Redox reactions and example

Overview:

Redox reactions, short for reduction-oxidation reactions, are fundamental chemical processes in which both oxidation (loss of electrons) and reduction (gain of electrons) occur simultaneously. These reactions are commonly encountered in various chemical transformations.

Understanding Redox Reactions:

Equations for Redox Reactions:

• Redox reactions can be represented by equations that illustrate the exchange of electrons between reacting species.
• These equations are produced by combining two ion-electron equations, one representing the reduction half-step and the other the oxidation half-step.
• Balancing these ion-electron equations and adding them together results in the overall redox reaction equation.

Displacement Reactions:

• Displacement reactions are excellent examples of redox reactions.
• In these reactions, a metal higher in the electrochemical series displaces a lower metal from a solution of its ions.
• For instance, when magnesium metal is added to a solution of copper sulfate, copper metal forms on the magnesium surface, and the blue copper sulfate solution decolourises.

Example:

• Let's consider the displacement reaction between silver nitrate and zinc.

Ion-Electron Equations for the Reaction:

Oxidation Step (Zinc):

• Zn(s) → Zn²⁺(aq) + 2e⁻

Reduction Step (Silver):

• 2Ag⁺(aq) + 2e⁻ → 2Ag(s)

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Redox reactions and example

Balancing the Equations:

• Balancing is essential for combining the equations effectively.
• To balance the number of electrons, the reduction equation is multiplied by two.
• Balanced Reduction Step:
  • 2Ag⁺(aq) + 2e⁻ → 2Ag(s)

Identifying Oxidising and Reducing Agents:

• In redox reactions, it is possible to identify substances acting as oxidising agents (causing oxidation) and reducing agents (causing reduction).
• This identification is done by analysing the ion-electron equations and understanding which substances gain or lose electrons.