Wize University Chemistry Textbook > Electrochemistry

Reduction Potentials

Reduction Potentials

Each half reaction has a potential.
For example, the first half reaction below has a reduction potential of –3.05 V!
We got these potentials from measuring the voltage when the other half reaction was the reference electrode (H, 0 V)
In the table below, are the reactions showing reductions or oxidations? ________
Therefore, these are ________ potentials!
If we flip a reaction, the Eo potential changes signs, let's do this for the first half reaction in the table so it is clear:

Li(s) --> Li+(aq) + e- Eo = 3.05

Is this reaction now describing reduction or oxidation? ________
So do you think the potential written would now be a reduction or oxidation potential? ________ potential
In general:
If we see a more positive reduction potential:
It means something is more likely to be oxidized/reduced: ________
and will act as a(n) reducing/oxidizing agent: ________ agent
If we see a more negative reduction potential:
It means something is more likely to be oxidized/reduced: ________
and will act as a(n) reducing/oxidizing agent: ________ agent

Based on this table, answer the following questions:

1) What is the strongest oxidizing agent? ________________________________________

2) What is the weakest reducing agent? ________________________________________

3) What is the strongest reducing agent? ________________________

How Potentials Relate to Spontaneity

The overall cell potential will determine the spontaneity of redox reactions:

$\Delta G^o=-nFE_{redox}^o$

Eoredox = potential associated with a redox reaction

n = # of electrons

F = Faraday's constant (96 500C/mol e–)

ΔG° = Gibbs free energy

According to this equation, when E°redox is positive, is ΔG° positive or negative? ________
This means the reaction would be spontaneous/nonspontaneous: ________
When this is the case we have a galvanic cell (more on this soon!)
K ________ 1
When E°redox is negative, is ΔG° positive or negative? ________
This means the reaction would be spontaneous/nonspontaneous: ________
When this is the case we have an electrolytic cell (more on this soon!)
K ________ 1

Apply: If we had the following overall reaction, would it be spontaneous or not?

Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)

1) Let's write out the oxidation and reduction half reactions:

a) Oxidation half reaction:

Zn(s) --> Zn2+(aq) + 2 e-

b) Reduction half reaction:

Cu2+(aq) + 2 e- --> Cu(s)

2) Given the above reactions that are occuring, write out an oxidation potential and a reduction potential using the table values below:

From the table:

Zn2+ (aq) + 2 e– → Zn(s) E° = –0.76 V

Cu2+(aq) + 2 e– → Cu(s) E° = +0.34 V

For us,________ is being oxidized, so E°ox = ________ (changes sign)

And ________________________ is being reduced, so E°red = ________ (stays the same)

3) Solve for E°redox:

$E_{redox}^o=E_{ox}^o\ +\ E_{red}^o$

E°redox = 0.76 V + 0.34 V

E°redox = 1.10 V

4) Figure out if the reaction is spontaneous using this equation: $\Delta G^o=-nFE_{redox}^o$

Our E°redox value is positive/negative: ________
This means our ΔG° value is positive/negative: ________
Therefore, the reaction is spontaneous/nonspontaneous:________

Will the following reaction proceed spontaneously or will current need to be added externally?

$Cu_{(s)}+Ca^{2+}_{(aq)} \to Ca_{(s)} +Cu^{2+} _{(aq)}$

Standard reduction potentials

$Ca^{2+}_{(aq)} +2e^- \to Ca_{(s)} \hspace{20pt} E^o=-2.868V$

$Cu^{2+}_{(aq)} +2e^- \to Cu_{(s)} \hspace{20pt} E^o=-0.340V$

For reaction 1: -2.868 V stays the same since that equation is not flipped.
For reaction 2: we get 0.340 V since that equation is flipped according to the overall equation.
For the Erxn = -2.868 + 0.340 = -2.528 (doing the full calculation isn't necessary since we just need to figure out the sign of E to tell us about the sign of G)
We know E is negative now, so deltaG would be positive.
The electromotive force (E) is negative so this reaction would be non-spontaneous. We would have to apply at least 2.528 V to force this reaction to proceed ie, current needs to be added externally.

Based on the following reduction potentials, which of these species is the strongest reducing agent?

$\def\arraystretch{2} \begin{array}{c|c} \hline \rm Reaction & E° (V) \\ \hline 2 HCℓO (aq) + 2 H^+ (aq) + 2 e^− → Cℓ_2 (g) + 2 H_2O (ℓ) & + 1.63 \\ \hline Cℓ_2(g) + 2 e^− → 2 Cℓ^− (aq) &+ 1.36\\ \hline Br_2(ℓ) + 2 e^− → 2 Br^− (aq) & + 1.08\\ \hline I_2 (s) + 2 e^− → 2 I^− (aq) & + 0.535\\ \hline Ag+ (aq) + e^− → Ag (s) &+ 0.799\\ \hline Sn^{4+} (aq) + 2 e^− → Sn^{2+} (aq) &+ 0.15\\ \hline \end{array}$

I don't know

Extra Practice

Use the data below to answer the question:

$Zn^{2+}(aq)+2e^-⇌Zn(s)\ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ E^o=-0.74V$

$Cu^{2+}(aq)+2e^-⇌Cu(s)\ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ E^o=+0.34V$

Which one of the species below is the best reducing agent?

A concentration cell consists of two H2/H+ half-cells. Half-cell A has H2 at 0.90 atm bubbling into 0.10 M HCl. Half-cell B has H2 atm bubbling into 2.0 M HCl. Which half-cell houses the anode? What is the voltage of this cell?

Electrochemistry: Galvanic Cells

The cell pictured below is prepared under standard conditions. In this cell, what will be the anode?

$𝐶𝑢^{2+} (𝑎𝑞)+𝐻_2 (𝑔)\rightleftharpoons 𝐶𝑢 (𝑠)+2 𝐻^+ (𝑎𝑞) \qquad𝐸°=0.339 𝑉$

Which one of the species above is the best oxidizing agent?

$\def\arraystretch{1.5}\begin{array}{cc}Zn^{2+}(aq)+2e^-⇌Zn(s)&&E^o=-0.74V\\ Cu^{2+}(aq)+2e^-⇌Cu(s)&&E^o=+0.34V\\ Co^{2+}(aq)+2e^-⇌Co(s)&&E^o=-0.28V\\ Ag^+(aq)+e^-⇌Ag(s)&&E^o=+0.80V\end{array}$

Reduction Potentials

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