2.1.4 - Acids and Bases

Acids

An acid is a chemical that releases protons, or hydrogen ions (H${}^{+}$) 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${}_2$SO${}_4$
  • Diprotic (2 protons)
  • Strong acid
• Hydrogen nitrate:
  • Dissolves to become nitric acid, HNO${}_3$
  • Monoprotic
  • Weak acid
• Hydrogen phosphate:
  • Dissolves to become phosphoric acid, H${}_3$PO${}_4$
  • Triprotic (3 protons)
  • Weak acid
• Ethanoic anhydride:
  • Reacts with water to produce ethanoic acid, CH${}_3$COOH
  • Monoprotic
  • Weak acid

Bases

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${}^{-}$) 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${}_3$:
  • 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${}_4$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:

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

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

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:

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

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

Neutralisation

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:

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

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:

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

or:

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

the Cl${}^{-}$ anion from the HCl and the Na${}^{+}$ cation from the NaOH or Na${}_2$O combine to form NaCl.

Another type of base is a carbonate, which contains the carbonate ion (CO${}_3^{2-}$). When it is involved in a neutralisation reaction, carbon dioxide (CO${}_2$) is formed.

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

Acid-Base Titrations

As seen in the ionic equation above, one mole of H${}^{+}$ ions and one mole of OH${}^{-}$ 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: