4.2 Standard potentials

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Standard electrode potentials

The voltages recorded earlier when zinc and copper were connected to a standard hydrogen electrode are in fact the standard electrode potentials for these two metals. It is important to remember that these are not absolute values, but are potentials that have been measured relative to the potential of hydrogen if the standard hydrogen electrode is taken to be zero.

Conventions and voltage sign

By convention, the hydrogen electrode is written on the left hand side of the cell. The sign of the voltage tells you the sign of the metal electrode.

In the examples we used earlier, zinc's electrode potential is actually -0.76 and copper is +0.34. So, if a metal has a negative standard electrode potential, it means it forms ions easily. The more negative the value, the easier it is for that metal to form ions. If a metal has a positive standard electrode potential, it means it does not form ions easily. This will be explained in more detail below.

Luckily for us, we do not have to calculate the standard electrode potential for every metal. This has been done already and the results are recorded in a table of standard electrode potentials ( [link] ).

Standard Electrode Potentials

Half-Reaction	$E^{0} V$
$L i^{+} + e^{-} ⇌ L i$	-3.04
$K^{+} + e^{-} ⇌ K$	-2.92
$B a^{2 +} + 2 e^{-} ⇌ B a$	-2.90
$C a^{2 +} + 2 e^{-} ⇌ C a$	-2.87
$N a^{+} + e^{-} ⇌ N a$	-2.71
$M g^{2 +} + 2 e^{-} ⇌ M g$	-2.37
$M n^{2 +} + 2 e^{-} ⇌ M n$	-1.18
$2 H 2 O + 2 e^{-} ⇌ H_{2} (g) + 2 O H^{-}$	-0.83
$Z n^{2 +} + 2 e^{-} ⇌ Z n$	-0.76
$C r^{2 +} + 2 e^{-} ⇌ C r$	-0.74
$F e^{2 +} + 2 e^{-} ⇌ F e$	-0.44
$C r^{3 +} + 3 e^{-} ⇌ C r$	-0.41
$C d^{2 +} + 2 e^{-} ⇌ C d$	-0.40
$C o^{2 +} + 2 e^{-} ⇌ C o$	-0.28
$N i^{2 +} + 2 e^{-} ⇌ N i$	-0.25
$S n^{2 +} + 2 e^{-} ⇌ S n$	-0.14
$P b^{2 +} + 2 e^{-} ⇌ P b$	-0.13
$F e^{3 +} + 3 e^{-} ⇌ F e$	-0.04
$2 H^{+} + 2 e^{-} ⇌ H_{2} (g)$	0.00
$S + 2 H^{+} + 2 e^{-} ⇌ H_{2} S (g)$	0.14
$S n^{4 +} + 2 e^{-} ⇌ S n^{2 +}$	0.15
$C u^{2 +} + e^{-} ⇌ C u^{+}$	0.16
$S O_{4}^{2 +} + 4 H^{+} + 2 e^{-} ⇌ S O_{2} (g) + 2 H_{2} O$	0.17
$C u^{2 +} + 2 e^{-} ⇌ C u$	0.34
$2 H_{2} O + O_{2} + 4 e^{-} ⇌ 4 O H^{-}$	0.40
$C u^{+} + e^{-} ⇌ C u$	0.52
$I_{2} + 2 e^{-} ⇌ 2 I^{-}$	0.54
$O_{2} (g) + 2 H^{+} + 2 e^{-} ⇌ H_{2} O_{2}$	0.68
$F e^{3 +} + e^{-} ⇌ F e^{2 +}$	0.77
$N O_{3}^{-} + 2 H^{+} + e^{-} ⇌ N O_{2} (g) + H_{2} O$	0.78
$H g^{2 +} + 2 e^{-} ⇌ H g (l)$	0.78
$A g^{+} + e^{-} ⇌ A g$	0.80
$N O_{3}^{-} + 4 H^{+} + 3 e^{-} ⇌ N O (g) + 2 H_{2} O$	0.96
$B r_{2} + 2 e^{-} ⇌ 2 B r^{-}$	1.06
$O_{2} (g) + 4 H^{+} + 4 e^{-} ⇌ 2 H_{2} O$	1.23
$M n O_{2} + 4 H^{+} + 2 e^{-} ⇌ M n^{2 +} + 2 H_{2} O$	1.28
$C r_{2} O_{7}^{2 -} + 14 H^{+} + 6 e^{-} ⇌ 2 C r^{3 +} + 7 H_{2} O$	1.33
$C l_{2} + 2 e^{-} ⇌ 2 C l^{-}$	1.36
$A u^{3 +} + 3 e^{-} ⇌ A u$	1.50
$M n O_{4}^{-} + 8 H^{+} + 5 e^{-} ⇌ M n^{2 +} + 4 H_{2} O$	1.52
$C o^{3 +} + e^{-} ⇌ C o^{2 +}$	1.82
$F_{2} + 2 e^{-} ⇌ 2 F^{-}$	2.87

A few examples from the table are shown in [link] . These will be used to explain some of the trends in the table of electrode potentials.

A few examples from the table of standard electrode potentials

Half-Reaction	$E^{0} V$
$L i^{+} + e^{-} ⇌ L i$	-3.04
$Z n^{2 +} + 2 e^{-} ⇌ Z n$	-0.76
$F e^{3 +} + 3 e^{-} ⇌ F e$	-0.04
$2 H^{+} + 2 e^{-} ⇌ H_{2} (g)$	0.00
$C u^{2 +} + 2 e^{-} ⇌ C u$	0.34
$H g^{2 +} + 2 e^{-} ⇌ H g (l)$	0.78
$A g^{+} + e^{-} ⇌ A g$	0.80

Refer to [link] and notice the following trends:

Metals at the top of series (e.g. Li) have more negative values. This means they ionise easily, in other words, they release electrons easily. These metals are easily oxidised and are therefore good reducing agents .
Metal ions at the bottom of the table are good at picking up electrons. They are easily reduced and are therefore good oxidising agents .
The reducing ability (i.e. the ability to act as a reducing agent) of the metals in the table increases as you move up in the table.
The oxidising ability of metals increases as you move down in the table.

The following half-reactions take place in an electrochemical cell:

${Cu}^{2 +} + 2 e^{-} ⇌ Cu$

${Ag}^{-} + e^{-} ⇌ Ag$

Which of these reactions will be the oxidation half-reaction in the cell?
Which of these reactions will be the reduction half-reaction in the cell?

Determine the electrode potential for each metal
From the table of standard electrode potentials, the electrode potential for the copper half-reaction is +0.34 V. The electrode potential for the silver half-reaction is +0.80 V.
Use the electrode potential values to determine which metal is oxidised and which is reduced
Both values are positive, but silver has a higher positive electrode potential than copper. This means that silver does not form ions easily, in other words, silver is more likely to be reduced . Copper is more likely to be oxidised and to form ions more easily than silver. Copper is the oxidation half-reaction and silver is the reduction half-reaction.

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Source: OpenStax, Siyavula textbooks: grade 12 physical science. OpenStax CNX. Aug 03, 2011 Download for free at http://cnx.org/content/col11244/1.2

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