Which of the following represents the oxidation half-reaction that occurs at the anode in a hydrogen fuel cell?

• A. 2 H_2O(l) → O_2(g) + 4 H^+(aq) + 4 e^-
• B. 2 H_2(g) → 4 H^+(aq) + 4 e^-
• C. 4 H^+(aq) + 4 e^- → 2 H_2(g)
• D. O_2(g) + 4 H^+(aq) + 4 e^- → 2 H_2O(l)

Answer: (B)

Verified step by step guidance

1. Understand that in a hydrogen fuel cell, the oxidation half-reaction occurs at the anode, where hydrogen gas is converted into protons and electrons.
2. Recall that oxidation involves the loss of electrons, so the species that loses electrons is the one being oxidized at the anode.
3. Identify the half-reactions given and determine which one shows hydrogen gas (H_2) losing electrons to form protons (H^+). This corresponds to the oxidation process.
4. Write the oxidation half-reaction at the anode as: \$2\ H_2(g) \rightarrow 4\ H^+(aq) + 4\ e^-$, showing hydrogen gas losing electrons.
5. Confirm that the other half-reactions represent either reduction (gain of electrons) or the overall reaction, and thus do not occur at the anode.