} \leftrightharpoons \text{NH$_4^+$} _\text{(aq)} + \text{OH$^-$} _\text{(aq)} $$


When an acid and a base or alkali are mixed together, they react to form a salt and water, with carbonates as bases also producing carbon dioxide. This is neutralisation, and produces a solution with pH equal to 7.0 at 298K. For every neutralisation reaction, the ionic equation is:

$$ \text{H$^+$} _\text{(aq)} + \text{OH$^-$} _\text{(aq)} \rightarrow \text{H$_2$O} _\text{(l)} $$

while the spectator ions are the only things that change.

As well as water, a salt is produced, which is a pH neutral ionic compound made up of a metal and a non-metal (except an ammonium salt, where both are non-metals). The metal comes from the base and the non-metal comes from the acid. For example:

$$ \text{HCl} _\text{(aq)} + \text{NaOH} _\text{(aq)} \rightarrow \text{NaCl} _\text{(aq)} + \text{H$_2$O} _\text{(l)} $$


$$ \text{2HCl} _\text{(aq)} + \text{Na$_2$O} _\text{(aq)} \rightarrow \text{2NaCl} _\text{(aq)} + \text{H$_2$O} _\text{(l)} $$

the Cl$\small^-$ anion from the HCl and the Na$\small^+$ cation from the NaOH or Na$\small_2$O combine to form NaCl.

Another type of base is a carbonate, which contains the carbonate ion (CO$\small_3^{2-}$). When it is involved in a neutralisation reaction, carbon dioxide (CO$\small_{2}$) is formed.

$$ \text{2HCl} _\text{(aq)} + \text{Na$_2$CO$_3$} _\text{(aq)} \rightarrow \text{2NaCl} _\text{(aq)} + \text{H$_2$O} _\text{(l)} + \text{CO$_2$} _\text{(g)} $$

Acid-Base Titrations

As seen in the ionic equation above, one mole of H$\small^+$ ions and one mole of OH$\small^-$ ions together form one mole of water. If neither ion is in excess, i.e., the number of hydrogen and hydroxide ions are exactly equal, a solution is perfectly neutralised.

An acid-base titration is an experiment to determine the exact volume of acid/base that is needed to neutralise a known volume of the other. This is usually done to find the concentration of one of the solutions. An example of a calculation was shown in notes of 2.1.3 - Amount of Substance.

Standard Solutions

A solution with a known concentration is known as a standard solution. One is usually made by measuring a specific mass of pure, dry solid substance and dissolving it carefully in a specific volume of distilled water such that the substance is dissolved equally throughout the solution. An example of a calculation was shown in notes of 2.1.3 - Amount of Substance.

Below is a video with further explanation of titrations, as well as how to set up and conduct an acid-base titration, and making a standard solution: