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reduction potentials
reduction potentials
reduction potential or electrode potential?
reduction potential or electrode potential?
The standard reduction potential is the tendency for a chemical species to be reduced, and is measured in volts at standard conditions. The more positive the potential is the more likely it will be reduced.
The standard reduction potential is the tendency for a chemical species to be reduced, and is measured in volts at standard conditions. The more positive the potential is the more likely it will be reduced.
The standard electrode potential is a measure of the energy per unit charge which is available from the oxidation/reduction reactions to drive the reaction.
The standard electrode potential is a measure of the energy per unit charge which is available from the oxidation/reduction reactions to drive the reaction.
In an electrochemical cell, an electric potential is created between two dissimilar metals, acting as two electrodes. Symbol, Eo .
In an electrochemical cell, an electric potential is created between two dissimilar metals, acting as two electrodes. Symbol, Eo .
From this we can calculate the electromotive force (emf) of the cell (see further down this page).
This table of standard reduction potentials, above, tells the relative strengths of oxidants and reductants:
This table of standard reduction potentials, above, tells the relative strengths of oxidants and reductants:
The higher the value of Eo the more likely it is to be reduced, ie it is a stronger oxidant, conversely the lower the value of Eo the more likely it is that the species will be oxidised; ie it will act as a strong reductant
The higher the value of Eo the more likely it is to be reduced, ie it is a stronger oxidant, conversely the lower the value of Eo the more likely it is that the species will be oxidised; ie it will act as a strong reductant
Lets look at a couple of examples:
Lets look at a couple of examples:
1) If we look at the F2/F- pair the value of Eo = +2.87V. This is the highest value of Eo meaning that it will act as a strong oxidant as F2 is going to be reduced to F- very easily; however F- will not readily be oxidised to F2
1) If we look at the F2/F- pair the value of Eo = +2.87V. This is the highest value of Eo meaning that it will act as a strong oxidant as F2 is going to be reduced to F- very easily; however F- will not readily be oxidised to F2
2) If we look at the Li+/Li pair the Eo value is -3.05V; the lowest reduction potential. This means the oxidised version (Li+) will not be readily reduced to Li metal. However the reduced species (Li) will readily be oxidised to Li+, meaning that Li is a strong reductant
2) If we look at the Li+/Li pair the Eo value is -3.05V; the lowest reduction potential. This means the oxidised version (Li+) will not be readily reduced to Li metal. However the reduced species (Li) will readily be oxidised to Li+, meaning that Li is a strong reductant
key points
key points
The species are grouped in pairs: the reductant and conjugate oxidant.
The species are grouped in pairs: the reductant and conjugate oxidant.
Oxidised species is on the left; reduced species is on the right.
Oxidised species is on the left; reduced species is on the right.
The Eo value is measured against the H+/H2 cell which is given an arbitrary value of 0V. A more positive value,ie >0 indicates that the species gains e-1 more easily than H2, and is a stronger oxidant (it becomes reduced). A more negative value would indicate the species will release e-1 more readily, thereby being a stronger reductant.
The Eo value is measured against the H+/H2 cell which is given an arbitrary value of 0V. A more positive value,ie >0 indicates that the species gains e-1 more easily than H2, and is a stronger oxidant (it becomes reduced). A more negative value would indicate the species will release e-1 more readily, thereby being a stronger reductant.
How can we use these values to predict whether a reaction will occur spontaneously? By working out the EMF of the cell (from the electrode potential).
How can we use these values to predict whether a reaction will occur spontaneously? By working out the EMF of the cell (from the electrode potential).
As an example lets see if Cu metal will react with K2SO4 in a redox reaction....( we already know it won't from our previous knowledge of the activity series) but bear with me and lets look at the Eo values and what they tell us, as the half reactions are at equilibrium and when we connect the half reactions (half cells) together equilibrium is disturbed and Le Chatelier's principles apply:
As an example lets see if Cu metal will react with K2SO4 in a redox reaction....( we already know it won't from our previous knowledge of the activity series) but bear with me and lets look at the Eo values and what they tell us, as the half reactions are at equilibrium and when we connect the half reactions (half cells) together equilibrium is disturbed and Le Chatelier's principles apply:
Cu2+ + 2e- ⇋ Cu Eo = +0.34 V
Cu2+ + 2e- ⇋ Cu Eo = +0.34 V
K+ + e- ⇋ K Eo = -2.92V
K+ + e- ⇋ K Eo = -2.92V
In equilibria the more negative the Eo value the further the equilibrium lies to the left (relative to H+/H2 )
In equilibria the more negative the Eo value the further the equilibrium lies to the left (relative to H+/H2 )
In equilibria the more positive the Eo the further equilibria lies to the right
In equilibria the more positive the Eo the further equilibria lies to the right
If we 'connect' the 2 half cells and electrons transfer, in effect we are creating 2 one way reactions. This means that the more negative Eo will move further to the left. It is already to the left and the species present is K+; hence K cannot be produced. Conversely the Cu2+ + 2e- ⇋ Cu half reaction is already to the right and will tend to shift further right, ie the species will be Cu
If we 'connect' the 2 half cells and electrons transfer, in effect we are creating 2 one way reactions. This means that the more negative Eo will move further to the left. It is already to the left and the species present is K+; hence K cannot be produced. Conversely the Cu2+ + 2e- ⇋ Cu half reaction is already to the right and will tend to shift further right, ie the species will be Cu
The key is to look at what we started with, in this instance Cu and K+
The key is to look at what we started with, in this instance Cu and K+
Therefore this reaction cannot happen spontaneously
Therefore this reaction cannot happen spontaneously
We can also show this by using standard electrode (cell) potential calculations, using:
We can also show this by using standard electrode (cell) potential calculations, using:
Eo = Rcell - Lcell , which is
Eo = Rcell - Lcell , which is
Eo = Reduction cell - Oxidation cell
Eo = Reduction cell - Oxidation cell
REMINDER:
ERL
ERL
Eo = Right - Left
Eo = Right - Left
Before writing the numbers down think what the question is asking: can Cu displace K+, is Cu going to oxidise to Cu2+ and will K+reduce to K!:
Before writing the numbers down think what the question is asking: can Cu displace K+, is Cu going to oxidise to Cu2+ and will K+reduce to K!:
Eo = Right Hand Cell = Left Hand Cell
Eo = Right Hand Cell = Left Hand Cell
Eo = Reduction cell - Oxidation Cell
Eo = Reduction cell - Oxidation Cell
Eo = -2.92 - (+0.34) = -3.26V
Eo = -2.92 - (+0.34) = -3.26V
A negative value means it cannot happen spontaneously
A negative value means it cannot happen spontaneously
Now lets look at another reaction: Will Mg react with a dilute acid? Again we already know the answer but lets look at it using standard reduction potentials:
Now lets look at another reaction: Will Mg react with a dilute acid? Again we already know the answer but lets look at it using standard reduction potentials:
Using ERL we get:
Using ERL we get:
E = R cell - Left Cell
E = R cell - Left Cell
Eo = Red - Oxidation
Eo = Red - Oxidation
E0 = 0 - (-2.37) = +2.37V
E0 = 0 - (-2.37) = +2.37V
The + value shows this could occur spontaneously
The + value shows this could occur spontaneously
With ERL the right hand cell is the reduction cell and the left hand cell is the oxidation cell.
With ERL the right hand cell is the reduction cell and the left hand cell is the oxidation cell.
This is important to remember when writing cell notations, which are written as follows:
This is important to remember when writing cell notations, which are written as follows:
An explanation of standard reduction potentials referenced from the Standards hydrogen electrode (SHE)
An explanation of standard reduction potentials referenced from the Standards hydrogen electrode (SHE)
Remember: all Eo values are referenced to the SHE (standard hydrogen electrode): 1 molL-1 H+ , 1 atm, 25 Celsius
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