Standard Cell Potentials (HL) (DP IB Chemistry): Revision Note
Standard cell potentials
Voltmeters measure the potential on the right-hand side of the cell and subtract it from the potential on the left-hand side of the cell
EMF= Eright - Eleft
Sometimes this can be hard to remember, but it helps if you remember the phrase 'knives & forks'
Trick to remember how to calculate EMF
If the standard hydrogen electrode is placed on the left-hand side of the voltmeter, then by convention:
Eleft will be zero
The EMF of the cell will be the electrode potential of the right-hand electrode
For example, if the standard zinc electrode is connected to the standard hydrogen electrode and the standard hydrogen electrode is placed on the left, the voltmeter measures -0.76V
Zn2+(aq) + 2e- ⇌ Zn(s)
The Zn2+(aq) + 2e- ⇌ Zn(s) half-cell thus has an electrode potential of -0.76V
If the Cu2+(aq) + 2e- ⇌ Cu(s) electrode is connected to the standard hydrogen electrode and the standard hydrogen electrode is placed on the left, the voltmeter reads +0.34V
The Cu2+(aq) + 2e- ⇌ Cu(s) half-cell thus has an electrode potential of +0.34V
Standard electrode potential
The standard electrode potential of a half-reaction is the emf of a cell where the left-hand electrode is the standard hydrogen electrode and the right-hand electrode is the standard electrode in question
The equation EMF = ERHS - ELHS can be applied to electrochemical cells in two ways:
Calculating an unknown standard electrode potential
Calculating a cell EMF
To be a standard electrode potential the measurements must be made at standard conditions, namely:
1.0 mol dm-3 ions concentrations
100 kPa pressure
298.15 K
Calculating an unknown standard electrode potential
If the RHS and LHS electrode are specified, and the EMF of the cell measured accordingly, then if the Eθ of one electrode is known then the other can be deduced.
For example, if the standard copper electrode (+0.34 V) is placed on the left, and the standard silver electrode is placed on the right, the EMF of the cell is +0.46 V.
Calculate the standard electrode potential at the silver electrode.
EMF = ERHS - ELHS
+0.46 = EθAg - (+0.34 V)
EθAg = 0.46 + 0.34 = +0.80 V
Calculating a cell EMF
If both standard electrode potentials are known, the EMF of the cell formed can be calculated if the right-hand electrode and left-hand electrode are specified
For example, if in a cell the RHS = silver electrode (+0.80V) and LHS is copper electrode (+0.34 V), then:
EMF = ERHS - ELHS
EMF = +0.80 - 0.34 = +0.46 V
Determining the direction of spontaneity
To predict the spontaneous reaction, we simply need to find the relevant half equations and electrode potentials
From this information, we can deduce the spontaneous and non-spontaneous reaction
By using the convention:
EMF = ERHS – ELHS
A positive EMF is obtained from the spontaneous reaction which occurs when the most negative half cell is ELHS and the most positive is ERHS
The left side is always where oxidation takes place so we can also us an alternative form of the relationship:
EMF = Ereduction – Eoxidation
Worked Example
Using data from the IB Chemistry data booklet (Section 19), determine if the reaction shown is spontaneous at standard conditions:
Sn (s) + Mn2+ (aq) → Sn2+ (aq) + Mn (s)
Answer:
Section 19 of the IB Chemistry data booklet shows the following half-reactions:
Sn2+ (aq) + 2e- → Sn (s) Eθ = -0.14 V
Mn2+ (aq) + 2e- → Mn (s) Eθ = -1.18 V
Manganese is the more negative value, so will be ELHS or Eoxidation in the spontaneous reaction
EMF =ERHS – ELHS
EMF = (-0.14) - (-1.18) = +1.04 V
For oxidation to take place, the manganese must lose electrons and the tin(II) must gain electrons:
Mn (s) → Mn2+ (aq) + 2e-
Sn2+ (aq) + 2e- → Sn (s)
So, the spontaneous reaction is:
Mn (s) + Sn2+ (aq) → Mn2+ (aq) + Sn (s)
Therefore, the reaction in the question is not spontaneous
The Eꝋ values of a species indicate how easily they can get oxidised or reduced
In other words, they indicate the relative reactivity of elements, compounds and ions as oxidising agents or reducing agents
The electrochemical series lists half-reactions in order of standard electrode potentials, from most positive to most negative
A more positive Eꝋ value means:
The species is more easily reduced
It is a stronger oxidising agent
A more negative Eꝋ value means:
The species is more easily oxidised
It is a stronger reducing agent
These trends are also described as:
“Less negative” = closer to zero = more easily reduced
“Less positive” = closer to zero = more easily oxidised
Be careful: “more” and “less” can refer to direction from zero, so always look at the numerical value and the context
Diagram to show the trends in oxidising and reducing power
Examiner Tips and Tricks
A word of caution
Although the positive Eθ indicates a reaction should take place, you might not actually see anything taking place if you constructed a cell that is predicted to be spontaneous
This is because like free energy changes, Eθ only predicts the energetic feasibility of a reaction and it does not take into account the rate of a reaction
A reaction could have a really high activation energy making it impossibly slow at room temperature
'THERMODYNAMICS PREDICTS; KINETICS CONTROLS'
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