R3.2.1 Oxidation and reduction definitions

What Do Rusting Iron, Burning Wood, and Photosynthesis Have in Common?

Oxidation and Reduction: More Than Just Electron Transfer

Redox reactions involve two simultaneous processes: oxidation and reduction. These terms can be understood in different ways depending on the context.

Oxidation

Electron Loss: When an atom or ion loses electrons, it undergoes oxidation.

Example

$Na(s) \to Na⁺(aq) + e ⁻$ Sodium loses an electron and is oxidized.

Oxygen Gain: Oxidation can also refer to the addition of oxygen.

Example

$2 Mg(s) + O₂(g) \to 2 MgO(s)$ Magnesium gains oxygen and is oxidized.

Hydrogen Loss: Oxidation may involve the loss of hydrogen.

Example

$CH₄(g) + 2 O₂(g) \to CO₂(g) + 2 H₂O(g)$ Methane loses hydrogen during combustion.

Note

While oxygen gain and hydrogen loss are traditional definitions of oxidation, they are less commonly used today compared to the electron transfer definition. Always consider the context of the reaction.

Reduction

Electron Gain: Reduction occurs when an atom or ion gains electrons.

Example

$Cl₂(g) + 2 e ⁻ \to 2 Cl⁻(aq)$ Chlorine gains electrons and is reduced.

Oxygen Loss: Reduction can also mean the removal of oxygen.

Example

$CuO(s) + H₂(g) \to Cu(s) + H₂O(g)$ Copper(II) oxide loses oxygen and is reduced.

Hydrogen Gain: Reduction may involve the addition of hydrogen.

Example

$C₂H₄(g) + H₂(g) \to C₂H₆(g)$ Ethene gains hydrogen in this reaction.

Tip

Use the mnemonic OIL RIG: Oxidation Is Loss (of electrons), Reduction Is Gain (of electrons).

Oxidation States: A Tool for Tracking Electron Transfers

To systematically analyze redox reactions, chemists assign oxidation states to atoms. These numbers help track the movement of electrons in a reaction.

Rules for Assigning Oxidation States

Elements in Their Standard State: The oxidation state of an atom in its elemental form is always 0.
Ions: The oxidation state equals the ion’s charge.
Fluorine: Fluorine always has an oxidation state of -1 in compounds.
Oxygen: Oxygen typically has an oxidation state of -2, except in peroxides ( $- 1$ ) or when bonded to fluorine ( $+ 2$ ).
Hydrogen: Hydrogen usually has an oxidation state of +1, except in hydrides ( $- 1$ ).
Neutral Compounds: The sum of oxidation states in a neutral compound must equal 0.
Polyatomic Ions: The sum of oxidation states in a polyatomic ion must equal the ion’s charge.

Example

Consider the compound $H₂SO₄$ (sulfuric acid):

Hydrogen: Each H is +1 (total = +2).
Oxygen: Each O is -2 (total = -8).
Sulfur: To balance the total charge to 0, sulfur must be +6.

Self review

Assign oxidation states to all atoms in $K₂Cr₂O₇$ . (Answer: K = +1, Cr = +6, O = -2)

Recognizing Oxidizing and Reducing Agents

In redox reactions, the species that is oxidized acts as the reducing agent, and the species that is reduced acts as the oxidizing agent.

Recognizing reducing and oxidizing agents.

Example

Reaction Between Zinc and Copper(II) Sulfate

$Zn(s) + CuSO₄(aq) \to ZnSO₄(aq) + Cu(s)$

Zinc is oxidized (reducing agent).
Copper(II) ion is reduced (oxidizing agent).

Common Mistake

Do not confuse the oxidizing and reducing agents. Remember: the oxidizing agent is reduced, and the reducing agent is oxidized.

Transition Metals and Variable Oxidation States

Transition metals, such as iron and manganese, often exhibit multiple oxidation states, enabling them to participate in diverse redox reactions.

Example 1: Manganese in Permanganate Ion

${MnO}_{4}^{-} (a q) + 8 H^{+} (a q) + 5 e^{-} \to {Mn}^{2 +} (a q) + 4 H_{2} O (l)$

Manganese is reduced from +7 in $M n O_{4}^{-}$ to +2 in $M n^{2 +}$ .

Example 2: Iron in Iron(III) Reduction

${Fe}^{3 +} (a q) + e^{-} \to {Fe}^{2 +} (a q)$

Iron is reduced from +3 in $F e^{3 +}$ to +2 in $F e^{2 +}$ .

Example 3: Chromium in Dichromate Ion

${Cr}_{2} O_{7}^{2 -} (a q) + 14 H^{+} (a q) + 6 e^{-} \to 2 {Cr}^{3 +} (a q) + 7 H_{2} O (l)$

Chromium is reduced from +6 in $C r_{2} O_{7}^{2 -}$ to +3 in $C r^{3 +}$ .

Example 4: Copper in Copper(II) Reduction

${Cu}^{2 +} (a q) + 2 e^{-} \to Cu (s)$

Copper is reduced from +2 in $C u^{2 +}$ to 0 in solid copper metal.

Note

The variable oxidation states of transition metals are crucial in catalysis, batteries, and biological processes.

Reflection Questions

Self review

In the reaction $Fe(s) + HCl(aq) \to FeCl₂(aq) + H₂(g)$ , identify the oxidizing and reducing agents.
Why are transition metals particularly versatile in redox reactions?
How do oxidation states simplify the analysis of complex reactions?

Theory of Knowledge

How does the concept of oxidation evolve as science advances? Why might different definitions (electron transfer, oxygen gain/loss) coexist?

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R3.2.1 Oxidation and reduction definitions

What Do Rusting Iron, Burning Wood, and Photosynthesis Have in Common?

Oxidation and Reduction: More Than Just Electron Transfer

Oxidation

Reduction

Oxidation States: A Tool for Tracking Electron Transfers

Rules for Assigning Oxidation States

Recognizing Oxidizing and Reducing Agents

Reaction Between Zinc and Copper(II) Sulfate

Transition Metals and Variable Oxidation States

Example 1: Manganese in Permanganate Ion

Example 2: Iron in Iron(III) Reduction

Example 3: Chromium in Dichromate Ion

Example 4: Copper in Copper(II) Reduction

Reflection Questions

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Introduction to Redox Reactions