A metal ion \(\mathrm{M}^{3+}\) loses three electrons to produce another cation. The oxidation number of the metal in the cation is: (a) \(+3\) (b) \(+4\) (c) \(-3\) (d) \(+6\)

Redox reactions, short for reduction-oxidation reactions, are chemical processes that involve the transfer of electrons between two substances. In these reactions, one substance undergoes oxidation by losing electrons, and another substance undergoes reduction by gaining electrons.

Understanding redox reactions is essential because they play a vital role in various biological processes, industrial applications, and even in everyday life like in batteries. Oxidation and reduction always occur simultaneously because electrons are transferred from the oxidizing agent to the reducing agent, which connects the two processes in what is known as a redox couple.

The Role of Electron Transfer

Electrons are the currency of chemical reactions and in redox reactions, they are moved from one atom or molecule to another. The loss or gain of these electrons alters the overall charge of the substances involved, leading to changes in their chemical properties and behaviors.

Identifying Oxidation and Reduction

In order to identify which species is oxidized and which is reduced, one must track the movement of electrons. A species that loses electrons is said to be oxidized, while the one that gains electrons is reduced.

Chemical oxidation refers to the process where a substance loses electrons and increases its oxidation state. This can involve the interaction with oxygen, hence the name 'oxidation', but it isn't limited to it; any electron loss, regardless of the electron acceptor, is considered oxidation.

Oxidation can be part of a redox reaction or it can occur on its own in a process called disproportionation. Disproportionation is a type of redox reaction where a single substance undergoes both oxidation and reduction, transforming into two different products.

Examples of Oxidation

Common examples of oxidation include the rusting of iron, the browning of fruit, and the burning of fuels. In each of these examples, oxidation causes significant changes in the chemical composition and properties of the substance.

The oxidation state, or oxidation number, indicates the degree of oxidation of an atom in a compound or ion based on the hypothetical charge it would have if all bonds were considered ionic. Calculating oxidation states is a fundamental skill in chemistry that helps us understand the electron distribution within molecules and ions.

General rules to calculate oxidation states include the assignment of 0 to elements in their standard state, and -2 to oxygen in most compounds (except for peroxides and superoxides). Fluorine always has an oxidation state of -1 due to its high electronegativity. For ions, the oxidation state is equal to the charge of the ion.

Applying Oxidation State Rules

For example, in a compound like H2O, oxygen has an oxidation state of -2, while hydrogen has an oxidation state of +1. The sum of the oxidation states of all atoms in a neutral molecule must be zero, which helps when calculating the oxidation states of less obvious atoms in more complex molecules.