Reactions of ions in solution

There are essentially only four basic reactions of ions in solution:

1. Acid-base reactions.
2. Precipitation reactions.
3. Complexation reactions.
4. Reduction-oxidation (redox) reactions.

Acid-base, precipitation and complexation reactions are all examples of exchange (metathesis) reactions in which ions in solution 'exchange partners', for example:


Metathesis reactions are really equilibria between the ionic species, which are displaced to the right (to the reaction product) by a feature which defines the classification of the reaction type.

Acid-base reactions The most common acid-base reactions are exemplified by the neutralization reaction between hydrogen ions and hydroxide ions; for example, the reaction between dilute hydrochloric acid and dilute sodium hydroxide:
Since water is essentially a covalent compound its formation effectively removes H⁺ and OH⁻ from the equilibrium and drives the reaction to completion. Other examples of this general type of reaction include the removal of a molecule as a gas, such as reactions of acids with carbonates and bicarbonates, where unstable H₂C0₃ decomposes to H₂0 and CO₂.

Precipitation reactions
In these reactions between ions, one substance is removed from the ionic equilibrium by precipitation (see solubility product) and drives the equilibrium to the right.

Complexation reactions A complex ion is formed by the reaction of a metal cation, in particular transition metals, with an electron donor molecule (ligand), which can be neutral or have a negative charge. The cation can accept an electron pair and the ligand donates an electron pair to form a covalent donor (co-ordinate) bond between the ligand and the metal ion. The ligands are said to coordinate with the metal ion to give a complex. Many ligands are more powerful electron donors than water and thus the addition of a ligand to an aqueous solution of a metal cation displaces the equilibrium towards the more stable complex ion. The effects of complex formation are illustrated in Box 6.3.

The overall effect of complex formation is to 'remove' a hydrated metal ion from the mixture of ions in solution by displacing the equilibrium in favour of the complex, cf. the similar process in the formation of water in acid-base titrations and precipitation reactions.

Reduction-oxidation (redox) reactions
The concepts of oxidation and reduction are defined in terms of complete electron transfer from one atom, ion or molecule of a chemical to another:

1. Chemical oxidized - chemical loses electron(s).
2. Chemical reduced - chemical gains electron(s).

This approach is generally applicable to most reactions and avoids complications of the older definitions involving hydrogen and oxygen. You should realize that if a chemical is oxidized during a reaction, then another must be reduced: oxidation and reduction always occur together. Furthermore:

1. Oxidizing agent - gains electron(s) and is therefore reduced.
2. Reducing agent -loses electron(s) and is therefore oxidized.

The following reaction between magnesium metal and dilute acid illustrates these concepts:

Magnesium metal has lost two electrons in forming MgMg²⁺ ions and has therefore been oxidized. The two protons have each gained an electron to form hydrogen atoms (and then one hydrogen molecule) and have been reduced. Since magnesium metal has been oxidized, it is a reducing agent and because H+ has gained an electron, it is an oxidizing agent.

The stoichiometry of a redox reaction is defined by the number of electrons transferred between the oxidizing agent and the reducing agent since the number of electrons lost by the reducing agent must equal the number of electrons gained by the reducing agent, e.g.


So that you can work out titrations involving redox reactions, you will find it necessary to balance redox equations, and while it is easy for simple reactions such as those above, more complex redox reactions, such as the one below, require more thought and work.


Such problems can be broken down into several simple steps, each with its own set of rules:

• Identify the atoms, IOns or molecules which have been oxidized and reduced.
• Identify the ionic half-reactions for the species being oxidized and reduced and combine them.
• Balance the ionic half-reactions and combine them to give a balanced equation for the reaction.

The species which are oxidized and reduced can be identified using the concept of oxidation numbers. The rules for determining oxidation numbers and examples are given in Box 6.4 and the application of ionic half-reactions to balance redox equations is shown in Box 6.5. Note that the result of the use of partial ionic equations gives a balanced ionic equation for the redox reaction.

*Note: In simple acid-base, precipitation and complexation reactions, no change of oxidation number occurs at any of the atoms involved.