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Question

Answers

a.) 0.59 V

b.) 0.00 V

c.) -0.59 V

d.) -0.059 V

Answer
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The electrode potential can be calculated as -

${E_{cell}} = E_{cell}^0 - \dfrac{{0.059}}{1}\log \dfrac{{product}}{{reactant}}$

Where ${E_{cell}}$ is the potential difference between two half cells.

In this case, $E_{cell}^0$ is taken as zero.

We know that the word potential means the ability. The potential can be oxidation potential of hydrogen electrode or the reduction potential of hydrogen electrode.

The reduction potential can be defined as the ability of the element to accept electrons and get reduced while the oxidation potential is the ability of an element to get oxidized by donating the electrons.

The reduction potential can be calculated as -

We know the formula

${E_{cell}} = E_{cell}^0 - \dfrac{{0.059}}{1}\log \dfrac{{product}}{{reac\tan t}}$

We have, $E_{cell}^0$ = 0

Thus, the formula is -

\[{E_{cell}} = - 0.059\log \dfrac{{product}}{{reac\tan t}}\]

The reaction for hydrogen electrode can be written as -

${H^ + } + {e^ - } \to \dfrac{1}{2}{H_2}$

Further, we have been given pH = 10

Thus, $\left[ {{H^ + }} \right]$ = ${10^{ - 10}}M$

Now, putting values in the formula; we have

\[{E_{cell}} = - 0.059\log \dfrac{{\left[ {{H_2}} \right]}}{{\left[ {{H^ + }} \right]}}\]

\[{E_{cell}} = - 0.059\log \dfrac{1}{{{{10}^{ - 10}}}}\]

\[{E_{cell}} = - 0.059 \times 10\]

\[{E_{cell}}\] = -0.59 V

Thus, this is the reduction potential of hydrogen electrodes.

So, option c.) is the correct answer.

Now, let us calculate the oxidation potential of hydrogen electrodes.

The reaction for oxidation potential may be written as -

$\frac{1}{2}{H_2} \to {H^ + } + {e^ - }$

So, now again using the formula; we can write

${E_{cell}} = E_{cell}^0 - \dfrac{{0.059}}{1}\log \dfrac{{product}}{{reac\tan t}}$

We have, $E_{cell}^0$ = 0

Thus, the formula is -

\[{E_{cell}} = - 0.059\log \dfrac{{product}}{{reac\tan t}}\]

We have pH = 10

Thus, $\left[ {{H^ + }} \right]$ = ${10^{ - 10}}M$

Now, putting values in the formula

\[{E_{cell}} = - 0.059\log \dfrac{{\left[ {{H^ + }} \right]}}{{\left[ {{H_2}} \right]}}\]

\[{E_{cell}} = - 0.059\log \dfrac{{{{10}^{ - 10}}}}{1}\]

\[{E_{cell}} = - 0.059 \times ( - 10)\]

\[{E_{cell}}\] = 0.59 V

Thus, this is the oxidation potential of hydrogen electrodes.

So, option a.) is the correct answer.

Thus, we can overall write that the option a.) and option c.) both can be the correct answer depending on the oxidation electrode potential and reduction electrode potential.

Further, mostly reduction electrode potential is used during calculations. So, we use (-0.591 V) as our value for electrode potential of hydrogen.