Predictions from Electrode Potentials

Predicting Feasibility: Can predict feasibility of reaction from standard electrode potentials. A
reaction takes place when the one being reduced (and on left) has a more positive E
o
value than
one being oxidised (and on right).
 This works as higher E
o
value means greater tendency to be reduced so stronger oxidising agent.
 Overall Cell Equation: With both half equations of the half cells, can write an overall redox
equation using methods learnt at start of this topic. But make sure more negative or oxidation E
o
half equation equilibria is switched round, so electrons are on RHS as losing electrons.
 Limitations of Predictions using E
o
value…
– Reaction may have very large activation energy, so slow rate of reaction.
– Or conditions not standard. For example
 Concentration: Standard electrode potentials measured using 1 mol dm-3. If not then electrode
potential, different value.
 Other Factors: Actual conditions may not be standard conditions. Will affect electrode potential
values. When conditions, the temperature, conc or pressure not standard, shifts equilibrium and
movement of electrons and may change cell potential.
 Also standard electrode potential apply to aqueous equilibria but many reactions take place that
are not aqueous.

Why does cell potential increase, in terms of equilibrium when water added- [Cu2+]
concentration decreases, equilibrium shifts to left, more electrons are released by Cu, The cell
has a bigger difference in E so cell potential increases. Look at half equation only not full. Have to
think if cell potential Ecell increases, means E
o
of oxidation half-cell must have increased, value
more negative.
 If concentration of Zn2+ greater than 1 moldm-3, equilibrium shifts right, removing electrons to
form Zn, making electrode potential less negative.
 Cell potential slowly changes over time because concentration of reactants change as reaction
occurs (becomes non- standard).
 Why does pH of solution in hydrogen half cell decrease? H redox system more negative,
releases electrons. Equilibrium shifts to increase H+.
 Why should acidic conditions not be used? H+ reacts with something to form something else.
Equilibrium shifts.
 Mass of electrode changes over course of reaction because metal formed from its aqueous ions
OR aqueous ions are formed from the metal. Include equations.
– In acidic conditions, conc of H+ increases, equilibrium position moves right to reduce
concentration of H+, less electrons so make electrode potential more positive.
– In alkaline conditions, alkali reacts and removes H+, equilibrium position moves left to
increase concentration of H+, electrode potential becomes more negative.