Redox Reactions: Concepts, Oxidation Numbers, and Applications

9.- Reacciones de Oxidación-Reducción (Redox Reactions)

9.1 Concepto de oxidación - reducción

Redox reactions are fundamental chemical processes involving the transfer of electrons between substances. These reactions are essential in biological systems, industrial processes, and everyday phenomena such as corrosion and energy production.

• Oxidation: The process in which a substance loses electrons.

• Reduction: The process in which a substance gains electrons.

• Redox Reaction: A chemical reaction where oxidation and reduction occur simultaneously.

• Example: The reaction between zinc and copper(II) ions: $\text{Zn} + \text{Cu}^{2+} \rightarrow \text{Zn}^{2+} + \text{Cu}$ Here, Zn is oxidized (loses electrons), and Cu2+ is reduced (gains electrons).

Redox reactions are crucial for processes such as respiration, combustion, corrosion, and electrochemical cells.

9.2 Índice o número de oxidación (Oxidation Number)

The oxidation number (or oxidation state) is a practical tool to determine whether a substance is oxidized or reduced in a reaction. It represents the hypothetical charge an atom would have if all bonds were ionic.

• Definition: The oxidation number of an element is the number of electrons gained or lost compared to the neutral atom.

• Rules for Assigning Oxidation Numbers:

  1. For monoatomic ions, the oxidation number equals the ion charge (e.g., Na+ = +1, Cl- = -1).

  2. In compounds, group 1 metals are always +1, group 2 metals are +2.

  3. Hydrogen is +1 (except in metal hydrides, where it is -1).

  4. Oxygen is usually -2 (except in peroxides, where it is -1, and in OF2, where it is +2).

  5. Fluorine is always -1.

  6. The sum of oxidation numbers in a neutral compound is zero; in a polyatomic ion, it equals the ion charge.

• Example: In H2O, H = +1, O = -2; sum = 0.

• Example: In SO42-, O = -2 (total -8), so S = +6 to balance the charge.

Oxidation numbers help identify which atoms are oxidized and which are reduced in a reaction.

9.3 Concepto de semirreacción (Half-Reaction Concept)

Redox reactions can be split into two half-reactions: one for oxidation and one for reduction. This approach clarifies the electron transfer process.

• Oxidation Half-Reaction: Shows the loss of electrons.

• Reduction Half-Reaction: Shows the gain of electrons.

• Example:

  • Oxidation: $\text{Zn} \rightarrow \text{Zn}^{2+} + 2e^-$

  • Reduction: $\text{Cu}^{2+} + 2e^- \rightarrow \text{Cu}$

Balancing redox reactions often involves writing and balancing these half-reactions separately.

9.4 Ajuste de reacciones Redox (Balancing Redox Reactions)

Balancing redox reactions requires ensuring both mass and charge are conserved. The process differs depending on whether the reaction occurs in acidic or basic medium.

• Steps for Balancing:

  1. Write the oxidation and reduction half-reactions.

  2. Balance all elements except H and O.

  3. Balance O by adding H2O; balance H by adding H+ (acidic) or OH- (basic).

  4. Balance charge by adding electrons.

  5. Combine the half-reactions, ensuring electrons cancel.

• Example (Acidic Medium):

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

  • Reduction: $\text{MnO}_4^- + 8H^+ + 5e^- \rightarrow \text{Mn}^{2+} + 4H_2O$

In basic medium, add OH- to neutralize H+ and form water.

9.5 Electrólisis. Ley de Faraday (Electrolysis and Faraday's Law)

Electrolysis is a process where electrical energy drives a non-spontaneous redox reaction. Faraday's Law quantifies the relationship between the amount of substance altered at an electrode and the quantity of electricity passed through the cell.

• Faraday's Law: The amount of substance produced at an electrode is proportional to the total electric charge passed.

• Equation: $m = \frac{Q}{F \cdot n} \cdot M$ where $m$ = mass of substance, $Q$ = total charge, $F$ = Faraday constant ($96,485$ C/mol), $n$ = number of electrons, $M$ = molar mass.

• Application: Used in electroplating, purification of metals, and production of chemicals.

9.6 Potenciales de electrodo. Pares oxidación - reducción (Electrode Potentials and Redox Pairs)

Electrode potential measures the tendency of a chemical species to gain or lose electrons. Standard electrode potentials are referenced to the standard hydrogen electrode (SHE).

• Redox Pair: A conjugate pair consisting of an oxidized and reduced form (e.g., Fe3+/Fe2+).

• Standard Electrode Potential ($E^\circ$): The voltage measured under standard conditions (1 M, 1 atm, 25°C).

• Table of Standard Electrode Potentials:

Redox Pair	$E^\circ$ (V)
Cu2+/Cu	+0.34
Zn2+/Zn	-0.76
Fe3+/Fe2+	+0.77
Ag+/Ag	+0.80
Additional info: Values inferred from standard tables.

The cell potential is calculated by subtracting the anode potential from the cathode potential.

9.7 La ecuación de Nernst (Nernst Equation)

The Nernst equation relates the cell potential to the concentrations of reactants and products, allowing calculation of non-standard cell potentials.

• Equation: $E = E^\circ - \frac{0.0592}{n} \log Q$ where $E$ = cell potential, $E^\circ$ = standard cell potential, $n$ = number of electrons transferred, $Q$ = reaction quotient.

• Application: Used to predict cell voltage under non-standard conditions.

9.8 Valoraciones Redox (Redox Titrations)

Redox titrations are analytical techniques used to determine the concentration of an analyte by reacting it with a titrant in a redox reaction.

• Key Points:

  • Indicator is often used to detect the endpoint.

  • Common titrants include KMnO4, K2Cr2O7.

  • Stoichiometry is based on the balanced redox equation.

• Example: Determining iron(II) concentration using permanganate titration.

9.9 Problemas y cuestiones (Problems and Questions)

This section typically includes practice problems and conceptual questions to reinforce understanding of redox concepts, calculations of oxidation numbers, balancing redox reactions, and applications of the Nernst equation.

• Example Problem: Calculate the cell potential for a Zn/Cu cell with given ion concentrations using the Nernst equation.

• Example Question: Assign oxidation numbers to all atoms in K2Cr2O7.

Additional info: These notes are based on textbook-style content and include inferred standard values and examples for completeness.