Electrochemistry: Redox Reactions, Oxidation Numbers, and Balancing

Electrochemistry

Introduction

Electrochemistry is the study of chemical processes that cause electrons to move, which is the basis for redox (oxidation-reduction) reactions. These reactions are fundamental to understanding how chemical energy is converted to electrical energy and vice versa.

Redox Reactions

General Concepts

• Redox (oxidation-reduction) reaction: A chemical reaction in which electrons are transferred from one reactant to another.

• Oxidation: Loss of electrons by a substance. Mnemonic: OIL (Oxidation Is Loss)

• Reduction: Gain of electrons by a substance. Mnemonic: RIG (Reduction Is Gain)

• Reducing agent: The species that donates electrons and is itself oxidized.

• Oxidizing agent: The species that accepts electrons and is itself reduced.

Example:

• $\mathrm{Zn}(s) + \mathrm{Cu}^{2+}(aq) \rightarrow \mathrm{Zn}^{2+}(aq) + \mathrm{Cu}(s)$

• Zn loses 2 electrons (is oxidized); Zn is the reducing agent.

• Cu2+ gains 2 electrons (is reduced); Cu2+ is the oxidizing agent.

Identifying Agents in Redox Reactions

• In $\mathrm{Fe}(s) + \mathrm{Cu}^{2+}(aq) \rightarrow \mathrm{Fe}^{2+}(aq) + \mathrm{Cu}(s)$, Fe is the reducing agent.

• In $\mathrm{Cu}(s) + 2\mathrm{Ag}^{+}(aq) \rightarrow \mathrm{Cu}^{2+}(aq) + 2\mathrm{Ag}(s)$, Ag+ is the oxidizing agent.

Half-Reactions

Every redox reaction can be split into two half-reactions: one for oxidation and one for reduction.

• Oxidation half-reaction: Shows the loss of electrons.

• Reduction half-reaction: Shows the gain of electrons.

Example:

• Overall: $\mathrm{Zn}(s) + \mathrm{Cu}^{2+}(aq) \rightarrow \mathrm{Zn}^{2+}(aq) + \mathrm{Cu}(s)$

• Oxidation: $\mathrm{Zn}(s) \rightarrow \mathrm{Zn}^{2+}(aq) + 2e^-$

• Reduction: $\mathrm{Cu}^{2+}(aq) + 2e^- \rightarrow \mathrm{Cu}(s)$

Oxidation Numbers

Definition and Assignment Rules

The oxidation number (or oxidation state) of an atom in a compound is the charge it would have if the compound were ionic. Assigning oxidation numbers helps identify which atoms are oxidized and which are reduced.

• Elemental form: Oxidation number is 0.

• Monatomic ion: Oxidation number equals the ion's charge.

• Common nonmetals:

  • Fluorine: -1

  • Hydrogen: +1 (with nonmetals), -1 (with metals)

  • Oxygen: -2 (most compounds), -1 (peroxides)

  • Group I metals: +1; Group II metals: +2

  • Cl, Br, I: -1 (except when combined with oxygen or other halogens)

• Sum of oxidation numbers equals the charge on the molecule or ion.

Example: Assign oxidation numbers in the following compounds:

• $\mathrm{N}_2\mathrm{O}_5$

• $\mathrm{HSO}_4^-$

• $\mathrm{NaIO}_3$

• $\mathrm{NH}_4^+$

Balancing Redox Reactions

Stepwise Procedure (Acidic and Basic Conditions)

Balancing redox reactions ensures both mass and charge are conserved. The process involves several steps:

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

2. Separate the equation into two half-reactions (oxidation and reduction).

3. Balance all atoms except O and H in each half-reaction.

4. Add $\mathrm{H}_2\mathrm{O}$ to balance oxygen atoms.

5. Add $\mathrm{H}^+$ to balance hydrogen atoms (for acidic solution).

6. Add electrons ($e^-$) to balance charges in each half-reaction.

7. Multiply half-reactions by appropriate factors so the number of electrons lost equals the number gained.

8. Add the half-reactions together, canceling like terms.

9. For basic solutions, add $\mathrm{OH}^-$ ions to both sides to neutralize $\mathrm{H}^+$ and form $\mathrm{H}_2\mathrm{O}$.

10. Verify that all atoms and charges are balanced.

Example:

• Balance: $\mathrm{Fe}^{2+} + \mathrm{Cr}_2\mathrm{O}_7^{2-} \rightarrow \mathrm{Fe}^{3+} + \mathrm{Cr}^{3+}$ (acidic conditions)

• Balance: $\mathrm{Cu}(s) + \mathrm{HNO}_3(aq) \rightarrow \mathrm{Cu}^{2+}(aq) + \mathrm{NO}(g)$ (acidic and basic conditions)

Additional info:

• Balancing in basic solution requires converting $\mathrm{H}^+$ to $\mathrm{H}_2\mathrm{O}$ using $\mathrm{OH}^-$.

• Always check that both mass and charge are balanced in the final equation.