What reaction can occur, if any, when the following experiments are carried out under standard-state conditions? (b) An acidic solution of FeSO4 is exposed to oxygen.

Verified step by step guidance

1. Identify the species involved in the reaction: FeSO4, H+ (from the acidic solution), and O2.
2. Consider the possible oxidation and reduction reactions: Fe2+ can be oxidized to Fe3+, and O2 can be reduced to H2O in an acidic solution.
3. Write the half-reactions for the oxidation and reduction processes: \( \text{Fe}^{2+} \rightarrow \text{Fe}^{3+} + e^- \) and \( \text{O}_2 + 4\text{H}^+ + 4e^- \rightarrow 2\text{H}_2\text{O} \).
4. Balance the electrons in the half-reactions to ensure the number of electrons lost in oxidation equals the number gained in reduction.
5. Combine the balanced half-reactions to write the overall balanced redox reaction.

Key Concepts

Oxidation-Reduction Reactions

Oxidation-reduction (redox) reactions involve the transfer of electrons between species, leading to changes in oxidation states. In the context of the question, FeSO4 can undergo oxidation when exposed to oxygen, where iron (Fe) in the +2 oxidation state can be oxidized to iron (Fe) in the +3 state.

Standard-State Conditions

Standard-state conditions refer to a set of specific conditions (1 bar pressure, 25°C temperature) under which the properties of substances are measured. These conditions are crucial for predicting the behavior of chemical reactions, as they provide a consistent reference point for comparing thermodynamic data and reaction feasibility.

Acidic Solutions and Their Effects

Acidic solutions, characterized by a high concentration of hydrogen ions (H+), can influence the reactivity of metal ions. In the case of FeSO4 in an acidic environment, the presence of H+ can stabilize the +2 oxidation state of iron, but upon exposure to oxygen, the iron can still be oxidized to the +3 state, leading to the formation of Fe2(SO4)3.