Oxidation-Reduction Reactions and Stoichiometry: Study Notes

Oxidation-Reduction Reactions

Introduction to Oxidation-Reduction (Redox) Reactions

Oxidation-reduction (redox) reactions are chemical processes in which electrons are transferred between substances. These reactions are fundamental to many chemical and biological processes, including combustion, respiration, and corrosion.

• Oxidation: The loss of electrons by a substance.

• Reduction: The gain of electrons by a substance.

• Mnemonic: OIL RIG (Oxidation Is Loss, Reduction Is Gain of electrons).

Example: In the reaction of hydrogen with oxygen to form water:

• Hydrogen is oxidized (loses electrons).

• Oxygen is reduced (gains electrons).

Examples of Oxidation and Reduction Reactions

• Oxidation Examples:

  • Combustion of magnesium: $2Mg(s) + O_2(g) \rightarrow 2MgO(s)$

  • Formation of hydrogen ions: $H_2(g) \rightarrow 2H^+(aq) + 2e^-$

  • Iron(II) to iron(III): $Fe^{2+}(aq) \rightarrow Fe^{3+}(aq) + e^-$

• Reduction Examples:

  • Formation of hydrogen gas: $2H^+(aq) + 2e^- \rightarrow H_2(g)$

  • Reduction of iron(III): $Fe^{3+}(aq) + e^- \rightarrow Fe^{2+}(aq)$

Identifying Oxidation and Reduction

• Oxidation involves an increase in oxidation number.

• Reduction involves a decrease in oxidation number.

• Oxidizing agent: Substance that causes oxidation (is reduced).

• Reducing agent: Substance that causes reduction (is oxidized).

Oxidation Numbers

Definition and Purpose

Oxidation numbers (or states) are assigned to atoms in molecules or ions to keep track of electron transfer in redox reactions. They help identify which atoms are oxidized and which are reduced.

Rules for Assigning Oxidation Numbers

• The oxidation number of an atom in its elemental form is 0 (e.g., $O_2$, $H_2$, $Na$).

• The oxidation number of a monatomic ion equals its charge (e.g., $Na^+$ is +1, $Cl^-$ is -1).

• In compounds:

  • Group 1 metals: +1

  • Group 2 metals: +2

  • Fluorine: -1

  • Hydrogen: +1 (except in metal hydrides, where it is -1)

  • Oxygen: -2 (except in peroxides, where it is -1)

  • Chlorine, bromine, iodine: usually -1, except when combined with oxygen or fluorine

• The sum of oxidation numbers in a neutral compound is 0; in a polyatomic ion, it equals the ion's charge.

Examples

• Oxidation number of S in $SO_2$:

  • Let S = x; O = -2

  • $x + 2(-2) = 0 \implies x = +4$

• Oxidation number of N in $NO_3^-$:

  • Let N = x; O = -2

  • $x + 3(-2) = -1 \implies x = +5$

Balancing Oxidation-Reduction Equations: The Half-Reaction Method

Overview

The half-reaction method is a systematic approach to balancing redox equations, especially in aqueous solutions. It involves separating the oxidation and reduction processes and balancing each for mass and charge.

Steps in the Half-Reaction Method

1. Assign oxidation numbers to all atoms to identify what is oxidized and what is reduced.

2. Write separate half-reactions for oxidation and reduction.

3. Balance all elements except H and O.

4. Balance O by adding $H_2O$; balance H by adding $H^+$ (in acidic solution) or $OH^-$ (in basic solution).

5. Balance charge by adding electrons ($e^-$).

6. Multiply half-reactions by appropriate factors so electrons lost = electrons gained.

7. Add the half-reactions and cancel identical species.

Example: Balancing a Redox Equation

• Balance the following in acidic solution: $MnO_4^-(aq) + Fe^{2+}(aq) \rightarrow Mn^{2+}(aq) + Fe^{3+}(aq)$

• Step 1: Write half-reactions:

  • Oxidation: $Fe^{2+} \rightarrow Fe^{3+} + e^-$

  • Reduction: $MnO_4^- + 8H^+ + 5e^- \rightarrow Mn^{2+} + 4H_2O$

• Step 2: Multiply oxidation half-reaction by 5 to balance electrons.

• Step 3: Add and simplify:

$5Fe^{2+} + MnO_4^- + 8H^+ \rightarrow 5Fe^{3+} + Mn^{2+} + 4H_2O$

Practice Problems

• Classify each as oxidation or reduction:

  • $Zn \rightarrow Zn^{2+} + 2e^-$ (Oxidation)

  • $Cl_2 + 2e^- \rightarrow 2Cl^-$ (Reduction)

• Assign oxidation numbers to all atoms in $K_2Cr_2O_7$, $SO_4^{2-}$, $NO_3^-$, $NH_4^+$, $H_2O_2$.

• Balance the following equations using the half-reaction method:

  • $Fe^{2+} + Cr_2O_7^{2-} \rightarrow Fe^{3+} + Cr^{3+}$

  • $MnO_4^- + C_2O_4^{2-} \rightarrow Mn^{2+} + CO_2$

Summary Table: Common Oxidation Numbers

Element	Common Oxidation Number(s)	Notes
Group 1 metals (e.g., Na, K)	+1	In all compounds
Group 2 metals (e.g., Mg, Ca)	+2	In all compounds
Hydrogen	+1, -1	+1 except in metal hydrides (-1)
Oxygen	-2, -1	-2 except in peroxides (-1)
Fluorine	-1	In all compounds
Chlorine, Bromine, Iodine	-1	Except with O or F

Additional info: These notes are based on standard college-level General Chemistry content, specifically focusing on redox reactions, oxidation numbers, and balancing redox equations, which are essential for understanding chemical reactions and stoichiometry.