Question

a. Distinguish between the processes of oxidation and reduction. b. Write an equation to illustrate each process.

Short answer

Oxidation involves electron loss (e.g., \(\text{Mg} \rightarrow \text{Mg}^{2+} + 2e^{-}\)), while reduction involves electron gain (e.g., \(\text{O}_2 + 4e^{-} \rightarrow 2\text{O}^{2-}\)).

Step-by-step solution

Define Oxidation

Oxidation is a chemical reaction in which a substance loses electrons. This can also be described as an increase in the oxidation state of the element involved. For example, when magnesium reacts with oxygen to form magnesium oxide, magnesium undergoes oxidation.

Provide an Oxidation Equation

An example of an oxidation reaction can be represented by the equation: \[ \text{Mg} \rightarrow \text{Mg}^{2+} + 2e^{-} \]In this reaction, magnesium (Mg) loses two electrons to form a magnesium ion (\(\text{Mg}^{2+}\)).

Define Reduction

Reduction is the process in which a substance gains electrons. This can also be described as a decrease in the oxidation state of the element involved. For instance, when oxygen gains electrons to form oxide ions during the formation of magnesium oxide, oxygen undergoes reduction.

Provide a Reduction Equation

An example of a reduction reaction can be represented by the equation: \[ \text{O}_2 + 4e^{-} \rightarrow 2\text{O}^{2-} \]In this reaction, oxygen (\(\text{O}_2\)) gains four electrons to form oxide ions (\(\text{O}^{2-}\)).

Key concepts

Oxidation Reaction

An oxidation reaction is a chemical process where a substance loses electrons. This is often referred to as an increase in the oxidation state. Imagine it as a substance getting rid of its extra electrons, kind of like losing weight. A common example is when magnesium reacts with oxygen to form magnesium oxide.

In the reaction, the magnesium atom loses two electrons, making its oxidation state go from 0 (neutral atom) to +2 (ion form). This can be illustrated with this equation: \[ \text{Mg} \rightarrow \text{Mg}^{2+} + 2e^{-} \] Here, magnesium (Mg) loses two electrons (2e⁻) and becomes a magnesium ion (\text{Mg}^{2+}).

Oxidation reactions are crucial in many biological and chemical processes, such as respiration and combustion.

• Respiration: Cells use oxygen to convert glucose into energy.
• Combustion: Burning fuels like gasoline releases energy.

Reduction Reaction

A reduction reaction is the exact opposite of an oxidation reaction. In this process, a substance gains electrons, and its oxidation state decreases. Think of it as adding electrons to make the substance more 'electron-rich'.

For example, in the formation of magnesium oxide, oxygen undergoes reduction. It gains electrons to form oxide ions. The equation for this reaction is: \[ \text{O}_2 + 4e^{-} \rightarrow 2\text{O}^{2-} \] In this equation, oxygen (\text{O}_2) gains four electrons (4e⁻) to form two oxide ions (\text{O}^{2-}).

Reduction reactions are essential in various chemical processes, such as:

• Photosynthesis: Plants convert carbon dioxide and water into glucose and oxygen.
• Metal extraction: Extracting metals from their ores generally involves reduction.

Electron Transfer

Electron transfer is at the heart of both oxidation and reduction reactions. It's simply the movement of electrons from one substance to another. When an atom loses electrons, those electrons have to go somewhere, and that's where electron transfer comes in.

In an oxidation reaction, electrons are lost from one substance. These electrons are then gained by another substance in a reduction reaction. This transfer creates a balanced overall reaction called a redox (reduction-oxidation) reaction.

• In magnesium oxide formation, oxidation and reduction occur together. Magnesium loses electrons (oxidized) while oxygen gains them (reduced).
• Redox reactions are fundamental in energy transfer processes, like in batteries where chemical energy is converted to electrical energy.

Understanding electron transfer helps us grasp how both oxidation and reduction reactions are interlinked and essential for various life and industrial processes.