2.7 Relative Reactivity as Oxidizing or Reducing Agents

Standard electrode potentials ($E^{\circ }$) provide a quantitative basis for comparing the relative reactivity of species as oxidizing or reducing agents.

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Standard Hydrogen Electrode (SHE)

Before comparing electrode potentials, a reference point is needed. The Standard Hydrogen Electrode (SHE) serves this purpose.

Construction:

• A platinum electrode (inert, does not react)
• Immersed in 1 mol dm${}^{-3}$ H${}^{+}$(aq) solution
• H${}_2$ gas bubbled over the electrode at 1 atm pressure
• Temperature: 298 K (25°C)

Assigned value: $E^{\circ }=0.00\text{ V}$ (by international convention)

The half-reaction at the SHE is: $2H^{+}(aq)+2e^{-}\rightleftharpoons H_2(g)E^{\circ }=0.00\text{ V}$

Measuring standard electrode potentials: To measure the $E^{\circ }$ of any half-cell, it is connected to the SHE under standard conditions (1 mol dm${}^{-3}$ ion concentration, 298 K, 1 atm). The measured cell voltage equals the $E^{\circ }$ of the half-cell being tested.

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Standard Electrode Potential and Reactivity

The standard reduction potential ($E^{\circ }$) measures the tendency of a species to be reduced (gain electrons) under standard conditions.

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Oxidizing Agents

An oxidizing agent gains electrons and is itself reduced. The species with the more positive $E^{\circ }$ is the stronger oxidizing agent.

Example — Halogens as oxidizing agents:

Trend: Oxidizing power decreases down Group VII: $F_2>C{l}_2>B{r}_2>I_2$

$F_2$ is the strongest oxidizing agent because it has the highest electronegativity and smallest atomic radius, giving it the greatest tendency to attract electrons.

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Reducing Agents

A reducing agent loses electrons and is itself oxidized. The species with the more negative $E^{\circ }$ (for the reduction half-reaction) is the stronger reducing agent.

Example — Halide ions as reducing agents:

The reducing power of halide ions is the reverse of the oxidizing power of halogens: $I^{-}>B{r}^{-}>C{l}^{-}>F^{-}$

• $I^{-}$ has the largest ionic radius → outermost electron is furthest from nucleus → held least tightly → most easily lost → strongest reducing agent
• $F^{-}$ has the smallest ionic radius → electron held most tightly → weakest reducing agent

Note: $F^{-}$ is such a weak reducing agent that it cannot be oxidized by any common reagent in aqueous solution.

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Predicting Spontaneity of Redox Reactions

Using $E^{\circ }$ values, we can predict whether a redox reaction will occur spontaneously:

$E_{cell}^{\circ }=E_{cathode}^{\circ }-E_{anode}^{\circ }$

• If $E_{cell}^{\circ }>0$: reaction is spontaneous (feasible)
• If $E_{cell}^{\circ }<0$: reaction is non-spontaneous

Rule: A species higher in the electrochemical series (more positive $E^{\circ }$) will oxidize a species lower in the series (more negative $E^{\circ }$).

Example: Will $C{l}_2$ oxidize $B{r}^{-}$?

$C{l}_2+2e^{-}\to 2C{l}^{-}{E}^{\circ }=+1.36\text{ V (cathode)}$ $2B{r}^{-}\to B{r}_2+2e^{-}{E}^{\circ }=-1.07\text{ V (anode, reversed)}$ $E_{cell}^{\circ }=+1.36-(+1.07)=+0.29\text{ V}>0✓\text{ Spontaneous}$

Yes — $C{l}_2$ will displace $B{r}^{-}$ from solution.

Example: Will $I_2$ oxidize $C{l}^{-}$?

$I_2+2e^{-}\to 2I^{-}{E}^{\circ }=+0.54\text{ V (cathode)}$ $2C{l}^{-}\to C{l}_2+2e^{-}{E}^{\circ }=-1.36\text{ V (anode, reversed)}$ $E_{cell}^{\circ }=+0.54-(+1.36)=-0.82\text{ V}<0\times \text{ Non-spontaneous}$

No — $I_2$ cannot oxidize $C{l}^{-}$.

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Summary Table