​​Acids and bases

​​In a nutshell

Acids are proton donors and bases are proton acceptors. They can be classified as strong or weak depending on how much they dissociate in a solution. They are used in many household items such as in soaps. Acids and bases can undergo a neutralisation reaction.   

Equations

These are general chemical equations. You will need to be able to use these to write equations with specific substances.

​$$acid \>+\>base\>→\>salt\>+water$$

$$acid \>+\>water\>→\>acid\>anion\>+\>hydroxonium\>ion$$

​$$base\>+\>water\>→\>base\>cation\>+\>hydroxide\>ion$$

$$acid\>→\>hydrogen\>ion\>+\>anion$$

$$base\>→\>cation\>+\>hydroxide\>ion$$

$$acid\>+\>base\>⇌\>conjugate\>base\>+conjugate\>acid$$​​​​​​

Acids, bases and protons

Brønsted-Lowry acids donate protons. When they mix with water, the Brønsted-Lowry acid donates a proton, $$H^+$$​, to water. Water turns into a hydroxonium ion.

​$$acid \>+\>water\>→\>acid\>anion\>+\>hydroxonium\>ion$$​​

​$$HA(aq)+H_2O(l)→A^-(aq)+H_3O^+(aq)$$

The Brønsted-Lowry acid, $$HA$$​, donated an $$H^+$$​ to water.

Example

$$hydrochloric\>acid \>+\>water\>→\>hydrochloric\>acid\>anion\>+\>hydroxonium\>ion$$

$$HCl(aq)+H_2O(l)→Cl^-(aq)+H_3O^+(aq)$$​​​

Brønsted-Lowry bases accept protons. When they mix with water, they accept a proton, $$H^+$$​, from water. Water turns into a hydroxide ion. 

​$$base\>+\>water\>→\>base\>cation\>+\>hydroxide\>ion$$​​

​$$B(aq)+H_2O(l)→BH^+(aq)+OH^-(aq)$$

The Brønsted-Lowry base, B, accepted an $$H^+$$​ from water.​

Example

$$ammonia\>+\>water\>→\>ammonium\>+\>hydroxide\>ion$$

$$NH_3+H_2O→NH_4^++OH^-$$​

​​

Strong or weak acids and bases

The strength of acids and bases depends on how strongly they dissociate or ionise in water. 

Strong acids

The amount of protons that an acid donates, depends on the acid. Strong acids completely dissociate. Monoprotic acids can only donate one proton, diprotic acids can donate two protons and triprotic acids can donate three protons. 

​$$acid\>→\>hydrogen\>ion\>+\>anion$$​​

Example

​$$hydrochloric\>acid\>→\>hydrogen\>ion\>+\>hydroxide\>ion$$​​

​$$HCl→H^++Cl^-$$

Hydrochloric acid is a monoprotic acid, it is a strong acid as it almost completely dissociates. 

Weak acids

Weak acids only slightly dissociates in water. This means only a small proportion of protons are formed. This forms an equilibrium that shifts more to the left (reactants).

$$acid\>→\>hydrogen\>ion\>+\>anion$$​​

Example

​$$benzoic\>acid\>→\>benzoic\>acid\>anion\>+\>hydrogen\>ion$$​​

$$C_6H_5COOH⇌C_6H_5COO^-+H^+$$​​

Benzoic acid only slightly dissociates. It is a weak acid so an equilibrium is formed. 

Strong bases

Strong bases almost completely gets dissociated in water.

​$$base\>→\>cation\>+\>hydroxide\>ion$$​​

Example

​$$lithium\>hydroxide\>→\>lithium\>cation\>+\>hydroxide\>ion$$​​

​$$LiOH→Li^++OH^-$$​​

The half equation of lithium hydroxide is shown. Lithium hydroxide is almost completely ionised. 

Weak bases

Weak bases only slightly get ionised in water. This forms an equilibrium that shifts more to the left (reactants).

​$$base\>+\>water\>→\>base\>cation\>+\>hydroxide\>ion$$​​

Example

​$$ammonia\>+\>water\>⇌\>ammonium\>+\>hydroxide\>ion$$​​

​$$NH_3+H_2O⇌NH_4^++OH^-$$​​

Ammonia is a weak base that only gets slightly ionised. 

Conjugate pairs

In order for acids to donate protons and bases to accept protons, they need a species to donate or accept this proton from. Previously, water was shown and it acted as either an acid or a base depending on what it reacted with. This means water is amphiprotic. Acids and bases can react together to fulfil both their roles. An acid donates protons to the base and the base accepts the protons.

Conjugate pairs are formed when acids and bases react. They are linked by proton transfer. Conjugate bases are species that lost a proton and conjugate bases are species that gain a proton.

​$$acid\>+\>base\>⇌\>conjugate\>base\>+conjugate\>acid$$​​

​$$HA(aq)+B(aq)⇌A^-(aq)+BH^+(aq)$$

Example

$$methyl \>acetate\>+\>water\>→\>methyl\>acetate\>anion\>+\>hydroxonium\>ion$$

$$CH_3CH_2CO_2H\>+\>H_2O\>⇌\>CH_3CH_2COO^−+H_3O^+$$​​​

An equilibrium is formed. If you add more acid and/or base, the equilibrium will shift to the right according to Le Chatelier's principle. 

Water acts as a base when an acid is added. An acid will donate protons to water and water will accept the protons. This forms conjugate pairs.

$$H_3O^+$$​ is the conjugate acid and $$A^-$$​ is the conjugate base. Labels "A" and "B" are showing the conjugate pairs. 

Water acts as an acid when a base is added. A base will accept protons from water and water will donate the protons. This forms conjugate pairs.

$$OH^-$$​ is the conjugate base and $$BH^+$$​ is the conjugate acid. Labels "A" and "B" are showing the conjugate pairs. 

Neutralisation 

A neutralisation reaction is when acids and bases react to form water and a salt. When there's an equal concentration of protons and hydroxide ions, the acids and bases get completely neutralised.

​$$acid\>+\>base\>→\>salt\>+\>water$$

Example

$$hydrochloric\>acid\>+\>lithium\>hydroxide\>→\>water\>+\>lithium\>chloride$$​​

$$HCl(aq)+LiOH(aq)→H_2O(l)+LiCl(aq)$$​​

Water and a salt, lithium chloride, are formed. 

Neutralisation causes a change in enthalpy. This is called the enthalpy change of neutralisation. This reaction is always exothermic. Therefore, the enthalpy change of neutralisation will always be negative. 

The standard enthalpy change of neutralisation is defined as the enthalpy change of when acid and base solutions react together, to produce one mole of water, under standard conditions.

Strong acids and strong bases can fully dissociate and ionise in solution. They don't have a dissociation enthalpy, there's only enthalpy for the reaction between the protons and hydroxide ions. This means the standard enthalpy of neutralisation for strong acids and strong bases is similar. 

Weak acids and weak bases don't completely dissociate in solution. This means that the reaction is reversible. Due to the small amount of protons and hydroxides being produced, they get used up very quickly. Acids or bases are therefore constantly dissociating to compensate the loss of protons or hydroxide ions meaning enthalpy is needed for dissociation. The enthalpy of dissociation varies depending on the acid and base used. Therefore, the standard enthalpy change of neutralisation changes depending on which weak acid and weak base is used.