At one time on Earth, iron was present mostly as iron(II). Later, once plants had produced a significant quantity of oxygen in the atmosphere, the iron became oxidized to iron(III). Show that Fe2+(aq) can be spontaneously oxidized to Fe3+(aq) by O2(g) at 25°C assuming the following reasonable environmental conditions:

Key Concepts

Oxidation and Reduction

Oxidation refers to the loss of electrons from a substance, while reduction is the gain of electrons. In the context of the question, Fe2+ is oxidized to Fe3+ as it loses an electron, and O2 acts as the oxidizing agent, facilitating this process. Understanding these concepts is crucial for analyzing redox reactions.

Standard Electrode Potentials

Standard electrode potentials (E°) provide a measure of the tendency of a chemical species to be reduced. The difference in E° values between the half-reactions of Fe2+/Fe3+ and O2/H2O can be used to determine the spontaneity of the reaction. A positive cell potential indicates that the reaction can occur spontaneously under standard conditions.

Gibbs Free Energy

Gibbs free energy (ΔG) is a thermodynamic potential that helps predict the spontaneity of a reaction. A negative ΔG indicates that a reaction can occur spontaneously. The relationship between ΔG, the standard electrode potential, and temperature is essential for determining whether the oxidation of Fe2+ by O2 is thermodynamically favorable.