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:

Verified step by step guidance

1. Identify the half-reactions involved in the oxidation of Fe^{2+} to Fe^{3+} and the reduction of O_2 to H_2O.
2. Write the oxidation half-reaction: Fe^{2+} \rightarrow Fe^{3+} + e^{-}.
3. Write the reduction half-reaction: O_2 + 4H^+ + 4e^{-} \rightarrow 2H_2O.
4. Determine the standard reduction potentials (E^\circ) for both half-reactions from a standard reduction potential table.
5. Calculate the standard cell potential (E^\circ_{cell}) using the formula: E^\circ_{cell} = E^\circ_{cathode} - E^\circ_{anode}. If E^\circ_{cell} is positive, the reaction is spontaneous.

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.