1. Engineering Chemistry CHM 406
Electrochemistry
Engineering Chemistry
CHM 406

3. Electrochemical cell Invented in early 1800’s (Volta)
Daniell Cell shown (1830’s)
Two half cells: Zn/Zn2+ and Cu/Cu2+
“Salt bridge” connecting them internally
External circuit
Current flows in the external circuit: electrons
Current is carried internally (aqueous solution) by ions: Zn2+, Cu2+, K+, Cl-.

4. Half cell reactions Anode: loss of electrons by zinc – oxidation.
Zn (s) → Zn2+ (aq) e-
Cathode: gain of electrons by copper – reduction.
Cu2+ (aq) e- → Cu (s)
Overall reaction (redox reaction):
Zn (s) + Cu2+ (aq) → Zn2+ (aq) + Cu (s)

5. Electrolytes
When an electric current is passed through an aqueous solution, the current is carried by ions in the solution.
Any substance which dissociates into ions when dissolved in water is called an electrolyte. A solution of an electrolyte conducts electricity.
Pure water itself is a very weak electrolyte – contains traces of H3O+ and OH- ions.

6. Types of electrolytes
Salts: ionic compounds whose ions dissolve in water.
Acids: covalent compounds which ionise and dissociate when dissolved in water, giving H3O+ (aqueous H+) ions.
Bases: molecules or ions which react with acids, removing H+ ions.

7. Acids Strong acids: completely dissociated.
Sulphuric acid: H2SO4(aq) → H+(aq) + HSO4-(aq)
Nitric acid: HNO3(aq) → H+(aq) + NO3-(aq)
Hydrochloric acid: HCl(aq) → H+(aq) + Cl-(aq)
Weak acids are partially dissociated.
Phosphoric acid: H3PO4(aq) H+(aq) + H2PO42-(aq)
Bisulphate: HSO4-(aq) H+(aq) + SO42-(aq)

8. Bases
React with acids, removing H+ ions, or react directly with H+ (aq).
Hydroxide: OH-(aq) + H+(aq) → H2O (l)
Ammonia: NH3 (aq) + H+(aq) → NH4+ (aq)
Carbonate: CO32- (aq) + 2 H+(aq) → CO2(g) + H2O(l)
Compounds that dissociate into OH- ions in water, e.g., NaOH, KOH, Ba(OH)2, are strong bases.

9. Electrolysis
When an electric current is passed through a solution of an electrolyte, or a molten salt, electrolysis occurs.
Different products are formed at each electrode.
Reverse of an electrochemical cell; electri- city is used to force a non-spontaneous or unfavoured reaction to occur.
Can be used to produce substances otherwise difficult to obtain, e.g., Na, Mg, Al.

10. Electrolysis of water Usually done using a dilute solution of H2SO4.
Inert electrodes – platinum, graphite, etc.
Cathode: reduction of H+.
2H+(aq) + 2e- → H2 (g)
Anode: H2O oxidised rather than SO42-.
2 H2O (l) → 4 H+(aq) + O2 (g) + 4e-
Overall:
2 H2O → 2 H2 + O2

11. Electrolysis of brine
Brine is concentrated aqueous NaCl, e.g., salt water.
Cathode: reduction of H2O.
2 H2O (aq) + 2e- → H2 (g) + 2 OH- (aq)
Anode: Cl- oxidised rather than H2O.
2 Cl- (aq) → Cl2 (g) + 2e-
Overall:
2 H2O + 2 Cl- → H2 + Cl OH-
H2, Cl2, and aqueous NaOH are produced.

13. Electrolysis of molten salts and oxides
No water – no H2 formed.
Used to manufacture metals which cannot be obtained by chemical means.
Highly reactive, e.g., Na (from NaCl), K (from KCl), etc.
Form very stable oxides, e.g., Al (from Al2O3), Ti (from TiO2) etc.

14. Electroplating
Electrochemistry is used to deposit a layer of a metal on an object, usually also made of metal.
The object to be plated must be de-greased and thoroughly cleaned, then immersed in a bath containing ions of the metal to be plated, e.g., Cu2+, Ag+, Cr3+, etc.
It is then connected to an external DC source, such that it is the cathode.

16. Electroplating (contd.)
The anode is usually a piece of the metal to be plated, though it can be inert.
When the current is passed, the anode gradually dissolves, while the cathode (object to be plated) acquires a coating of metal.
E.g., silver (Ag) plating:
Anode: Ag (s) → Ag+ (aq) + e-
Cathode: Ag+ (aq) + e- → Ag (s)

17. Applications of electroplating
Aesthetic or decorative, e.g., gold or silver.
Strengthening or protection against corrosion, e.g., chromium, copper, zinc.
Purification of impure metal: used as anode, with a piece of pure metal as cathode. The impure anode dissolves, and pure metal is deposited on the cathode.

18. Redox reactions
Reactions in which electrons are transferred from a reducing agent (reductant) to an oxidising agent (oxidant).
Combustion of fuels
Corrosion of metals
Conversion of metal ores to metals.
Examples of oxidants: O2, Cl2, metal cations, metal oxides and oxy-anions (e.g., CrO42-).
Examples of reductants: H2, compounds rich in H and C, pure metals and cations with low charge, anions such as Cl-, Br-, S2-, etc.

19. Half reactions
All redox reactions can be written as the sum of two half reactions (whether or not they are compartmentalised as in a cell).
Oxidation half reaction.
Reduction half reaction.
Example:
Oxidation: Zn (s) → Zn2+ (aq) e-
Reduction: Cu2+ (aq) + 2 e- → Cu (s)
Overall: Zn + Cu2+ → Zn Cu

20. Common oxidation half-reactions
M (any metal) → Mn+ + n e-
H2 → 2 H e-
Fe2+ → Fe e-
2 Cl- → Cl e-
2 S2O32- → S4O e-
2 H2O → 4 H+ + O e-

21. Common reduction half-reactions
Mn+ + n e- → M
O e H+ → 2 H2O
F e- → 2 F-
Cu e- → Cu+
PbO H+ + SO e- → PbSO H2O
MnO H e- → Mn H2O
CrO H e- → Cr H2O

22. Oxidation states (oxidation numbers)
How can we determine how many electrons are gained or lost? Sometimes not obvious.
The oxidation state of an atom is a number that indicates the degree to which the atom is oxidised (or reduced) compared to its elemental state.
The oxidation state of an elemental atom is zero.
The difference in oxidation states is the number of electrons gained or lost.

23. Calculating oxidation states
The O.S. of a charged atom is equal to its charge.
For covalently bonded atoms, polar covalent bonds are treated like ionic bonds.
F is always -1, except in F2.
O is always -2, except when it is bonded to F or another O.
H bonded to C, N, O, or a halogen (Group 17) is considered to be +1.
The charge of a polyatomic ion is the sum of the O.S.’s of its atoms.
The O.S.’s of any other atom can be calculated from the above.

24. Example: chromate
CrO42- O.S. (Cr) + 4 x O.S.(O) = -2 O.S. (Cr) + 4 x (-2) = -2 O.S. (Cr) = +6 CrO42- (+6) → Cr3+ (+3) : 3 electrons needed. Reaction takes place in acid; balance by adding H+ to left hand side and H2O to right. CrO H+ + 3 e- → Cr H2O

25. Balancing redox equations
Multiply both oxidation and reduction equations (must be balanced!) by numbers such that the number of electrons gained and lost is equal.
If necessary, balance using H2O and either H+ ions or OH- ions.
Add the equations, so that the electrons cancel out. Overall equation will be balanced.

26. Example MnO4- + Fe2+ → Mn2+ + Fe3+ (in acidic medium)
Oxidation: Fe2+ → Fe e-
Reduction: MnO4- (+7) → Mn2+ (+2)
(5 electrons needed)
MnO H e- → Mn H2O
5 Fe2+ → 5 Fe e-
Overall: MnO H Fe2+ →
Mn H2O Fe3+

27. Electromotive force / cell potential
For a current to flow in the external circuit between two electrodes, there must be an electromotive force. This emf is known as the cell potential, E.
Expressed in volts (V).
Measured using a potentiometer or an ultra-low current voltmeter.
The cell potential is the potential difference between the two electrodes.

28. Electrode potentials
We can assign to every electrode (half cell) a number in volts, called the electrode potential, such that the difference between two such numbers for two electrodes in a cell, is the cell potential.
The electrode potential is a measure of the tendency of the half reaction at that electrode (oxidation or reduction) to take place.
By convention, electrode potentials are tabulated as reduction potentials.

29. Measuring electrode potentials
Absolute value cannot be measured.
Values are measured relative to a standard electrode.
The Standard Hydrogen Electrode (S.H.E.) is used as the reference electrode for this purpose.
By definition, the electrode potential for the S.H.E. under standard conditions is zero.

30. Standard Hydrogen Electrode

31. Components of S.H.E.
Inert electrode (Pt foil – H2 easily adsorbed on the surface).
Immersed in dilute H2SO4, such that H+ concentration is 1 mol L-1.
H2 gas bubbled into the solution past the electrode, such that the H2 pressure is 1 bar (105 Pa).
Temperature is 25oC (298 K)
Salt bridge connects the half cell to the other half cell, whose potential is to be determined.
2 H+ (aq) e H2 (g) ; Eo = 0 V

32. Standard Conditions Temperature = 25oC (298 K)
Partial pressure of every gaseous reactant or product = 1 bar (105 Pa)
Concentration of every aqueous reactant or product = 1 mol L-1
Reduction potentials measured under these conditions, relative to the S.H.E. are called “standard reduction potentials,” Eo.

33. Standard Reduction Potentials
Li+ (aq) + e- Li (s) Eo = V Mg2+ (aq) + 2e- Mg (s) Eo = V Al3+ (aq) + 3e- Al (s) Eo = V Zn2+ (aq) + 2e- Zn (s) Eo = V Cr3+ (aq) + 3e- Cr (s) Eo = V Fe2+ (aq) + 2e- Fe (s) Eo = V Ni2+ (aq) + 2e- Ni (s) Eo = V 2 H+ (aq) + 2e- H2 (g) Eo = 0.00 V Cu2+ (aq) + 2e- Cu (s) Eo = 0.34 V I2 (s) + 2e- 2 I- (aq) Eo = 0.54 V Ag+ (aq) + e- Ag (s) Eo = 0.80 V O2 (g) + 4H+ + 4e- 2 H2O (l) Eo = 1.23 V F2 (s) + 2e- 2 F- (aq) Eo = 2.87 V

34. Electrochemical series
Lists half reactions in order of reactivity.
More positive Eo  greater tendency of the reaction to occur.
For oxidation potentials – reverse the reaction & change the sign of Eo.
Half reactions at the top of the series are more likely to occur in reverse: best reducing agent is Li.
Half reactions at the bottom are more likely to occur as shown: best oxidising agent is F2.

35. Types of electrodes
Reactive metals in contact with solutions of their cations.
E.g., Zn(s) | aq. ZnSO4
Reactive gases bubbling through solutions of cations or anions, in contact with an inert electrode.
E.g., Pt(s) | Cl2(g) | aq. KCl
Inert electrodes (Pt) in contact with aqueous solutions of the reactants & products of the half reaction
E.g., Pt(s) | Fe2+(aq), Fe3+(aq)

36. Types of electrodes (contd.)
Reactive metals coated with an insoluble salt of that metal, the salt in contact with a solution of its anions.
E.g., Ag(s) | AgCl(s) | aq. KCl
AgCl (s) + e Ag (s) + Cl- (aq)
Membrane electrodes.
E.g., glass electrode, used to measure concentration of H+ ions in solution (pH).

37. The calomel electrode

38. The calomel electrode (contd.)
The S.H.E. is not a convenient electrode for regular use as a reference.
A reference electrode needs to be easy to use, stable, reproducible, and reliable.
The calomel electrode is the electrode of choice: Pt (s) | Hg (l) | Hg2Cl2 (s) | saturated aq. KCl.
Hg2Cl2(s) + 2e Hg(l) Cl-(aq)
E = V (relative to S.H.E.)

39. Standard Cell Potentials
In principle, any two half-cells can be combined to give a cell.
The cell potential Ecell is the difference between two reduction potentials, or equivalently, the sum of a reduction and an oxidation potential.
If determined under standard conditions, it is a standard cell potential, Ecello.
Ecell positive  spontaneous reaction.

40. Electrochemical work
When an electric current is passed between 2 electrodes, the electrical work done is
w = - q x Ecell
q = charge (coulombs; q = i (in A) x t (in s) ).
For a redox reaction, q can be determined in terms of moles of electrons transferred.
q = n F
n = number of moles of electrons
F = charge per mole of electrons

41. Electrochemical work (contd.)
F is known as Faraday’s constant.
F = 96,480 C mol-1 (J V-1 mol-1).
The work that can be done is referred to as free energy, DG.
Therefore
DG = - n F Ecell

42. The Nernst Equation
The Nernst equation is used to determine Ecell when conditions are not standard, i.e., concentrations ≠ 1 mol L-1.
Q = quotient
R = gas constant = J K-1 mol-1
F = Faraday’s constant
n = number of electrons transferred.

43. What is the “quotient?”
Consider any reaction (redox or otherwise), which may be denoted aA + bB + cC + … xX + yY + zZ + … where A, B, C, etc. are the reactants / products, and a, b, c, etc. are the coefficients in the balanced equation. Let [A], [B], [C], etc., be the concentrations of A, B, C, etc., in mol L-1. Then Thus, if the various concentrations are known, Ecell can be calculated.