Definitions [2]
Any reaction that involves both oxidation and reduction occurring simultaneously is called an oxidation-reduction reaction or simply a redox reaction.
or
The chemical reaction in which both oxidation and reduction occur simultaneously is called a redox reaction.
Oxidation number (also called oxidation state) is the charge that an atom of an element appears to have when present in a combined state with other atoms. It is a hypothetical charge assigned by assuming all bonds are ionic — atoms in real molecules like H₂O do not actually carry these charges.
Key Points
Redox Reactions:
- A substance that oxidises another substance (and is itself reduced) is called an oxidising agent.
- A substance that reduces another substance (and is itself oxidised) is called a reducing agent.
What is Oxidation and Reduction?
| Perspective | Oxidation | Reduction |
|---|---|---|
| In terms of oxygen | Gain of one or more O atoms | Loss of one or more O atoms |
| In terms of hydrogen | Loss of hydrogen | Gain of hydrogen |
| In terms of electropositive element | Loss of electropositive element | Gain of electropositive element |
| In terms of electronegative element | Gain of electronegative element | Loss of electronegative element |
| In terms of electrons | Loss of electrons | Gain of electrons |
| In terms of oxidation number | Increase in oxidation number | Decrease in oxidation number |
Redox in Terms of Electron Transfer:
A reaction in which electrons are lost by one substance and gained by another is called a redox reaction.
- Oxidising agent = electron acceptor
- Reducing agent = electron donor
Example:
(Hg₂²⁺ gains electrons → reduced; Sn²⁺ loses electrons → oxidised)
- Combination reaction — Two atoms or molecules combine to form a product
- Decomposition reaction — A compound breaks down into simpler substances
- Displacement reaction — An atom or ion displaces another atom or ion from a compound
- Disproportionation reaction — A single substance undergoes both oxidation and reduction simultaneously
Rules for Assigning Oxidation Numbers:
| Species | Rule |
|---|---|
| Free elements | Oxidation number = 0 (e.g., Na, O₂, O₃, Hg, S₈, P₄) |
| Monoatomic ions | Oxidation number = charge on the ion (e.g., Mn²⁺ = +2, Cr³⁺ = +3) |
| Fluorine | Always −1 in all compounds |
| Oxygen | Usually −2; Exceptions: −1 in peroxides (H₂O₂), −1/2 in superoxides (KO₂), +2 in oxygen fluoride (OF₂) |
| Hydrogen | Usually +1; Exception: −1 in metal hydrides (e.g., CaH₂, NaH) |
| Halogens (Cl, Br, I) | Usually −1 in binary compounds; can be positive when bonded to a more electronegative element or oxygen |
| Neutral compound | Sum of all oxidation numbers = 0 |
| Polyatomic ion | Sum of all oxidation numbers = charge on the ion |
- Oxidation number of N can be −3 (bonded to less electronegative atoms) or +3 (bonded to more electronegative atoms)
- Oxidation number of halogens is always −1 in metal halides
- In interhalogen compounds, the more electronegative halogen gets the oxidation number of −1
- Oxidation number of metals in amalgams and carbonyls is zero (e.g., Fe in [Fe(CO)₅] = 0)
- In complex ions, the algebraic sum of oxidation numbers of all atoms = net charge on the ion
- Oxidation number can be positive, negative, zero, a whole number, or a fraction
- Oxidation number greater than +6 or less than −4 is unusual — double-check for errors
Stock Notation
Variable oxidation states are indicated using Roman numerals in parentheses after the element symbol:
| Formula | Name | Stock Notation |
|---|---|---|
| Cu₂O | Cuprous oxide | Copper (I) oxide |
| Fe₂O₃ | Ferric oxide | Iron (III) oxide |
| HgCl₂ | Mercuric chloride | Mercury (II) chloride |
| SnCl₂ | Stannous chloride | Tin (II) chloride |
Two methods are used to balance redox reactions:
Method 1: Oxidation Number Method
The change in oxidation number is used to balance electron gain and loss.
Steps (Acidic Medium):
- Write the skeleton equation; balance all atoms except O and H first
- Identify which atoms change oxidation number; calculate the net increase and decrease
- Multiply coefficients to make total increase in oxidation number = total decrease
- Balance O atoms by adding H₂O to the side with fewer O atoms
- Balance H atoms by adding H⁺ ions
- Check that all atoms and charges are balanced
Method 2: Ion Electron Method (Half-Reaction Method)
The reaction is split into two half-reactions (oxidation and reduction) which are balanced separately and then combined.
Steps:
- Write the redox reaction in ionic form
- Split into oxidation half-reaction and reduction half-reaction
- Balance atoms in each half-reaction (except O and H first)
- Balance O by adding H₂O; balance H by adding H⁺ (acidic) or OH⁻ (basic)
- Balance charge by adding electrons to the appropriate side
- Equalise electrons transferred — multiply one or both half-reactions by suitable factors so electrons cancel
- Add both half-reactions; cancel identical species on both sides
- Check that all atoms and charges are balanced
For Basic Medium (Ion Electron Method):
After balancing in acidic conditions:
- Add OH⁻ ions to both sides equal to the number of H⁺ ions
- H⁺ + OH⁻ → H₂O (combine)
- Eliminate H₂O molecules appearing on both sides
- Final check: all elements and charges must balance
Electrochemical Cell:
An electrochemical cell converts chemical energy (from a redox reaction) into electrical energy. The classic example is the Daniell cell (Zn–Cu cell).
Key terms:
| Term | Description |
|---|---|
| Flow of electrons | Electrons travel through the external circuit from anode (−) to cathode (+) |
| Flow of current | Conventional current flows from cathode to anode (opposite to electron flow) |
| Electrode potential | The electrical potential developed at an electrode when it is in contact with its ion solution |
| Electrode reaction | The half-reaction (oxidation or reduction) that takes place at an electrode |
| Redox couple | The two chemical species (oxidised and reduced forms) linked by electron transfer, forming a half-cell |
- A large negative value of E° means the redox couple is a strong reducing agent (readily loses electrons)
- Fluorine has the highest positive E° value — it has the greatest tendency to gain electrons (strongest oxidising agent)
- A strong oxidising agent and a strong reducing agent should never be stored together — they can react violently
