Standard Cell Potentials (HL) (HL IB Chemistry)

• 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= E_right - E_left

• Sometimes this can be hard to remember, but it helps if you remember the phrase 'knives & forks'

Trick to remember how to calculate EMF 

You hold your knife in your right hand and your fork in your left hand. EMF is right minus left

• If the standard hydrogen electrode is placed on the left-hand side of the voltmeter, then by convention E_left will be zero and 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 = E_RHS - E_LHS  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 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 E^θ 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 = E_RHS - E_LHS

+0.46 = E^θ_Ag - (+0.34 V)

E^θ_Ag = 0.46 + 0.34 = +0.80 V

Calculating a cell EMF

• If both SEP's 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 = E_RHS - E_LHS

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 = E_RHS – E_LHS

• A positive EMF is obtained from the spontaneous reaction which occurs when the most negative half cell is E_LHS and the most positive is E_RHS
• The left side is always where oxidation takes place so we can also us an alternative form of the relationship:

EMF = E_reduction – E_oxidation

• The E^ꝋ 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 is a list of various redox equilibria in order of decreasing E^ꝋ values
• More positive (less negative) E^ꝋ values indicate that: 
  • The species is easily reduced
  • The species is a better oxidising agent
• Less positive (more negative) E^ꝋ values indicate that:
  • The species is easily oxidised
  • The species is a better reducing agent

Diagram to show the trends in oxidising and reducing power

Metals with the most negative E^θ values are the strongest reducing agents and non-metals with the more positive E^θ values are the strongest oxidising agents