1. Thermodynamics at the electrode
Electrochemistry
Thermodynamics at the electrode

2. Learning objectives You will be able to:
Identify main components of an electrochemical cell
Write shorthand description of electrochemical cell
Calculate cell voltage using standard reduction potentials
Apply Nernst equation to determine free energy change
Apply Nernst equation to determine pH
Calculate K from electrode potentials
Calculate amount of material deposited in electrolysis

3. Energy in or energy out
Galvanic (or voltaic) cell relies on spontaneous process to generate a potential capable of performing work – energy out
Electrolytic cell performs chemical reactions through application of a potential – energy in

4. Redox Review Oxidation is... Reduction is...
Loss of electrons
Reduction is...
Gain of electrons
Oxidizing agents oxidize and are reduced
Reducing agents reduce and are oxidized

5. Redox at the heart of the matter
Zn displaces Cu from CuSO4(aq)
In direct contact the enthalpy of reaction is dispersed as heat, and no useful work is done
Redox process:
Zn is the reducing agent
Cu2+ is the oxidizing agent

6. Separating the combatants
Each metal in touch with a solution of its own ions
External circuit carries electrons transferred during the redox process
A “salt bridge” containing neutral ions completes the internal circuit.
With no current flowing, a potential develops – the potential for work
Unlike the reaction in the beaker, the energy released by the reaction in the cell can perform useful work – like lighting a bulb

7. Labelling the parts

8. Odes to a galvanic cell Cathode Anode Where reduction occurs
Where electrons are consumed
Where positive ions migrate to
Has positive sign
Anode
Where oxidation occurs
Where electrons are generated
Where negative ions migrate to
Has negative sign

9. The role of inert electrodes
Not all cells start with elements as the redox agents
Consider the cell
Fe can be the anode but Fe3+ cannot be the cathode.
Use the Fe3+ ions in solution as the “cathode” with an inert metal such as Pt

10. Anode
Cathode
Oxidation
Reduction

11. Cell notation Anode on left, cathode on right
Electrons flow from left to right
Oxidation on left, reduction on right
Single vertical = electrode/electrolyte boundary
Double vertical = salt bridge
Cathode:
Cu2+ + 2e →Cu
Anode:
Zn →Zn2+ + 2e

12. Vertical │denotes different phase
Fe(s)│Fe2+(aq)║Fe3+(aq),Fe2+(aq)│Pt(s)
Cu(s)│Cu2+(aq)║Cl2(g)│Cl-(aq)│C(s)

13. Connections: cell potential and free energy
The cell in open circuit generates an electromotive force (emf) or potential or voltage. This is the potential to perform work
Energy is charge moving under applied voltage

14. Relating free energy and cell potential
The Faraday:
F = C/mol e
Standard conditions (1 M, 1 atm, 25°C)

15. Standard Reduction Potentials
The total cell potential is the sum of the potentials for the two half reactions at each electrode
Ecell = Ecath + Ean
From the cell voltage we cannot determine the values of either – we must know one to get the other
Enter the standard hydrogen electrode (SHE)
All potentials are referenced to the SHE (=0 V)

16. Unpacking the SHE
The SHE consists of a Pt electrode in contact with H2(g) at 1 atm in a solution of 1 M H+(aq).
The voltage of this half-cell is defined to be 0 V
An experimental cell containing the SHE half-cell with other half-cell gives voltages which are the standard potentials for those half-cells
Ecell = 0 + Ehalf-cell

17. Zinc half-cell with SHE
Cell measures 0.76 V
Standard potential for Zn(s) = Zn2+(aq) + 2e = 0.76 V

18. Where there is no SHE
In this cell there is no SHE and the measured voltage is 1.10 V

19. Standard reduction potentials
Any half reaction can be written in two ways:
Oxidation:
M = M+ + e (+V)
Reduction:
M+ + e = M (-V)
Listed potentials are standard reduction potentials

21. Applying standard reduction potentials
Consider the reaction
What is the cell potential?
The half reactions are:
E° = 0.80 V – (-0.76 V) = 1.56 V
NOTE: Although there are 2 moles of Ag reduced for each mole of Zn oxidized, we do not multiply the potential by 2.

22. Extensive v intensive
Free energy is extensive property so need to multiply by no of moles involved
But to convert to E we need to divide by no of electrons involved
E is an intensive property

23. The Nernst equation
Working in nonstandard conditions

24. Electrode potentials and pH
For the cell reaction
The Nernst equation
Half-cell potential is proportional to pH

25. The pH meter is an electrochemical cell
Overall cell potential is proportional to pH
In practice, a hydrogen electrode is impractical

26. Calomel reference electrodes
The potential of the calomel electrode is known vs the SHE. This is used as the reference electrode in the measurement of pH
The other electrode in a pH probe is a glass electrode which has a Ag wire coated with AgCl dipped in HCl(aq). A thin membrane separates the HCl from the test solution

27. Cell potentials and equilibrium
Lest we forget…
So then
and

28. Cell potential a convenient way to measure K

29. Many pathways to one ending
Measurement of K from different experiments
Concentration data
Thermochemical data
Electrochemical data

30. Batteries The most important application of galvanic cells
Several factors influence the choice of materials
Voltage
Weight
Capacity
Current density
Rechargeability

31. Running in reverse
Recharging a battery requires to run the process in reverse by applying a voltage
In principle any reaction can be reversed
In practice it will depend upon many factors
Reversibility depends on kinetics and not thermodynamics
Cell reactions that involve minimal structural rearrangement will be the easiest to reverse

32. Lithium batteries Lightweight (Molar mass Li = 6.94 g) High voltage
Reversible process

33. Fuel cells – a battery with a difference
Reactants are not contained within a sealed container but are supplied from outside sources

34. Store up not treasures on earth where moth and rust…
An electrochemical mechanism for corrosion of iron. The metal and a surface water droplet constitute a tiny galvanic cell in which iron is oxidized to Fe2+ in a region of the surface (anode region) remote from atmospheric O2, and O2 is reduced near the edge of the droplet at another region of the surface (cathode region). Electrons flow from anode to cathode through the metal, while ions flow through the water droplet. Dissolved O2 oxidizes Fe2+ further to Fe3+ before it is deposited as rust (Fe2O3·H2O).

35. Mechanisms Why does salt enhance rusting?
Improves conductivity of electrolyte
Standard reduction potentials indicate which metals will “rust”
Aluminium should corrode readily. It doesn’t. Is thermodynamics wrong?
No, the Al2O3 provides an impenetrable barrier

36. No greater gift than to give up your life for your friend
A layer of zinc protects iron from oxidation, even when the zinc layer becomes scratched. The zinc (anode), iron (cathode), and water droplet (electrolyte) constitute a tiny galvanic cell. Oxygen is reduced at the cathode, and zinc is oxidized at the anode, thus protecting the iron from oxidation.

37. Electrolysis Electrolysis of a molten salt using inert electrodes
Signs of electrodes:
In electrolysis, anode is positive because electrons are removed from it by the battery
In a galvanic cell, the anode is negative because is supplies electrons to the external circuit

38. Electrolysis in aqueous solutions – a choice of process
There are (potentially) competing processes in the electrolysis of an aqueous solution
Cathode
Anode

39. Thermodynamics or kinetics?
On the basis of thermodynamics we choose the processes which are favoured energetically
But…chlorine is evolved at the anode

40. The role of overpotentials
Thermodynamic quantities prevail only at equilibrium – no current flowing
When current flows, kinetic considerations come into play
Overpotential represents the additional voltage that must be applied to drive the process

41. In the NaCl(aq) solution the overpotential for evolution of oxygen is greater than that for chlorine, and so chlorine is evolved preferentially
Overpotential will depend on the electrolyte and electrode. By suitable choices, overpotentials can be minimized but are never eliminated
The limiting process in electrolysis is usually diffusion of the ions in the electrolyte (but not always)
Driving the cell at the least current will give rise to the smallest overpotential

42. Electrolysis of water
In aqueous solutions of most salts or acids or bases the products will be O2 and H2

43. Quantitative aspects of electrolysis
Quantitative analysis uses the current flowing as a measure of the amount of material
Charge = current x time
Moles = charge/Faraday