1. Electrochemistry

2. Outline (2 lectures, 1 tutorial)
Reactions at electrodes
Half equations
Electrode Potentials
Electrolytic and Galvanic cells
Cell emf
The Standard Hydrogen Electrode (SHE)
Standard Electrode Potential, SEP (E0)

3. Outline The Nernst Equation
Relationship between Eo, Gibb’s Free Energy (G) and equilibrium Constants (K).
Measurement of pH.

4. Electrochemistry
Is the interchange between electricity and materials that may involve charge (electron) or ion transfer.

5. Electron Transfer reactions (Redox)
The exchange of charge in reactions is not always easy to see, so the concept of oxidation number was devised.
Example
2 Mg(s) + O2(g) → MgO(s)
Mg is oxidized:
2 Mg(s) → 2 Mg2+(aq) + 4e-
O2 is reduced:
O2(g) + 4e- → 2O2-(aq)
Oxidation : loss of electrons
Reduction: Gain in electrons
OIL RIG
Oxidation Is Loss Reduction Is Gain

6. Electrochemistry Observations:
a spontaneous reaction occurs (ΔG is –ve) & will continue until it reaches equilibrium.
Gradually a grayish white deposit forms on the copper
The solution becomes pale blue as hydrated Cu2+ ions enter the solution.
Cu0(s) → Cu2+(aq) and Ag+(aq) → Ag0(s)
Overall reaction : Cu(s) + Ag+(aq) → Cu2+(aq) + Ag(s)
Redox Reaction

7. Redox Reaction (Electron Transfer)
Half equations
Oxidation
Cu0(s) → Cu2+(aq)+ 2e-
Reduction
Ag+(aq) + e- → Ag0(s)

8. Electrochemistry
It would be useful to be able to convert this chemical energy to electrical energy instead of heat energy.
This is done by an electrochemical cell.
An Electrochemical Cell - is a device used to convert chemical energy (produced in a redox reaction) into electrical energy.

9. Galvanic or Voltaic Cell
Electrochemical Cells are also known as Galvanic or Voltaic Cells.
Named after Luigi Galvani( ) an Italian anatomist who discovered that electricity can cause the contraction of muscles.
Also called voltaic cells after another Italian scientist Alessandro Volta ( ), the inventor of the battery.
Galvanic cell is one in which the energy released in a spontaneous redox reaction is used to perform electrical work (drive e-s through an external circuit).

10. Example of an Early Galvanic Cell – Daniell Cell (Porous Pot).
Example of an Early Galvanic Cell – Daniell Cell (Porous Pot)
Galvanic cells are formed when we separate the sites where oxidation and reduction in a spontaneous reaction occur.
Higher electrode potential
How does the Daniell Cell work?
the copper strip attracts electrons from the zinc strip.
these electrons pass through the wires of our external circuit.
As the copper electrode receives electrons, free positive ions in the solution arrive to equalize the charges.
Positive Cu2+ are attracted to the charged copper electrode where they receive two electrons and become neutral and deposit on the electrode in metallic form.
The use of a porous barrier allows SO42- to move from right to left (to balance e- flow in circuit) but prevents the Cu2+ and Zn2+ from moving between electrodes.
For each copper atom that is deposited on the copper electrode, a zinc atom goes into solution, giving up two electrons to the zinc electrode.
In order for the voltaic cell to continue to produce an external electric current, there must be a movement of the sulfate ions in solution from the right to the left to balance the electron flow in the external circuit. The metal ions themselves must be prevented from moving between the electrodes, so some kind of porous membrane or other mechanism must provide for the selective movement of the negative ions in the electrolyte from the right to the left.

11. Electromotive Force (emf)
The electromotive force is the maximum potential difference between two electrodes of a galvanic or voltaic cell. (the energy converted into electrical energy when unit charge passes through the source)
It measures the pulling tendency of the two electrodes.
That is the tendency to acquire (i.e. gain) or release (loss) electrons.

12. Daniell Cell
The maximum potential between Zn and Cu of the Daniell cell has been measured to be 1.10 V, but this is not observed in practice.
Factors that affect the measured potential in a cell (e.g. Daniell Cell)?
Porosity of the pot
Cleanliness of the electrodes
Electrical Resistance of the Measuring Device
Minimize
Make pot as porous as possible without allowing solutions to mix
Electrodes clean
High Resistance Voltmeter
Internal Resistance {there is build up of charges inside pot }
Liquid Junction Potential (LJP) {difference in ion mobility}
We do not see1.10V b/c of internal resistance ( a small current is pulled thru’ voltmeter; there is a voltage drop across the voltmeter). V= E-ir
E=1.10V and V is what voltmeter measures.

13. Liquid Junction Potential (LJP)
The liquid junction, porous pot, is also a source of “lost” potential
Why is this?
Build-up of charge results from the difference in mobility of the ions as they move across the wall of the porous pot to neutralize the charge.
Potential difference exists between the inner & outer surfaces of the wall of the porous pot .
This potential that results is called a LIQUID JUNCTION POTENTIAL
In any circuit there are junction potentials whenever two dissimilar materials come into contact.

14. How do we overcome the liquid junction potential?
We need to provide a good conducting path
Use a salt bridge with ions of similar mobilities to reduce the effect of liquid junction potentials (i.e. potentials which arise because of the difference in mobilities of the ions).
Once these precautions are taken the emf of the cell depends solely on the concentrations of the solutions & the metals used as the electrodes.

15. Or No electrical current could be produced by the cell
Salt Bridge
For a galvanic cell to work, the solution in both half cells must remain electrically neutral (charge must be balance).
Therefore ions must be permitted to enter and leave the solutions.
Salt Bridge: a tube filled with an electrolyte solution(viscous), such as KNO3 or KCl and fitted with a porous plug at each end or simply a strip of filter paper soaked in saturated KNO3.
Function:
Maintains electrolytic contact (w/o mixing solutions)
Completes the circuit
Reduces liquid junction potential (LJP)
Viscosity prevents mixing of solutions

16. Daniell Cell – Salt Bridge Galvanic or Voltaic Cell
Cu2+ a better oxidizing agent than Zn2+
Cu2+ has a greater tendency to pull electrons than Zn2+ and this difference is what appears as a difference in potential

17. Electrochemical Cell Each half-cell has a characteristic voltage.
Different choices of substances for each half-cell give different potential differences.
Each reaction is undergoing an equilibrium reaction between different oxidation states of the ions—when equilibrium is reached the cell cannot provide further voltage.
The spontaneous reaction that drives the Daniell cell is:
Cu2+(aq) + Zn(s) → Cu(s) + Zn2+(aq)

18. Galvanic or Voltaic Cell
Oxidation
Reduction Ag
Ag+ has a greater tendency to pull electrons than Cu2+
Ag+ is a better oxidizing agent than Cu2+
The spontaneous reaction: Cu + 2Ag+ → Cu2+ + Ag

19. Electrolytic Cell
An electrochemical cells which uses energy from other source (e.g. DC) to induce redox reactions and in which electrical energy converts into chemical energy.

20. Electrolytic vs Galvanic Cell
NOTE: It is the nature of the chemical change and not the electrical charge that determines whether we label an electrode as cathode or anode.
Electrolytic Cell
Galvanic Cell
Cathode is negative (reduction)
Cathode is positive (reduction)
Anode is positive (oxidation)
Anode is negative (oxidation)

21. Electrode (Cell) Potential
What is potential (E)?
Potential is the amount of work required to bring one unit of charge from infinity to a fixed distance in space.
Electrode Potential or Cell Potential (emf or
electromotive force) : is the difference in
potential between two half-reactions or electrodes.
i.e. Ecell = ΔE= Ecathode– Eanode = Eoxidation + Ereduction
Ecell is expressed in Volts.
Potential is a relative quantity and a standard state needs to be defined.

22. Electrode Potential & Half Cells
Each electrode, i.e. the ion & its neutral atom
Contributes a characteristic potential to the overall cell potential
Independent of the other electrode in the pair
Cu | Cu2+ half cell has a characteristic potential
Zn | Zn2+ half cell has a characteristic potential
Ag | Ag+ half cell has a characteristic potential
To assign a potential to each half cell one must assign an electrode as a “standard electrode” & measure each electrode relative to this standard electrode.

23. Standard Hydrogen Electrode (SHE)
The standard to which all electrodes are compared is the Standard Hydrogen Electrode
Its characteristic potential is ZERO at ALL temperatures
Potentials measured against the SHE are called Reduction Potentials and are represented by Eo in Volts
The SHE is represented as:
Pt(s) | H2(g) | H+(aq)
H2+ 2e- → 2H+
E0 = 0.0Volts

24. Standard Electrode Potentials
Standard potentials are measured with the test electrode on the right hand side
The measured potential is +ve if the electrode has a greater tendency to pull electrons than the H2 electrode (SHE) and –ve if it has a lower tendency
Reduction Potentials
Cu2+ + 2e-  Cu Eθ = V
Zn2+ + 2e-  Zn Eθ = V
Ag+ + e-  Ag Eθ = V
Pb2+ + 2e- Pb Eθ = V
Pb4+ + 2e- Pb Eθ = V

25. The standard-state cell potentials for some common half-reactions
More –ve E0
Species at the top of the series are more readily oxidized
Half-Reaction Eored/V
K+ + e- ↔ K
Ba e- ↔ Ba
Ca e- ↔ Ca
Na+ + e- ↔ Na
Mg e- ↔ Mg
H2 + 2 e- ↔ 2 H
Al e- ↔ Al
Mn e- ↔ Mn
Zn e- ↔ Zn
Cr e- ↔ Cr
S e- ↔ S
CO2 + 2 H+ + 2 e- ↔ H2C2O
Cr3+ + e- ↔ Cr
Fe e- ↔ Fe
Co e- ↔ Co
Ni e- ↔ Ni
Sn e- ↔ Sn
Pb e- ↔ Pb
Fe e- ↔ Fe
2H+ + 2 e- ↔H
Increasing reducing strength
Zn2+ is oxidized when paired with the hydrogen electrode
E°Cell
1.Measured against SHE
2.Concentration 1 Molar
3.Pressure 1 atmosphere
4.Temperature 25°C

26. The standard-state cell potentials for some common half-reactions
Copper is reduced when paired with hydrogen
S4O e- ↔ 2S2O
Sn e- ↔ Sn
Cu2+ + e- ↔ Cu
Cu e- ↔ Cu
O2 + 2 H2O + 4 e- ↔ 4 OH
Cu+ + e- ↔ Cu
I e- ↔ 3 I
MnO H2O + 3 e- ↔ MnO2 + 4OH
O2 + 2 H+ + 2 e- ↔ H2O
Fe3+ + e- ↔ Fe
Hg e- ↔ Hg
Ag+ + e- ↔ Ag
Hg e- ↔ Hg
H2O2 + 2 e- ↔ 2 OH
HNO3 + 3 H+ + 3 e- ↔ NO + 2 H2O
Br2(aq) + 2 e- ↔ 2 Br
IO H e- ↔ I2 + 6 H2O
CrO H+ + 3 e- ↔ Cr H2O
Increasing reducing strength
Increasing oxidizing strength
More +ve E0
*The more positive the Eo value for a half-reaction, the greater the tendency for that reaction to occur as written.
More easily reduced

27. Standard Cell Notation for Galvanic Cell
A s a matter of convenience chemists have devised a shorthand way of representing a galvanic cell.
For example the copper-silver cell is represented as follows.
Cu(s) Cu2+(aq) Ag+(aq) Ag(s)
By convention anode half cell is specified on left
Exercise : Write the anode and cathode half reactions for the
following galvanic cell.
Al(s) Al3+(aq) Pb2+(aq) Pb(s)
Salt bridge
Phase boundary
Al(s) Al3+(aq)+ 3e-
Pb2+(aq) + 2e Pb(s)

28. Daniell Cell - Cell Notation
Daniell Cell can be written as
Zn(s) | ZnSO4(aq) || CuSO4(aq) | Cu(s) or
Zn(s) | Zn2+ (aq) || Cu2+ (aq) | Cu(s)

29. Cell Potential - Calculating E0 of Cells
To calculate voltages for any two electrochemical cell we can do the following:
Daniell Cell:
Zn(s) | Zn2+ (aq) || Cu2+ (aq) | Cu(s)
a) Locate the two half cells reactions in the table of standard reduction potentials.
1. Cu2+ (aq) + 2e-  Cu(s) E0 = V
Zn2+ (aq) + 2e-  Zn(s) E0 = V

30. Cell Potential - Calculating E0 of Cells
b) The half reaction that has the higher reduction potential will reduce and can be written as you find it in the table.
Cu2+ + 2e-  Cu Eθ = V
c) The half reaction that has the lower reduction potential must be reversed and written as an oxidation. (The sign of the Eo value of the lower half reaction is changed ).
Reverse Eqn. 2
Zn  Zn2+ + 2e- Eθ = V

31. Cell Potential - Calculating E0 of Cells
The two half reactions are balanced for the number of electrons exchanged but the value of each Eo remains unchanged.
Cu2+ + 2e-  Cu Eθ = V
Zn  Zn2+ + 2e Eθ = V
E0oxi= - E0red

32. Cell Potential - Calculating E0 of Cells
g) The two half reactions are then added together and so are the Eo values. ( This value will always be positive for an electrochemical cell).
Cu2+ + 2e-  Cu E0 = V
Zn  Zn2+ + 2e E0 = V
Overall eqn. Cu2+ + Zn  Zn2+ + Cu
Overall Ecell0= E0oxi + E0red= = 1.10V
 When E0 is +ve the reaction is spontaneous (thermodynamically favourable) in the direction written.

33. Cell Potential -Exercise
Determine the cell potential for a galvanic cell based on the redox reaction.
Cu(s) + Fe3+ (aq) ® Cu2+ (aq) + Fe2+ (aq)
Fe3+ (aq) + e-® Fe2+ (aq) Eº = 0.77 V
Cu2+ (aq)+2e- ® Cu(s) Eº = 0.34 V
2Fe3+ (aq) + 2e-® 2Fe2+ (aq) Eº = 0.77 V
Cu(s) ® Cu2+ (aq)+2e Eº = V
Overall eqn. Cu(s) + 2Fe3+ (aq) ® Cu2+ (aq) + 2Fe2+ (aq)
Ecello = E0xi + Ered = 0.43V

34. Free Energy and Electrode Potentials
The cell potential of a voltaic cell is a measure of the maximum amount of energy per unit charge which is available to do work when charge is transferred through an external circuit.
This maximum work is equal to the change in Gibbs free energy, ΔG, in the reaction.
Maximum work = ΔG = -nFE°cell
n= number of electrons transferred per mole of reactant (after balancing)
F = Faraday’s constant (96485 Cmol-1)
E°cell= standard cell potential

35. Free Energy and Electrode Potentials
Consider the historic Daniell cell in which zinc and copper were used as electrodes.
E°cell = 1.1 Volts
n= 2 electrons are transferred per mole of reactant.
ΔG = -nFE°cell = -2 x 96,485 Cmol-1 x 1.10 JC-1 = -212 kJ
Voltage is defined as the work done per unit charge.
(1V = 1JC-1)
Cathode (Reduction) Half-Rxn
Standard Potential E° (volts)
Zn2+(aq) + 2e- → Zn(s)
-0.76
Cu2+(aq) + 2e- → Cu(s)
+0.34

36. Gibb's Free Energy
The Gibb's free energy DG is the negative value of maximum electric work,
DG = - W       = - q DE
W = maximum electric work is the product of charge q in Coulomb (C), and the potential DE in Volt (= J/C)
A redox reaction equation represents definite amounts of reactants in the formation of also definite amounts of products.
The number (n) of electrons in such a reaction equation, is related to the amount of charge transferred when the reaction is completed.
Since each mole of electron has a charge of C (known as the Faraday's constant, F),
q = n F
and, DG = - n F DE
At standard conditions,
DG° = - n F DE°

37. G0& E0 - Spontaneity G0 = -n F E0
When E0 is +ve, G0is -ve = reaction is spontaneous
When E0 is -ve, G0 is +ve = reaction is not spontaneous

38. Nernst Equation
As a voltaic cell is discharged, its emf falls until E = 0, at which point we say that the cell is dead. 
Studies show that the emf depends on the concentrations of the reactants and products in the cell reaction. 
Increasing the [reactants] will increase the cell emf.
Increasing the [products] will decrease the cell emf. 
The emf generated under nonstandard conditions can be calculated by using an equation first derived by Walther Hermann Nernst ( ).

39. Nernst Equation
The dependence of the cell emf on concentration can be obtained from the dependence of the ΔG on concentration. 
Recall:
ΔGat any stage of rxn  = ΔGo + RT ln Qrxn ….……. ….eqn. 1
where R = Universal gas constant, T = temperature
and for a generalized equation of the form:
cC + bB+… mM+ nN+……
At equilibrium, Ecell = 0 hence DG = 0 and Qrxn corresponds to Keq.

40. Activity- effective concentration of species.
Some of the species that take part in these electrode reactions are pure solid compounds and pure liquid compounds. In dilute aqueous solutions, water can be treated as a pure liquid.
For pure solid compounds or pure liquid compounds, activities are constant and their values are considered to be unity. i.e. a= 1
The activities of gases are usually taken as their partial pressures and the activities (ai) of solutes such as ions are the product of the molar concentration and the activity coefficient of each chemical species :
ai = γ[i]i ≈ [i]
Activity coefficient depends on concentration of ions, charge and diameter of ions.
Activity arises because molecules in a non-ideal gas or solution interact with each other. Activity depends on temperature, pressure, and composition.

41. Nernst Equation Substituting ΔG = -nFEcell into eqn. 1 gives:
-nFEcell = -nFEocell + RT ln Q  
Solving this equation for E gives the Nernst equation:
The equation is usually written in term of common logarithms:

42. Nernst Equation At 298 K the quantity 2.303RT / F = 0.0592 V,
- so at 298 K a simplified form of the Nernst equation is:

43. Non-standard Conditions
The SEP values refer to standard conditions i.e. 1 Molar concentrations at 25ºC and atmospheric pressure.
If these conditions change then so does the electrode potential.
For example, according to standard electrode potentials, MnO2 will not react spontaneously with HCl, however this is the standard preparation of chlorine in the laboratory.
MnO2(s) + 4H+(aq) + 2e- Mn2+ + 2H2O(l) Eº = 1.23 V
Cl2(g) + 2e Cl-(aq) Eº = 1.36V
Predicting spontaneity, E = Ered + Eox = ( ) Eocell = V
Negative value therefore no reaction!!
In the lab preparation the MnO2 is heated with the concentrated HCl - these are not standard conditions, the temperature >>25ºC and [HCl] >>1 mol dm-3.
Under these new conditions the reaction becomes spontaneous and proceeds at a comfortable rate to collect the chlorine gas produced.
MnO2(s) + 4H+(aq) + 2Cl-(aq) Mn2+ (aq) + 2H2O(l) + Cl2(g)

44. Worked Example Calculate the emf of the following cell at 25°C;
Sn(s)|Sn2+(0.025 M)||Ag+(2.0 M)|Ag(s) 
Sn2+(aq) +2e Sn(s) V
Ag+ +e Ag(s) V
2Ag+ + Sn(s) Sn2+ + 2Ag(s)
E°cell = E°oxi + E°red = 0.80 V V = 0.94 V

45. Nernst Equation
If we have a Cu2+/Cu electrode in one half & the SHE in the other
Pt(s) | H2(g) | H+(aq) || Cu2+(aq) | Cu(s) E0 = V
Cu2+ + 2e-  Cu E0 = V
2H+ + 2e-  H2 E0 = V
(E0xi + Ered) Cu2+(aq) + H2(g) Cu(s) + 2H+(aq) E0 = V
Q = [aH+]2 [aCu]/[aH2][aCu2+]
Q = [1.00]2 [1]/[1.0][1.00] = 1

46. Nernst Equation - pH
When H+ concentration is NOT 1.00 M but everything remains the same.
log10Q, for pH measurements, can be expressed in terms of the hydrogen ion activity by :-log10aH+ which is the pH of a solution and the equation becomes:
The measured potential is related to the activity/ concentration of H+ and E0 of the cell
pH can be measured electrically
E.g. pH meter

47. Nernst Equation - pH Take for example:
MnO4- + 5e- + 8H+ ↔ Mn2+ + 4H2O E0 = 1.51V
Here the potential depends not only on the concentrations of the manganese species but also on the pH of the solution. 8

48. Applications Widespread applications of electrochemistry.
In the industry important chemical : liquid bleach(NaOCl) and lye (NaOH) are manufactured by electrochemical reactions.
Batteries (galvanic cells) which produce electrical energy by means of chemical reactions are used to power toys, flashlights, calculators, laptop computers, cellphones, clocks, watches, remote controllers etc. (Dry Cells)
Wet cells : lead acid batteries in some cars.
Fuel Cells e.g. hydrogen fuel cell