Title: Lesson 7 Standard Electrode Potential
Learning Objectives:
– Describe the standard hydrogen electrode
– Define the term standard electrode potential
– Use standard electrode potentials to calculate the potential of a cell
– Use standard electrode potentials to determine the feasibility of a reaction
Refresh
Which processes occur during the electrolysis of molten sodium chloride?
I.
II.
III.
A.
B.
C.
D.

Sodium and chloride ions move through the electrolyte.
Electrons move through the external circuit.
Oxidation takes place at the anode.
I and II only
I and III only
II and III only
I, II and III
Justify your answer.
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General Reminders…
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Comparisons of half-cell electrode potentials need a
reference point

Potential difference is known as the electromotive force (EMF)

Electrons tend to flow from half-cells: more negative potential  more positive potential

Potential generated is called the cell potential or electrode potential… Symbol is E.

Magnitude of this voltage depends on the difference in tendency of reduction of the halfcells.

Can’t measure an isolated half cell (no electron flow)

So we measure against a fixed reference point… STANDARD HYDROGEN
ELECTRODE
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Standard Electrode Potential, Eo
Half Cell
• This is the potential of a
standard electrode
relative to the standard
hydrogen electrode.
Standard Electrode Potential,
Eo / V
H+(aq) + e- ⇌ ½ H2(g)
0.00
Li+(aq)
-3.04
+
e-
⇌ Li(s)
Mn2+(aq) + 2e- ⇌ Mn(s)
-1.19
Cu2+(aq) + 2e- ⇌ Cu(s)
+0.34
½ Br2(l) + e- ⇌ Br-(aq)
+1.07
• Always measure the
potential of the reduction

• Measured in Volts, V
• Full table in the data
booklet
Look at the table in the data booklet:
 What trends do you notice?
 How do the values relate to your ideas
of reactivity?
 How do the values compare to the
reactivity series you constructed
earlier?
Comparisons of half-cell electrode potentials need
a reference point
• Potential difference is known as the electromotive force (EMF)
• Electrons tend to flow from half-cells: more negative potential  more positive potential
• Potential generated is called the cell potential or electrode potential… Symbol is E.
• Magnitude of this voltage depends on the difference in tendency of reduction of the half-cells.
• Can’t measure an isolated half cell (no electron flow)
• So we measure against a fixed reference point… STANDARD HYDROGEN ELECTRODE
The standard hydrogen electrode
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The Standard Hydrogen Electrode
 Platinum is used as the conducting
metal because it is fairly inert, will not
ionize and can b a catalyst for
reduction
 Surface of the metal is coated in finely
divided platinum to increase surface
area for absorption of hydrogen gas
 The electrode is bathed alternately in
H2(g) and H+(aq), setting up an
equilibrium
Forward - Reduction
Standard electrode potential defined as 0.00 V…
Backward - Oxidation
So we can measure and compare electrode potentials
or other half-cells
Measuring Standard Electrode Potentials
• Temperature = 298K
• Pressure = 100 kPa or 1 atm
• Substances must be pure
• If half cell does not include a solid metal, platinum is used
Connecting wire
Half-cells measured under these
conditions are known as
standard half-cells
Metal Electrode
Aqueous solution of
metal ions:
[M+] = 1.0 mol dm-3
Combining half cells 2
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Measuring Standard Electrode Potentials
• When the standard hydrogen electrode is connected to another standard hall-cell, the EMF generated is
known as the Standard Electrode Potential.
• Given the symbol Eө
+ve value for Eө
(+0.34V) indicates
that this has greater
tendency to be
reduced than H+
Electrons flow from
hydrogen half-cell
(oxidized) to copper
half cell (reduced)
Anode
Overall reaction equation:
Cathode
More reactive metals tend to lose their
electrons…
• H+ will be
reduced
• Electrons flow
toward Hydrogen
• -ve value for Eө
Overall reaction equation:
Anode
OXIDISED
From the examples we can see: Zinc has a lower Eө than hydrogen so can reduce H+…
From the examples we can see: Copper has a higher Eө than hydrogen so cannot reduce H+…
Cathode
REDUCED
Remember, Standard Electrode Potentials are give for
REDUCTION reaction
• Sometimes known as the Standard Reduction Potentials
• Oxidised species on the left, reduced species on the right.
Note:
• The Eө values do not depend on the total number of electrons, no need to
scale according to stoichiometry
The electrochemical series
The electrochemical series is a list of standard electrode
potentials (Eө). The equilibria are written with the electrons on
the left of the arrow, i.e. as a reduction.
Half cell
Half equation
Eө / V
Mg2+(aq) / Mg(s)
Mg2+(aq) + 2e-
Mg(s)
-2.36
Zn2+(aq) / Zn(s)
Zn2+(aq) + 2e-
Zn(s)
-0.76
2H+(aq) / H2(g)
2H+(aq) + 2e-
H2(g)
0
Cu2+(aq) / Cu(s)
Cu2+(aq) + 2e-
Ag+(aq) / Ag(s)
Ag+(aq) + e-
Cu(s)
Ag(s)
+0.34
+0.80
Electrodes with negative values of Eө are better at releasing
electrons (i.e. better reducing agents) than hydrogen.
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Calculating Ecell
The e.m.f of an electrochemical cell, Ecell, is the difference
between the standard electrode potentials of the two half cells.
Ecell = Eө (positive electrode) – Eө (negative electrode)
The positive electrode is taken to be the least negative half
cell, and the negative electrode is the most negative half cell.
This can be worked out from the electrode potentials values
in the electrochemical series.
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Calculating Ecell: worked example
An electrochemical cell is set up using the two half reactions
below. What potential difference Ecell would this cell generate?
Zn2+(aq) + 2e-
Zn(s)
Eө = -0.76 V
Cu2+(aq) + 2e-
Cu(s)
Eө = +0.34 V
Ecell = Eө (positive electrode) – Eө (negative electrode)
The zinc half cell has the more negative potential and so
forms the negative electrode. Therefore:
Ecell = (+0.34) – (-0.76)
= +1.10 V
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Calculating Ecell: combining half equations
To find the overall reaction occurring in the cell as a whole,
the two half equations are added together:
Cu2+(aq) + 2e-
Cu(s)
Because the zinc half cell forms the negative electrode of the
cell, oxidation occurs at this electrode and the half equation
must be reversed:
Zn(s)
Zn2+(aq) + 2e-
The two half equations are added to give the overall
cell reaction:
Zn(s) + Cu2+(aq)
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Zn2+(aq) + Cu(s)
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Calculating Ecell
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Remember, Standard Electrode Potentials are give for REDUCTION reaction
• Sometimes known as the Standard Reduction Potentials
• Oxidised species on the left, reduced species on the right.
• The more positive the Eө value, the more readily it is reduced
Note:
• The Eө values do not depend on the total number of electrons, no need to scale according to
stoichiometry
• A selection of electrode potential values are given in section 24 of the IB data booklet
Using Standard Electrode Potential Data
• Because Standard Electrode Potential data are all referenced to the
same point, it provides us a relative scale for us to make predictions
about redox reactions and the direction of electron flow.
There are three specific applications of using Eө:
1. Calculating the cell potential, Eө cell
2. Determining spontaneity of a reaction
3. Comparing relative oxidising and reducing power of half-cells
1. Calculating the cell potential, Eө cell
• From the Eө values for any two half-cells, we can calculate the EMF for the voltaic cell
• The higher Eө value will be REDUCED, the lower Eө value will be OXIDISED
Note:
• The Eө values must be the reduction potentials as supplied in the data tables
• The Eө values do not need to be scaled according to stoichiometry
2 Determining spontaneity of a reaction
A voltaic cell will always run in the direction that gives a positive value for the Eө values
• If a Eө cell is positive  reaction is spontaneous
• If a Eө cell is negative  reaction is non spontaneous (but spontaneous if the reaction is
reversed)
Electrode potential and free energy change (Eө cell and ΔG)
• We have seen two quantitative measures of the spontaneity of a reaction, Eө cell and ΔG (free energy
change – chapter 5)
• These values can be directly related through this equation:
Where:
• N = number of moles of electrons transferred in the reaction;
• F = the charge carried by 1 mole of electrons (Faraday’s constant = 96500 C mol-1)
• When Eө cell is positive, ΔG is negative  spontaneous reaction
• When Eө cell is negative, ΔG is positive  non-spontaneous reaction
• When Eө cell is 0 (zero), ΔG is 0 (zero)  reaction is at equilibrium
The more positive the value
of Eө cell, the more
energetically favourable.
A voltmeter can be an indirect
measure of free energy
change as well as of electrode
potential.
3 Comparing relative oxidizing and reducing power of
half-cells
• Calculations using Eө for half
reactions confirms the order of
the activity series.
• Trends are summarised here:
Final note about Eө data:
Feasibility and rate are two
different considerations.
A spontaneous reaction may
give no observable sign of
reaction because the activation
energy may be too high for the
reaction to occur at an
observable rate.
Solutions