Redox Fundamentals

Redox = Reduction + Oxidation (simultaneous processes)

Key Definitions:

Oxidation: Loss of electrons (oxidation number increases)
Reduction: Gain of electrons (oxidation number decreases)
Oxidizing agent: Causes oxidation, gets reduced itself
Reducing agent: Causes reduction, gets oxidized itself

Memory Device: “OIL RIG”

Oxidation Is Loss (of electrons)
Reduction Is Gain (of electrons)

Oxidation Numbers: Complete Guide

Definition

Oxidation number (oxidation state) = hypothetical charge an atom would have if all bonding electrons were assigned to the more electronegative atom.

Fundamental Rules

Rule 1: Free Elements

All atoms in elemental form = 0
Examples: $N a^{0}$ , $C l_{2}^{0}$ , $O_{2}^{0}$ , $P_{4}^{0}$ , $S_{8}^{0}$

Rule 2: Monatomic Ions

Oxidation number = ion charge
Examples: $N a^{+ 1}$ , $C l^{- 1}$ , $C a^{+ 2}$ , $O^{- 2}$ , $A l^{+ 3}$

Rule 3: Group Trends (Most Compounds)

Group 1 (alkali metals): Always +1
Group 2 (alkaline earth): Always +2
Group 17 (halogens): Usually -1 (except when bonded to O or more electronegative halogen)

Rule 4: Hydrogen

With nonmetals: +1 ( $H Cl$ , $H_{2} O$ , $N H_{3}$ )
With metals (hydrides): -1 ( $N a H$ , $C a H_{2}$ , $L i H$ )

Rule 5: Oxygen

Most compounds: -2 ( $H_{2} O$ , $C O_{2}$ , $S O_{4}^{2 -}$ )
Peroxides (O-O bond): -1 ( $H_{2} O_{2}$ , $N a_{2} O_{2}$ )
Superoxides: -1/2 ( $K O_{2}$ )
With fluorine: +2 ( $O F_{2}$ )

Rule 6: Sum Rule

Neutral compounds: Sum = 0
Polyatomic ions: Sum = ion charge

Advanced Cases and Exceptions

Fractional Oxidation Numbers

Some compounds have average oxidation numbers:

$F e_{3} O_{4}$ (magnetite):

Contains both $F e^{+ 2}$ and $F e^{+ 3}$
Average Fe oxidation number = +8/3
Actually: $F e^{+ 2} F e_{2}^{+ 3} O_{4}$

Superoxides ( $O_{2}^{-}$ ):

$K O_{2}$ : $K^{+ 1}$ , $O^{- 1/2}$
Each O atom in $O_{2}^{-}$ has -1/2 oxidation number

Peroxides vs Normal Oxides

Compound	O Oxidation Number	Structure
$H_{2} O$	-2	H-O-H
$H_{2} O_{2}$	-1	H-O-O-H
$N a_{2} O$	-2	$N a_{2} O$
$N a_{2} O_{2}$	-1	$N a_{2} [O - O]$

Complex Ions and Coordination Compounds

$[C r (N H_{3})_{6}] C l_{3}$ :

$N H_{3}$ is neutral: 0
$C l^{-}$ : -1 each, total -3
Cr must be +3 to balance

$K_{4} [F e (CN)_{6}]$ :

$K^{+}$ : +1 each, total +4
$C N^{-}$ : -1 each, $6 \times (- 1) = - 6$
Fe: $+ 4 + (- 6) = - 2$ , so Fe is +2

Systematic Approach for Complex Molecules

Step-by-Step Method:

Example: $H_{2} S O_{4}$

Identify known oxidation numbers:
- H: +1 (Rule 4)
- O: -2 (Rule 5)
Set up equation:
- Let S = x
- $2 (+ 1) + x + 4 (- 2) = 0$ (neutral compound)
Solve:
- $2 + x - 8 = 0$
- $x = + 6$
- S oxidation number = +6

Polyatomic Ion Example: $S O_{4}^{2 -}$

Known: O = -2
Equation: $x + 4 (- 2) = - 2$ (ion charge)
Solve: $x - 8 = - 2$ , so $x = + 6$
Result: S = +6

Common Oxidation Numbers by Element

Element	Common Oxidation Numbers	Examples
Carbon	-4, -2, 0, +2, +4	$C H_{4}$ (-4), $CO$ (+2), $C O_{2}$ (+4)
Nitrogen	-3, -2, -1, 0, +1, +2, +3, +4, +5	$N H_{3}$ (-3), $N_{2} O$ (+1), $NO$ (+2), $N O_{2}$ (+4), $H N O_{3}$ (+5)
Sulfur	-2, 0, +2, +4, +6	$H_{2} S$ (-2), $S O_{2}$ (+4), $H_{2} S O_{4}$ (+6)
Chlorine	-1, 0, +1, +3, +5, +7	$H Cl$ (-1), $Cl O^{-}$ (+1), $Cl O_{3}^{-}$ (+5), $Cl O_{4}^{-}$ (+7)
Manganese	+2, +3, +4, +6, +7	$M n^{2 +}$ (+2), $M n O_{2}$ (+4), $M n O_{4}^{-}$ (+7)
Chromium	+2, +3, +6	$C r^{2 +}$ (+2), $C r^{3 +}$ (+3), $C r O_{4}^{2 -}$ (+6)

Balancing Methods

Method 1: Half-Reaction Method

Best for: Complex ionic reactions, especially in acidic/basic solutions

See detailed steps in [#Method 1 Examples]] below.

Method 2: Oxidation Number Method

Best for: Tracking electron transfer, understanding redox process

Uses Oxidation Numbers: Complete Guide to balance equations systematically.

Method 3: Inspection Method

Best for: Simple redox reactions with obvious stoichiometry

Method 1 Examples: Half-Reaction Method

Acidic Solution Example

Balance: $C r_{2} O_{7}^{2 -} + F e^{2 +} \to C r^{3 +} + F e^{3 +}$ (acidic)

Step 1: Write half-reactions

Reduction: $C r_{2} O_{7}^{2 -} \to C r^{3 +}$
Oxidation: $F e^{2 +} \to F e^{3 +}$

Step 2: Balance reduction half-reaction

$C r_{2} O_{7}^{2 -} \to 2 C r^{3 +} (balance Cr)$

$C r_{2} O_{7}^{2 -} \to 2 C r^{3 +} + 7 H_{2} O (balance O with H_{2} O)$

$C r_{2} O_{7}^{2 -} + 14 H^{+} \to 2 C r^{3 +} + 7 H_{2} O (balance H with H^{+})$

$C r_{2} O_{7}^{2 -} + 14 H^{+} + 6 e^{-} \to 2 C r^{3 +} + 7 H_{2} O (balance charge)$

Step 3: Balance oxidation half-reaction

$F e^{2 +} \to F e^{3 +} + e^{-}$

Step 4: Equalize electrons and combine

$C r_{2} O_{7}^{2 -} + 14 H^{+} + 6 e^{-} \to 2 C r^{3 +} + 7 H_{2} O (\times1)$ $6 F e^{2 +} \to 6 F e^{3 +} + 6 e^{-} (\times6)$

Final: $C r_{2} O_{7}^{2 -} + 14 H^{+} + 6 F e^{2 +} \to 2 C r^{3 +} + 7 H_{2} O + 6 F e^{3 +}$

Basic Solution Conversion

For basic solutions, follow Method 1: Half-Reaction Method then:

Add $O H^{-}$ to both sides (equal to $H^{+}$ count)
Combine $H^{+} + O H^{-} \to H_{2} O$
Cancel excess water

Method 2 Examples: Oxidation Number Method

Detailed Example

Balance: $C u + H N O_{3} \to C u (N O_{3})_{2} + NO + H_{2} O$

Step 1: Assign oxidation numbers using Oxidation Numbers: Complete Guide

$C u^{0} + H^{+ 1} N^{+ 5} O_{3}^{- 2} \to C u^{+ 2} (N^{+ 5} O_{3}^{- 2})_{2} + N^{+ 2} O^{- 2} + H_{2}^{+ 1} O^{- 2}$

Step 2: Identify electron transfer

Cu: $0 \to + 2$ (loses $2 e^{-}$ , oxidized)
N: $+ 5 \to + 2$ (gains $3 e^{-}$ , reduced)

Step 3: Balance electron transfer

LCM of 2 and 3 = 6
Need 3 Cu ( $3 \times 2 e^{-} = 6 e^{-}$ lost)
Need 2 N reduced ( $2 \times 3 e^{-} = 6 e^{-}$ gained)

Step 4: Complete balancing

$3 C u + 8 H N O_{3} \to 3 C u (N O_{3})_{2} + 2 NO + 4 H_{2} O$

Strategy Selection

Reaction Type	Best Method	Reference
Simple reactions	Inspection	Quick visual balance
Complex ionic	Half-reaction
Need electron tracking	Oxidation number
Acidic/basic solutions	Half-reaction	Built-in pH handling
Disproportionation	Half-reaction	Multiple pathways

Final Checks

Mass balance: Equal atoms on both sides
Charge balance: Equal total charge
Electron balance: $e^{-}$ lost = $e^{-}$ gained (use Oxidation Numbers: Complete Guide to verify)

TODO:

Learn redox reactions with both methods, oxyidation numbers and yapyapyap
- Ion electron method
- Oxidation number Change Method

Jason Cameron

Explorer

Redox Balancing