Understanding Oxidation-Reduction (Redox) Reactions

Posted on Feb 16, 2025 in Chemistry

Oxidation: Gain of O2, loss of H, loss of electrons. Reduction: The reverse of oxidation. Oxidant: Element that oxidizes a compound. Reductant: Element that reduces a compound. Oxidation Number (Ox. No.): The charge an atom would have in a molecule or ion if electrons were completely transferred, as if they were free ions. Elements in their natural state have an oxidation number of zero. In monatomic ions, the oxidation number equals the charge of the ion. The oxidation number of O is usually -2 (except in H2O2 and O2-2, where it is -1). The oxidation number of H is +1, except in metal hydrides where it is -1. Group IA metals are +1, Group IIA metals are +2, and fluorine is always -1. The sum of the oxidation numbers of all atoms in a molecule equals the charge of the molecule or ion.

Balancing Redox Reactions

Acidic Medium:

  1. Write the unbalanced reaction and assign oxidation numbers.
  2. Separate the reaction into two half-reactions: oxidation and reduction.
  3. Balance all elements except O and H in each half-reaction.
  4. In acidic medium, add H2O to balance oxygen and H+ to balance hydrogen.
  5. Add electrons to balance the charge in each half-reaction.
  6. Multiply each half-reaction by a factor so that the number of electrons gained equals the number of electrons lost.
  7. Add the balanced half-reactions together, canceling out electrons.

Basic Medium:

  1. Balance the reaction as if it were in acidic medium.
  2. Add OH– ions to both sides of the equation to neutralize the H+ ions, forming water.
  3. Cancel out any water molecules that appear on both sides of the equation.

Electrochemical Cells (Pilas)

  • Anode: Oxidation occurs.
  • Cathode: Reduction occurs.
  • Electrochemical Cell: Generates electricity from spontaneous redox reactions.
  • Electrode: Surface where oxidation or reduction occurs. The potential difference between the anode and cathode is called the cell voltage or potential.
  • Galvanic Cells: Produce electricity from spontaneous chemical reactions.
  • Electrolytic Cells: Use electricity to drive non-spontaneous chemical transformations.
  • Couple, M | Mn+: A pair of species related by a change in the number of electrons.

Standard Electrode Potentials

The voltage of each cell can be determined precisely. It is difficult to establish the potential of an individual electrode, so a standard reduction potential is chosen arbitrarily. The standard hydrogen electrode (SHE) is assigned a potential of zero. Standard conditions are 1M solutions and 1 atm for gases.

E0cell = E0cathode – E0anode

Relationship Between Ecell, ΔG, and Keq

The cell performs electrical work: Welect = -n * F * Ecell (where F = 96485 C/mol).

ΔG0 = -n * F * E0cell = -n * R * T * lnK

Nernst Equation

E = E0 – (RT / (nF)) * lnQ

Spontaneity

  • ΔG < 0, E0cell > 0, K > 1: The reaction is spontaneous in the forward direction under the specified conditions.
  • E0cell = 0: The reaction is at equilibrium.
  • ΔG > 0, E0cell < 0: The reaction is non-spontaneous in the forward direction (spontaneous in reverse).

Concentration Cells

Batteries comprising an anode and a cathode of the same material but with different ion concentrations.

Types of Batteries

Primary (dry cell Leclanche), secondary (lead-acid battery, mercury battery, lithium batteries, fuel cells).

Electrolysis

The process in which electrical energy is used to drive non-spontaneous reactions.

Faraday’s Laws

  1. The amount of substance oxidized at the anode or reduced at the cathode is proportional to the quantity of electricity passed through the cell.
  2. During electrolysis, one Faraday of electricity reduces and oxidizes one equivalent of reducing agent and oxidizing agent.

Redox Titrations

The equivalence point is indicated by a color change (equivalent oxidant = equivalent reductant).

Corrosion

An irreversible interfacial reaction of a material with its environment, implying consumption of the material or the introduction of constituents from the environment. The corrosion cell is formed in metals due to anodic areas (where oxides are produced; the metal has a greater tendency to oxidize due to the presence of impurities) and cathodic areas (corrosion-resistant). In pure metals, the anode and cathode areas are alternately distributed homogeneously. When the potential reaches a value of 0, corrosion stops, or continues until all the metal is consumed.

Types of Corrosion

  • Uniform: Not particularly dangerous.
  • Localized: Anode and cathode areas are concentrated in specific parts of the metal. Causes include structural defects or exposure to aggressive agents.
  • Pitting: Due to the presence of aggressive anions.
  • Galvanic: Occurs when two dissimilar metals are in contact in the presence of an electrolyte.