Oxidation Numbers and Redox Reactions

• Oxidation number measures the degree of oxidation of an atom in a compound.

Rules for Assigning Oxidation Numbers

• The oxidation number of any pure element is 0.
• For monoatomic ions, the oxidation number equals the charge of the ion.
• In compounds, hydrogen is usually +1, and oxygen is usually -2.
• The sum of the oxidation numbers in a neutral compound is 0; in an ion, it equals the charge of the ion.

Oxidation vs. Reduction

• Oxidation is when an atom or ion loses electrons and its oxidation number increases.
• Key example: Fe becomes Fe³⁺ + 3e⁻, showing iron being oxidized.
• Reduction is when an atom or ion gains electrons, leading to a decrease in its oxidation number.
• Key example: MnO₄⁻ becomes Mn²⁺ + 5e⁻, showing manganese being reduced.
• LEO the lion says GER: Loss of Electrons is Oxidation, Gain of Electrons is Reduction.

Redox Reactions

• Redox reactions involve simultaneous reduction of one species and oxidation of another.

Steps to Identify Redox Reactions

• Identify the oxidation numbers of elements in the reactants and products.
• Determine which element is oxidized (increase in oxidation number) and which is reduced (decrease in oxidation number).
• Iron is oxidized (loses electrons) in the reaction Fe → Fe³⁺ + 3e⁻.
• Mn is reduced (gains electrons) in the reaction MnO₄⁻ → Mn²⁺.

Electrochemical Cells

• Voltaic (Galvanic) Cells convert chemical energy into electrical energy and consist of two half-cells connected by a wire and a salt bridge.
• Anode is the site where oxidation occurs (electrons are lost).
• Cathode is the site where reduction occurs (electrons are gained).
• Electrolytic Cells use electrical energy to drive a non-spontaneous chemical reaction.
• Anode is positive (oxidation occurs here).
• Cathode is negative (reduction occurs here).

Standard Electrochemical Potentials

• Standard Electrode Potential is the tendency of a half-reaction to occur as a reduction, measured under standard conditions.
• The more positive the electrode potential, the more likely the substance will be reduced.
• The more negative the electrode potential, the more likely the substance will be oxidized.

Autooxidation

• Autooxidation is a spontaneous oxidation reaction, often with oxygen, where a substance is oxidized without an external oxidizing agent.
• Key example: When fats or oils are exposed to air, they undergo autooxidation, leading to rancidity.

Electroplating

• Electroplating uses electrolysis to deposit a layer of metal onto a surface.
• The metal to be plated is reduced at the cathode.
• The metal source (in ionic form) is present in the electrolyte.
• Key example: Plating silver onto a spoon.

Balancing Redox Reactions in Acidic Solution

• Write separate half-reactions for oxidation and reduction.
• Balance atoms (except hydrogen and oxygen).
• Balance oxygen by adding H₂O, and balance hydrogen by adding H⁺.
• Balance the charges by adding electrons.
• Combine the half-reactions.

Balancing Redox Reactions in Basic Solution

• After balancing in acidic conditions, add OH⁻ ions to both sides to neutralize H⁺ ions.

Key Terms and Concepts

• Oxidation Number: The charge an atom would have if the compound were ionic.
• Oxidation: Loss of electrons, increase in oxidation number.
• Reduction: Gain of electrons, decrease in oxidation number.
• Redox Reaction: A chemical reaction where one substance is oxidized and another is reduced.
• Electrochemical Cell: A device that generates electricity through redox reactions.
• Electroplating: Using electrolysis to coat an object with a thin layer of metal.