A Galvanic Cell Consists of Two Beakers Which of the Following Is the Oxidation Half-Reaction That Occurs

A galvanic cell consists of two beakers. One beaker contains a strip of lead immersed in aqueous sulfuric acid, and the other contains a platinum electrode immersed in aqueous nitric acid. The two solutions are connected by a salt bridge, and the electrodes are connected by a wire. Current begins to flow, and bubbles of a gas appear at the platinum electrode. The spontaneous redox reaction that occurs is described by the following balanced chemical equation: $$3 \mathrm {~Pb} ( \mathrm {~s} ) + 2 \mathrm { NO } _ { 3 } ^ { - } ( \mathrm { aq } ) + 8 \mathrm { H } ^ { + } ( \mathrm { aq } ) \rightarrow 3 \mathrm {~Pb} ^ { 2 + } ( \mathrm { aq } ) + 2 \mathrm { NO } ( \mathrm { g } ) + 4 \mathrm { H } _ { 2 } \mathrm { O } ( \mathrm { l } )$$ Which of the following is the oxidation half-reaction that occurs in the anode?

A) $$\mathrm { Pb } ( \mathrm { s } ) \rightarrow \mathrm { Pb } ^ { 2 + } ( \mathrm { aq } ) + 2 \mathrm { e } ^ { - }$$
B) $\mathrm { NO } _ { 3 } ^ { - } ( \mathrm { aq } ) + 3 \mathrm { e } ^ { - } \rightarrow \mathrm { NO } ( \mathrm { g } )$
C) $\mathrm { NO } _ { 3 } ^ { - } ( \mathrm { aq } ) + 4 \mathrm { H } ^ { + } ( \mathrm { aq } ) + 3 \mathrm { e } ^ { - } \rightarrow \mathrm { NO } ( \mathrm { g } ) + 2 \mathrm { H } _ { 2 } \mathrm { O } ( \mathrm { l } )$
D) $\mathrm { Pb } ^ { 2 + } ( \mathrm { aq } ) + \mathrm { H } _ { 2 } \mathrm { O } ( \mathrm { l } ) \rightarrow \mathrm { PbO } ( \mathrm { s } ) + 2 \mathrm { H } ^ { + } ( \mathrm { aq } ) + 2 \mathrm { e } ^ { - }$
E) $3 \mathrm {~Pb} ( \mathrm {~s} ) + 2 \mathrm { NO } _ { 3 } ^ { - } ( \mathrm { aq } ) \rightarrow 3 \mathrm {~Pb} ^ { 2 + } ( \mathrm { aq } ) + 2 \mathrm { NO } ( \mathrm { g } ) + 3 \mathrm { e } ^ { - }$