Standard reduction potentials

What a reduction potential measures

• Every half reaction has a tendency to pull electrons towards itself. The standard reduction potential, $E°$, is a number measuring that tendency.
• It is always quoted for the reduction direction — electrons on the left:

$$\text{Cu}^{2+}(aq) + 2e^- \rightarrow \text{Cu}(s) \qquad E° = +0.34 \text{ V}$$

• More positive $E°$ → the species pulls electrons more strongly → it is a stronger oxidant.
• More negative $E°$ → it holds its electrons less willingly in the reduced form → the reduced species is a stronger reductant.

Why "standard"

• A potential only has a fixed value under fixed conditions. Standard conditions are:
  • Concentration of all solutions 1 mol L⁻¹
  • Pressure of any gas 100 kPa
  • Temperature 25 °C (298 K)
• Change the conditions and the potential changes. At Level 3 you use the standard values as given and do not adjust them.

Measured against what

• A voltage is always a difference — you cannot measure the potential of one half cell alone.
• So every value is measured against the same reference, the standard hydrogen electrode:

$$2\text{H}^+(aq) + 2e^- \rightarrow \text{H}_2(g) \qquad E° = 0.00 \text{ V (by definition)}$$

• The zero is a convention, not a measurement. It is the origin everything else is quoted from, which is why the values can be negative.

The table

Listed from strongest oxidant at the top to strongest reductant at the bottom.

Half reaction (reduction)	$E°$ / V
$\text{F}_2 + 2e^- \rightarrow 2\text{F}^-$	+2.87
$\text{H}_2\text{O}_2 + 2\text{H}^+ + 2e^- \rightarrow 2\text{H}_2\text{O}$	+1.78
$\text{MnO}_4^- + 8\text{H}^+ + 5e^- \rightarrow \text{Mn}^{2+} + 4\text{H}_2\text{O}$	+1.51
$\text{Cl}_2 + 2e^- \rightarrow 2\text{Cl}^-$	+1.36
$\text{Cr}_2\text{O}_7^{2-} + 14\text{H}^+ + 6e^- \rightarrow 2\text{Cr}^{3+} + 7\text{H}_2\text{O}$	+1.33
$\text{O}_2 + 4\text{H}^+ + 4e^- \rightarrow 2\text{H}_2\text{O}$	+1.23
$\text{Br}_2 + 2e^- \rightarrow 2\text{Br}^-$	+1.07
$\text{Ag}^+ + e^- \rightarrow \text{Ag}$	+0.80
$\text{Fe}^{3+} + e^- \rightarrow \text{Fe}^{2+}$	+0.77
$\text{I}_2 + 2e^- \rightarrow 2\text{I}^-$	+0.54
$\text{Cu}^{2+} + 2e^- \rightarrow \text{Cu}$	+0.34
$2\text{H}^+ + 2e^- \rightarrow \text{H}_2$	0.00
$\text{Pb}^{2+} + 2e^- \rightarrow \text{Pb}$	−0.13
$\text{Sn}^{2+} + 2e^- \rightarrow \text{Sn}$	−0.14
$\text{Ni}^{2+} + 2e^- \rightarrow \text{Ni}$	−0.25
$\text{Fe}^{2+} + 2e^- \rightarrow \text{Fe}$	−0.44
$\text{Zn}^{2+} + 2e^- \rightarrow \text{Zn}$	−0.76
$2\text{H}_2\text{O} + 2e^- \rightarrow \text{H}_2 + 2\text{OH}^-$	−0.83
$\text{Al}^{3+} + 3e^- \rightarrow \text{Al}$	−1.66
$\text{Mg}^{2+} + 2e^- \rightarrow \text{Mg}$	−2.37
$\text{Na}^+ + e^- \rightarrow \text{Na}$	−2.71
$\text{Li}^+ + e^- \rightarrow \text{Li}$	−3.04

Reading the table

• The left-hand side of each half reaction is the oxidised form. Going up the table, these become stronger oxidants.

• The right-hand side is the reduced form. Going down the table, these become stronger reductants.

• F₂ is the strongest oxidant listed; Li is the strongest reductant.

• A useful check on your understanding: the metals at the bottom of the table are the reactive ones (Li, Na, Mg, Al) and the metals near the top are the unreactive ones (Ag, Cu). The reactivity series is the reduction potential table, read from the reduced side.

$E°$ does not depend on how many electrons

• If you multiply a half equation to balance electrons, you do not multiply $E°$.
• $E°$ is an intensive property — like temperature or density, it does not depend on how much material there is.

$$\text{Ag}^+ + e^- \rightarrow \text{Ag} \qquad E° = +0.80 \text{ V}$$

$$2\text{Ag}^+ + 2e^- \rightarrow 2\text{Ag} \qquad E° = +0.80 \text{ V, still}$$

• Reversing a half equation does flip the sign, however:

$$\text{Ag} \rightarrow \text{Ag}^+ + e^- \qquad E° = -0.80 \text{ V}$$