Redox Reactions

1. Classical Idea of Redox Reactions – Oxidation and Reduction

Initially, the terms "oxidation" and "reduction" had a more limited scope.

Oxidation

Classically, oxidation was defined as the addition of oxygen or another electronegative element to a substance, or the removal of hydrogen or another electropositive element from a substance.

Reduction

Reduction was defined as the opposite: the removal of oxygen or another electronegative element from a substance, or the addition of hydrogen or another electropositive element to a substance.

It was quickly realized that oxidation and reduction always occur simultaneously, leading to the coinage of the term "redox" reactions.

2. Redox Reactions in Terms of Electron Transfer Reactions

A more modern and precise definition of redox reactions is based on electron transfer. Each process can be considered as two separate steps, called half-reactions:

• Oxidation reactions: Half-reactions that involve the loss of electrons.
• Reduction reactions: Half-reactions that involve the gain of electrons.

Oxidising agent (oxidant): The species that accepts electron(s) and causes oxidation in another species. It itself gets reduced.

Reducing agent (reductant): The species that donates electron(s) and causes reduction in another species. It itself gets oxidized.

3. Competitive Electron Transfer Reactions

Redox reactions often involve a competition for electron release. When a zinc strip is placed in copper nitrate solution, zinc loses electrons to form Zn²⁺ (oxidation), and copper ions gain electrons to form metallic copper (reduction).

Conversely, placing a copper strip in zinc sulfate solution shows no visible reaction. These observations demonstrate that metals have different tendencies to release electrons. The electron-releasing tendency of metals is Zn > Cu > Ag. This concept forms the basis for developing a metal activity series or electrochemical series.

4. Oxidation Number

The oxidation number is a practical method for tracking electron shifts in chemical reactions, particularly for covalent compounds. It denotes the oxidation state of an element in a compound, determined by a set of rules, assuming that electrons in a covalent bond belong entirely to the more electronegative element.

Rules for Assigning Oxidation Numbers:

• Elements in free or uncombined state: Oxidation number is zero.
• Monatomic ions: Oxidation number equals the charge on the ion.
• Alkali metals (Group 1): Always +1 in compounds.
• Alkaline earth metals (Group 2): Always +2 in compounds.
• Oxygen: Usually -2, except in peroxides (-1), superoxides (-½), and with fluorine (+1 or +2).
• Hydrogen: Usually +1, except in metal hydrides (-1).
• Fluorine: Always -1 in all its compounds.
• The algebraic sum of oxidation numbers in a neutral compound must be zero; in a polyatomic ion, it must equal the charge on the ion.

The Stock notation system uses a Roman numeral in parenthesis after a metal's symbol to indicate its oxidation state (e.g., Au(I)Cl, Sn(IV)Cl₄).

Using oxidation numbers, we can redefine redox terms:

• Oxidation: An increase in the oxidation number.
• Reduction: A decrease in the oxidation number.

5. Types of Redox Reactions

Combination Reactions

Two or more substances combine to form a single compound (A + B → C). It's a redox reaction if at least one reactant is in its elemental form.

Decomposition Reactions

A compound breaks down into two or more components. It's a redox reaction if at least one product is in its elemental state. Not all decomposition reactions are redox reactions.

Displacement Reactions

An ion or atom in a compound is replaced by an ion or atom of another element (X + YZ → XZ + Y). This includes metal displacement and non-metal (e.g., hydrogen, halogen) displacement.

Disproportionation Reactions

A special type of redox reaction where an element in one oxidation state is simultaneously oxidised and reduced. The reacting substance must contain an element that can exist in at least three oxidation states, and the element in the reactant is in an intermediate state.

Here, oxygen in H₂O₂ (oxidation state -1) is reduced to -2 (in H₂O) and oxidized to 0 (in O₂).

6. Balancing of Redox Reactions

Two primary methods are used to balance redox equations:

• Oxidation Number Method: Balances the equation by identifying changes in oxidation numbers and ensuring the total increase equals the total decrease.
• Half-Reaction Method (Ion-Electron Method): Splits the reaction into oxidation and reduction half-reactions, balances each separately for atoms and charge (using H⁺/OH⁻ and H₂O as needed), and then combines them.

7. Redox Reactions as the Basis for Titrations

Similar to acid-base titrations, redox titrations determine the strength of a reductant or oxidant using a redox-sensitive indicator. Some reagents like the permanganate ion (MnO₄⁻) are intensely colored and act as their own indicators. In other cases, external indicators or specific indicators like starch for iodine are used.

8. Redox Reactions and Electrode Processes

Redox reactions are fundamental to electrode processes and the functioning of electrochemical cells. In a direct redox reaction, electrons are transferred directly. To harness electrical energy, an indirect transfer is set up in cells like the Daniell cell.

A redox couple consists of the oxidized and reduced forms of a substance taking part in a half-reaction (e.g., Zn²⁺/Zn). The potential difference between electrodes is the electrode potential.

The Standard Electrode Potential (E°) is the electrode potential when the concentration of each species is unity (or 1 atm for gases) and the reaction is at 298 K. By convention, the standard electrode potential of a hydrogen electrode is 0.00 V. A negative E° means the redox couple is a stronger reducing agent than the H⁺/H₂ couple, while a positive E° means it is a weaker reducing agent.