Standard Electrode Potentials

Standard electrode potentials

In order to use the hydrogen electrode, it needs to be attached to the electrode system that you are investigating. For example, if you are trying to determine the electrode potential of copper, you will need to connect the copper half-cell to the hydrogen electrode; if you are trying to determine the electrode potential of zinc, you will need to connect the zinc half-cell to the hydrogen electrode and so on. Let's look at the examples of zinc and copper in more detail.

Zinc

Zinc has a greater tendency than hydrogen to form ions (to be oxidised), so if the standard hydrogen electrode is connected to the zinc half-cell, the $\color{blue}{\textbf{zinc}}$ will be relatively $\color{blue}{\textbf{more negative}}$ because the electrons that are released when zinc is oxidised will accumulate on the metal.

$\text{Zn}^{2+}(\text{aq}) + 2\text{e}^{-}$ $\to$ $\text{Zn}(\text{s})$

$2\text{H}^{+}(\text{aq}) + 2\text{e}^{-}$ $\to$ $\text{H}_{2}(\text{g})$

The solid zinc is more likely to form zinc ions than the hydrogen gas is to form ions. A simplified representation of the cell is shown in the figure below.

When zinc is connected to the standard hydrogen electrode, relatively few electrons build up on the platinum (hydrogen) electrode. There are lots of electrons on the zinc electrode.

Fact:

When determining standard electrode reduction potentials the standard hydrogen electrode is considered to be on the left ($\text{Pt}(\text{s})|\text{H}_{2}(\text{g}),\text{H}^{+}(\text{aq})||$). So a negative value means that the other element or compound has a greater tendency to oxidise, and a positive value means that the other element or compound has a greater tendency to be reduced.

The voltmeter measures the potential difference between the charge on these electrodes. In this case, the voltmeter would read $-\text{0.76}$ $\text{V}$ as the $\text{Zn}$ electrode has a relatively higher number of electrons.

Copper

Copper has a lower tendency than hydrogen to form ions, so if the standard hydrogen electrode is connected to the copper half-cell, the $\color{red}{\textbf{copper}}$ will be relatively $\color{red}{\textbf{less negative}}$.

$\text{Cu}^{2+}(\text{aq}) + 2\text{e}^{-}$ $\to$ $\text{Cu}(\text{s})$

$2\text{H}^{+}(\text{aq}) + 2\text{e}^{-}$ $\to$ $\text{H}_{2}(\text{g})$

The copper ions are more likely to form solid copper than the hydrogen ions are to form hydrogen gas. A simplified representation of the cell is shown in the figure below.

When copper is connected to the standard hydrogen electrode, relatively few electrons build up on the copper electrode. There are lots of electrons on the hydrogen electrode.

The voltmeter measures the potential difference between the charge on these electrodes. In this case, the voltmeter would read $\text{+0.34}$ $\text{V}$ as the $\text{Cu}$ electrode has a relatively lower number of electrons.

The voltages recorded 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.

Tip:

By convention, we always write the reduction half-reaction when giving the standard electrode potential.

Tip:

In the examples we used earlier, zinc's electrode reduction potential is $-\text{0.76}$ and copper's is $\text{+0.34}$. So, if an element or compound has a negative standard electrode reduction potential, it means it forms ions easily. The more negative the value, the easier it is for that element or compound to form ions (be oxidised, and be a reducing agent). If an element or compound has a positive standard electrode potential, it means it does not form ions as easily.

Luckily for us, we do not have to determine the standard electrode potential for every metal. This has been done already and the results are recorded in a table of standard electrode potentials. The table below is presented as the standard electrode reduction potentials.

Half-Reaction	E° V
$\text{Li}^{+} + \text{e}^{-}$ $\rightleftharpoons$ $\text{Li}$	$-\text{3.04}$
$\text{K}^{+} + \text{e}^{-}$ $\rightleftharpoons$ $\text{K}$	$-\text{2.92}$
$\text{Ba}^{2+} + 2\text{e}^{-}$ $\rightleftharpoons$ $\text{Ba}$	$-\text{2.90}$
$\text{Ca}^{2+} + 2\text{e}^{-}$ $\rightleftharpoons$ $\text{Ca}$	$-\text{2.87}$
$\text{Na}^{+} + \text{e}^{-}$ $\rightleftharpoons$ $\text{Na}$	$-\text{2.71}$
$\text{Mg}^{2+} + 2\text{e}^{-}$ $\rightleftharpoons$ $\text{Mg}$	$-\text{2.37}$
$\text{Al}^{3+} + 3\text{e}^{-}$ $\rightleftharpoons$ $\text{Al}$	$-\text{1.66}$
$\text{Mn}^{2+} + 2\text{e}^{-}$ $\rightleftharpoons$ $\text{Mn}$	$-\text{1.18}$
$2\text{H}_{2}\text{O} + 2\text{e}^{-}$ $\rightleftharpoons$ $\text{H}_{2}(\text{g}) + 2\text{OH}^{-}$	$-\text{0.83}$
$\text{Zn}^{2+} + 2\text{e}^{-}$ $\rightleftharpoons$ $\text{Zn}$	$-\text{0.76}$
$\text{Cr}^{2+} + 2\text{e}^{-}$ $\rightleftharpoons$ $\text{Cr}$	$-\text{0.74}$
$\text{Fe}^{2+} + 2\text{e}^{-}$ $\rightleftharpoons$ $\text{Fe}$	$-\text{0.44}$
$\text{Cr}^{3+} + 3\text{e}^{-}$ $\rightleftharpoons$ $\text{Cr}$	$-\text{0.41}$
$\text{Cd}^{2+} + 2\text{e}^{-}$ $\rightleftharpoons$ $\text{Cd}$	$-\text{0.40}$
$\text{Co}^{2+} + 2\text{e}^{-}$ $\rightleftharpoons$ $\text{Co}$	$-\text{0.28}$
$\text{Ni}^{2+} + 2\text{e}^{-}$ $\rightleftharpoons$ $\text{Ni}$	$-\text{0.25}$
$\text{Sn}^{2+} + 2\text{e}^{-}$ $\rightleftharpoons$ $\text{Sn}$	$-\text{0.14}$
$\text{Pb}^{2+} + 2\text{e}^{-}$ $\rightleftharpoons$ $\text{Pb}$	$-\text{0.13}$
$\text{Fe}^{3+} + 3\text{e}^{-}$ $\rightleftharpoons$ $\text{Fe}$	$-\text{0.04}$
$2\text{H}^{+} + 2\text{e}^{-}$ $\rightleftharpoons$ $\text{H}_{2}(\text{g})$	0.00
$\text{S} + 2\text{H}^{+} + 2\text{e}^{-}$ $\rightleftharpoons$ $\text{H}_{2}\text{S}(\text{g})$	$\text{+0.14}$
$\text{Sn}^{4+} + 2\text{e}^{-}$ $\rightleftharpoons$ $\text{Sn}^{2+}$	$\text{+0.15}$
$\text{Cu}^{2+} + \text{e}^{-}$ $\rightleftharpoons$ $\text{Cu}^{+}$	$\text{+0.16}$
$\text{SO}_{4}^{2-} + 4\text{H}^{+} + 2\text{e}^{-}$ $\rightleftharpoons$ $\text{SO}_{2}(\text{g}) + 2\text{H}_{2}\text{O}$	$\text{+0.17}$
$\text{Cu}^{2+} + 2\text{e}^{-}$ $\rightleftharpoons$ $\text{Cu}$	$\text{+0.34}$
$2\text{H}_{2}\text{O} + \text{O}_{2} + 4\text{e}^{-}$ $\rightleftharpoons$ $4\text{OH}^{-}$	$\text{+0.40}$
$\text{Cu}^{+} + \text{e}^{-}$ $\rightleftharpoons$ $\text{Cu}$	$\text{+0.52}$
$\text{I}_{2} + 2\text{e}^{-}$ $\rightleftharpoons$ $2\text{I}^{-}$	$\text{+0.54}$
$\text{O}_{2}(\text{g}) + 2\text{H}^{+} + 2\text{e}^{-}$ $\rightleftharpoons$ $\text{H}_{2}\text{O}_{2}$	$\text{+0.68}$
$\text{Fe}^{3+} + \text{e}^{-}$ $\rightleftharpoons$ $\text{Fe}^{2+}$	$\text{+0.77}$
$\text{NO}_{3}^{-} + 2\text{H}^{+} + \text{e}^{-}$ $\rightleftharpoons$ $\text{NO}_{2}(\text{g}) + \text{H}_{2}\text{O}$	$\text{+0.78}$
$\text{Hg}^{2+} + 2\text{e}^{-}$ $\rightleftharpoons$ $\text{Hg}(\text{l})$	$\text{+0.78}$
$\text{Ag}^{+} + \text{e}^{-}$ $\rightleftharpoons$ $\text{Ag}$	$\text{+0.80}$
$\text{NO}_{3}^{-} + 4\text{H}^{+} + 3\text{e}^{-}$ $\rightleftharpoons$ $\text{NO}(\text{g}) + 2\text{H}_{2}\text{O}$	$\text{+0.96}$
$\text{Br}_{2} + 2\text{e}^{-}$ $\rightleftharpoons$ $2\text{Br}^{-}$	$\text{+1.06}$
$\text{O}_{2}(\text{g}) + 4\text{H}^{+} + 4\text{e}^{-}$ $\rightleftharpoons$ $2\text{H}_{2}\text{O}$	$\text{+1.23}$
$\text{MnO}_{2} + 4\text{H}^{+} + 2\text{e}^{-}$ $\rightleftharpoons$ $\text{Mn}^{2+} + 2\text{H}_{2}\text{O}$	$\text{+1.28}$
$\text{Cr}_{2}\text{O}_{7}^{2-} + 14\text{H}^{+} + 6\text{e}^{-}$ $\rightleftharpoons$ $2\text{Cr}^{3+} + 7\text{H}_{2}\text{O}$	$\text{+1.33}$
$\text{Cl}_{2} + 2\text{e}^{-}$ $\rightleftharpoons$ $2\text{Cl}^{-}$	$\text{+1.36}$
$\text{Au}^{3+} + 3\text{e}^{-}$ $\rightleftharpoons$ $\text{Au}$	$\text{+1.50}$
$\text{MnO}_{4}^{-} + 8\text{H}^{+} + 5\text{e}^{-}$ $\rightleftharpoons$ $\text{Mn}^{2+} + 4\text{H}_{2}\text{O}$	$\text{+1.52}$
$\text{Co}^{3+} + \text{e}^{-}$ $\rightleftharpoons$ $\text{Co}^{2+}$	$\text{+1.82}$
$\text{F}_{2} + 2\text{e}^{-}$ $\rightleftharpoons$ $2\text{F}^{-}$	$\text{+2.87}$

Table: Table of standard electrode (reduction) potentials.

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

Half-Reaction	E° V
$\text{Li}^{+} + \text{e}^{-}$ $\rightleftharpoons$ $\text{Li}$	$-\text{3.04}$
$\text{Mg}^{2+} + 2\text{e}^{-}$ $\rightleftharpoons$ $\text{Mg}$	$-\text{2.37}$
$\text{Zn}^{2+} + 2\text{e}^{-}$ $\rightleftharpoons$ $\text{Zn}$	$-\text{0.76}$
$\text{Fe}^{3+} + 3\text{e}^{-}$ $\rightleftharpoons$ $\text{Fe}$	$-\text{0.4}$
$\text{Pb}^{2+} + 2\text{e}^{-}$ $\rightleftharpoons$ $\text{Pb}$	$-\text{0.13}$
$2\text{H}^{+} + 2\text{e}^{-}$ $\rightleftharpoons$ $\text{H}_{2}(\text{g})$	0.00
$\text{Cu}^{2+} + 2\text{e}^{-}$ $\rightleftharpoons$ $\text{Cu}$	$\text{+0.34}$
$\text{Ag}^{+} + \text{e}^{-}$ $\rightleftharpoons$ $\text{Ag}$	$\text{+0.80}$
$\text{Au}^{3+} + 3\text{e}^{-}$ $\rightleftharpoons$ $\text{Au}$	$\text{+1.50}$

Table: A few examples of standard electrode potentials.

A large negative value (e.g. $\text{Li}^{+} + \text{e}^{-}$ $\rightleftharpoons$ $\text{Li}$) means that the element or compound ionises easily, in other words, it releases electrons easily. This element or compound is $\color{blue}{\textbf{easily oxidised}}$ and is therefore a good $\color{blue}{\textbf{reducing agent}}$.
A large positive value (e.g. $\text{Au}^{3+} + \text{e}^{-}$ $\rightleftharpoons$ $\text{Au}$) means that the element or compound gains electrons easily. This element or compound is $\color{red}{\textbf{easily reduced}}$ and is therefore a good $\color{red}{\textbf{oxidising agent}}$.
The $\color{blue}{\textbf{reducing ability}}$ (i.e. the ability to act as a reducing agent) of the elements or compounds in the table decreases as you move down in the table.
The $\color{red}{\textbf{oxidising ability}}$ of elements or compounds increases as you move down in the table.

The voltages produced in the ice cube tray experiment may be significantly different from the E$^{\circ}$ calculations. This is because the salt-bridge is not that effective (being only a piece of string soaked in electrolyte). To get the best results use a saturated sodium nitrate solution when soaking the string.

Optional Experiment: Ice cube tray redox experiment

Aim

To determine relative reactivity of the respective metals and to gain understanding of the workings of a simple electrochemical cell.

Apparatus

Ice cube tray, voltmeter and connecting wires.
Lead ($\text{Pb}$), magnesium ($\text{Mg}$), zinc ($\text{Zn}$), copper ($\text{Cu}$) strips.
$\text{1}$ $\text{mol.dm$^{-3}$}$ solutions of lead (e.g. $\text{PbSO}_{4}$), magnesium (e.g. $\text{MgSO}_{4}$), zinc (e.g. $\text{ZnSO}_{4}$) and copper (e.g. $\text{CuSO}_{4}$).
String soaked in a sodium nitrate ($\text{NaNO}_{3}$) solution.

Pre-knowlege

Electrons move from the anode to the cathode. Conventional current moves from the cathode to the anode - therefore the positive terminal of the voltmeter will be on the cathode and the negative terminal will be on the anode.

Method

Place approximately $\text{15}$ $\text{cm$^{3}$}$ of the $\text{Pb}$, $\text{Zn}$, $\text{Cu}$ and $\text{Mg}$ solutions into four different ice cube depressions.

These should not be next to each other to avoid mixing of solutions.
Attach two different metals to the crocodile clips. Drape the wet string across the two solutions being used (so that each end of the string is in a solution). Then place the metal into its respective ion solution.

i.e. zinc electrode must go into the $\text{Zn}^{2+}$ solution, copper electrode must go into the $\text{Cu}^{2+}$ solution.

Use the cell combinations in the following order:

Pb/Zn; Pb/Cu; Pb/Mg; Zn/Cu; Zn/Mg; Cu/Mg
Determine the combinations of metals that give a positive reading.
Draw up a table that shows:
- the combination of the metals
- which metal is the anode in that pair of metals
- which metal is the cathode in that pair of metals
Metal combination

Anode

Cathode
Use this table to rank the metals from the strongest reducing agent (rank from the strongest to the weakest).
For each combination - write down the reduction half-reaction and oxidation half-reaction and then the overall cell reaction.
Note all observations for each cell.

Questions

Explain why the voltages appear to be lower/higher than expected.
What is the purpose of the string?

Conclusions

Depending on the electrode potential of each metal, the same metal could be the anode in one reaction and the cathode in another reaction. This can be seen by the positive, or negative, reading on the voltmeter.

For example, lead is more likely to be reduced than zinc, therefore in that pair lead will be the cathode and zinc will be the anode. However, lead is more likely to be oxidised than copper, therefore in that pair copper will be the cathode and lead will be the anode.

Half-Reaction	E° V
\(\text{Li}^{+} + \text{e}^{-}\) \(\rightleftharpoons\) \(\text{Li}\)	\(-\text{3.04}\)
\(\text{K}^{+} + \text{e}^{-}\) \(\rightleftharpoons\) \(\text{K}\)	\(-\text{2.92}\)
\(\text{Ba}^{2+} + 2\text{e}^{-}\) \(\rightleftharpoons\) \(\text{Ba}\)	\(-\text{2.90}\)
\(\text{Ca}^{2+} + 2\text{e}^{-}\) \(\rightleftharpoons\) \(\text{Ca}\)	\(-\text{2.87}\)
\(\text{Na}^{+} + \text{e}^{-}\) \(\rightleftharpoons\) \(\text{Na}\)	\(-\text{2.71}\)
\(\text{Mg}^{2+} + 2\text{e}^{-}\) \(\rightleftharpoons\) \(\text{Mg}\)	\(-\text{2.37}\)
\(\text{Al}^{3+} + 3\text{e}^{-}\) \(\rightleftharpoons\) \(\text{Al}\)	\(-\text{1.66}\)
\(\text{Mn}^{2+} + 2\text{e}^{-}\) \(\rightleftharpoons\) \(\text{Mn}\)	\(-\text{1.18}\)
\(2\text{H}_{2}\text{O} + 2\text{e}^{-}\) \(\rightleftharpoons\) \(\text{H}_{2}(\text{g}) + 2\text{OH}^{-}\)	\(-\text{0.83}\)
\(\text{Zn}^{2+} + 2\text{e}^{-}\) \(\rightleftharpoons\) \(\text{Zn}\)	\(-\text{0.76}\)
\(\text{Cr}^{2+} + 2\text{e}^{-}\) \(\rightleftharpoons\) \(\text{Cr}\)	\(-\text{0.74}\)
\(\text{Fe}^{2+} + 2\text{e}^{-}\) \(\rightleftharpoons\) \(\text{Fe}\)	\(-\text{0.44}\)
\(\text{Cr}^{3+} + 3\text{e}^{-}\) \(\rightleftharpoons\) \(\text{Cr}\)	\(-\text{0.41}\)
\(\text{Cd}^{2+} + 2\text{e}^{-}\) \(\rightleftharpoons\) \(\text{Cd}\)	\(-\text{0.40}\)
\(\text{Co}^{2+} + 2\text{e}^{-}\) \(\rightleftharpoons\) \(\text{Co}\)	\(-\text{0.28}\)
\(\text{Ni}^{2+} + 2\text{e}^{-}\) \(\rightleftharpoons\) \(\text{Ni}\)	\(-\text{0.25}\)
\(\text{Sn}^{2+} + 2\text{e}^{-}\) \(\rightleftharpoons\) \(\text{Sn}\)	\(-\text{0.14}\)
\(\text{Pb}^{2+} + 2\text{e}^{-}\) \(\rightleftharpoons\) \(\text{Pb}\)	\(-\text{0.13}\)
\(\text{Fe}^{3+} + 3\text{e}^{-}\) \(\rightleftharpoons\) \(\text{Fe}\)	\(-\text{0.04}\)
\(2\text{H}^{+} + 2\text{e}^{-}\) \(\rightleftharpoons\) \(\text{H}_{2}(\text{g})\)	0.00
\(\text{S} + 2\text{H}^{+} + 2\text{e}^{-}\) \(\rightleftharpoons\) \(\text{H}_{2}\text{S}(\text{g})\)	\(\text{+0.14}\)
\(\text{Sn}^{4+} + 2\text{e}^{-}\) \(\rightleftharpoons\) \(\text{Sn}^{2+}\)	\(\text{+0.15}\)
\(\text{Cu}^{2+} + \text{e}^{-}\) \(\rightleftharpoons\) \(\text{Cu}^{+}\)	\(\text{+0.16}\)
\(\text{SO}_{4}^{2-} + 4\text{H}^{+} + 2\text{e}^{-}\) \(\rightleftharpoons\) \(\text{SO}_{2}(\text{g}) + 2\text{H}_{2}\text{O}\)	\(\text{+0.17}\)
\(\text{Cu}^{2+} + 2\text{e}^{-}\) \(\rightleftharpoons\) \(\text{Cu}\)	\(\text{+0.34}\)
\(2\text{H}_{2}\text{O} + \text{O}_{2} + 4\text{e}^{-}\) \(\rightleftharpoons\) \(4\text{OH}^{-}\)	\(\text{+0.40}\)
\(\text{Cu}^{+} + \text{e}^{-}\) \(\rightleftharpoons\) \(\text{Cu}\)	\(\text{+0.52}\)
\(\text{I}_{2} + 2\text{e}^{-}\) \(\rightleftharpoons\) \(2\text{I}^{-}\)	\(\text{+0.54}\)
\(\text{O}_{2}(\text{g}) + 2\text{H}^{+} + 2\text{e}^{-}\) \(\rightleftharpoons\) \(\text{H}_{2}\text{O}_{2}\)	\(\text{+0.68}\)
\(\text{Fe}^{3+} + \text{e}^{-}\) \(\rightleftharpoons\) \(\text{Fe}^{2+}\)	\(\text{+0.77}\)
\(\text{NO}_{3}^{-} + 2\text{H}^{+} + \text{e}^{-}\) \(\rightleftharpoons\) \(\text{NO}_{2}(\text{g}) + \text{H}_{2}\text{O}\)	\(\text{+0.78}\)
\(\text{Hg}^{2+} + 2\text{e}^{-}\) \(\rightleftharpoons\) \(\text{Hg}(\text{l})\)	\(\text{+0.78}\)
\(\text{Ag}^{+} + \text{e}^{-}\) \(\rightleftharpoons\) \(\text{Ag}\)	\(\text{+0.80}\)
\(\text{NO}_{3}^{-} + 4\text{H}^{+} + 3\text{e}^{-}\) \(\rightleftharpoons\) \(\text{NO}(\text{g}) + 2\text{H}_{2}\text{O}\)	\(\text{+0.96}\)
\(\text{Br}_{2} + 2\text{e}^{-}\) \(\rightleftharpoons\) \(2\text{Br}^{-}\)	\(\text{+1.06}\)
\(\text{O}_{2}(\text{g}) + 4\text{H}^{+} + 4\text{e}^{-}\) \(\rightleftharpoons\) \(2\text{H}_{2}\text{O}\)	\(\text{+1.23}\)
\(\text{MnO}_{2} + 4\text{H}^{+} + 2\text{e}^{-}\) \(\rightleftharpoons\) \(\text{Mn}^{2+} + 2\text{H}_{2}\text{O}\)	\(\text{+1.28}\)
\(\text{Cr}_{2}\text{O}_{7}^{2-} + 14\text{H}^{+} + 6\text{e}^{-}\) \(\rightleftharpoons\) \(2\text{Cr}^{3+} + 7\text{H}_{2}\text{O}\)	\(\text{+1.33}\)
\(\text{Cl}_{2} + 2\text{e}^{-}\) \(\rightleftharpoons\) \(2\text{Cl}^{-}\)	\(\text{+1.36}\)
\(\text{Au}^{3+} + 3\text{e}^{-}\) \(\rightleftharpoons\) \(\text{Au}\)	\(\text{+1.50}\)
\(\text{MnO}_{4}^{-} + 8\text{H}^{+} + 5\text{e}^{-}\) \(\rightleftharpoons\) \(\text{Mn}^{2+} + 4\text{H}_{2}\text{O}\)	\(\text{+1.52}\)
\(\text{Co}^{3+} + \text{e}^{-}\) \(\rightleftharpoons\) \(\text{Co}^{2+}\)	\(\text{+1.82}\)
\(\text{F}_{2} + 2\text{e}^{-}\) \(\rightleftharpoons\) \(2\text{F}^{-}\)	\(\text{+2.87}\)

Half-Reaction	E° V
\(\text{Li}^{+} + \text{e}^{-}\) \(\rightleftharpoons\) \(\text{Li}\)	\(-\text{3.04}\)
\(\text{Mg}^{2+} + 2\text{e}^{-}\) \(\rightleftharpoons\) \(\text{Mg}\)	\(-\text{2.37}\)
\(\text{Zn}^{2+} + 2\text{e}^{-}\) \(\rightleftharpoons\) \(\text{Zn}\)	\(-\text{0.76}\)
\(\text{Fe}^{3+} + 3\text{e}^{-}\) \(\rightleftharpoons\) \(\text{Fe}\)	\(-\text{0.4}\)
\(\text{Pb}^{2+} + 2\text{e}^{-}\) \(\rightleftharpoons\) \(\text{Pb}\)	\(-\text{0.13}\)
\(2\text{H}^{+} + 2\text{e}^{-}\) \(\rightleftharpoons\) \(\text{H}_{2}(\text{g})\)	0.00
\(\text{Cu}^{2+} + 2\text{e}^{-}\) \(\rightleftharpoons\) \(\text{Cu}\)	\(\text{+0.34}\)
\(\text{Ag}^{+} + \text{e}^{-}\) \(\rightleftharpoons\) \(\text{Ag}\)	\(\text{+0.80}\)
\(\text{Au}^{3+} + 3\text{e}^{-}\) \(\rightleftharpoons\) \(\text{Au}\)	\(\text{+1.50}\)

Metal combination	Anode	Cathode

Standard Electrode Potentials