electrochemistry

1. Redox and Electrochemistry
By-
Saurav K. Rawat
(Rawat DA Greatt) 1

2. Redox Reactions and
Electrochemistry
I. Redox Reactions
a) Oxidation Number
b) Oxidizing and Reducing Reagents
II. Galavanic or Voltaic Cells
a) Anode/Cathode/Salt Bridge
b) Cell Notations
c) Determining Cell Potential/Cell Voltage/Electromotive
force (emf)
III. Relating Cell Potential to K and DG0
IV. Effect of Concentration on Cell Potential

3. Redox Reactions and
Electrochemistry
V. Corrosion
VI. Batteries
VII. Fuel Cells
VIII.Electrolytic Cells
a) Calculating amounts of substances reduced or
oxidized

4. Electrochemistry: Interconversion of electrical and
chemical energy using redox reactions
Redox (Oxidation-Reduction) Reaction: Type of electron
transfer reaction. One substance gives up electrons;
the other accepts electrons.
OIL RIG
•Oxidation Half-Reaction; Oxidation Involves Loss of electrons
2Mg 2Mg2+ + 4e-
•Reduction Half-Reaction; Reduction Involves Gain of electrons
O2 + 4e- 2O2-
Net Redox Rxn; 2Mg + O2 -> 2 Mg+2 + 2 O-2

6. Oxidation number
The charge the atom would have in a molecule (or an
ionic compound) if electrons were completely transferred
to the more electronegative atom.
1. Oxidation number equals ionic charge for monoatomic ions
in ionic compound
2. Metal ions in Family A have one, positive oxidation
number; Group IA metals are +1, IIA metals are +2
Li+, Li = +1; Mg+2, Mg = +2
4.4
CaBr2; Ca = +2, Br = -1

7. Oxidation number,continued
The charge the atom would have in a molecule (or an
ionic compound) if electrons were completely transferred
to the more electronegative atom.
3. The oxidation number of a transition metal ion is positive,
but can vary in magnitude.
4. Nonmetals can have a variety of oxidation numbers,both
positive and negative numbers which can vary in
magnitude.
4.4
5. Free elements (uncombined state) have an oxidation number
of zero. Each atom in O2, F2, H2, Cl2, K, Be has the same
oxidation number; zero.

8. 6. The oxidation number of fluorine is always –1.
(unless fluorine is in elemental form, F2)
7. The sum of the oxidation numbers of all the atoms in a
molecule or ion is equal to the charge on the molecule or
ion.
IF; F= -1; I = +1
8. The oxidation number of hydrogen is +1 except when it is
bonded to metals in binary compounds. In these cases, its
oxidation number is –1 or when it’s in elemental form (H2;
oxidation # =0).
HF; F= -1, H= +1
NaH; Na= +1, H = -1

9. 9. The oxidation number of oxygen is usually –2. In H2O2 and
O2
2- it is –1, in elemental form (O2 or O3) it is 0.
H2O ; H=+1, O= -2
SO3; O = -2; S = +6
HCO3
-
O = -2 H = +1
3x(-2) + 1 + ? = -1
C = +4
Oxidation numbers of all
the atoms in HCO- ?
3
4.4

10. NaIO3
Na = +1 O = -2
3x(-2) + 1 + ? = 0
I = +5
IF7
F = -1
7x(-1) + ? = 0
I = +7
K2Cr2O7
O = -2 K = +1
7x(-2) + 2x(+1) + 2x(?) = 0
Cr = +6
Oxidation numbers of all
the elements in the
following ?
4.4

11. Determination of Oxidizing and
Reducing Agents
I. Determine oxidation # for all atoms in
both the reactants and products.
II. Look at same atom in reactants and
products and see if oxidation # increased
or decreased.
• If oxidation # decreased; substance reduced
• If oxidation # increased; substance oxidized

12. Determination of Oxidizing and
Reducing Agents, continued
• Oxidizing Agent: Substance that oxidizes
the other substance by accepting electrons.
It is reduced in reaction.
• Reducing Agent: Substance that reduces the
other substance by donating electrons. It is
oxidized in reaction.

13. Spontaneous Redox Reaction
Zn(s) + Cu+2 (aq) -> Cu(s) + Zn+2(aq)
Zn
Cu+2
time Zn+2
Cu

15. Gets Smaller -> <- Gets Larger

16. Voltaic Cell Animation
Anode; Site of Oxidation
Cathode; Site of Reduction
AnOx or both vowels
Red Cat or both consonants
Direction of electron flow; anode to cathode (alphabetical)
Salt Bridge; Maintains electrical neutrality
+ ion migrates to cathode
- ion migrates to anode

17. Cell Notation
1. Anode
2. Salt Bridge
3. Cathode
Anode | Salt Bridge | Cathode
| : symbol is used whenever there is a different phase

18. 19.2
Cell Notation
Zn (s) + Cu2+ (aq) Cu (s) + Zn2+ (aq)
[Cu2+] = 1 M & [Zn2+] = 1 M
Zn (s) | Zn2+ (1 M) || Cu2+ (1 M) | Cu (s)
anode cathode
More detail..
Zn (s)| Zn+2 (aq, 1M)| K(NO3) (saturated)|Cu+2(aq, 1M)|Cu(s)
anode Salt bridge cathode

19. Zn (s) + 2 H+(aq) -> H2 (g) + Zn+2 (aq)
K(NO3)
Zn(s)| Zn+2|KNO3|H+(aq)|H2(g)|Pt

20. Electrochemical Cells
19.2
The difference in electrical
potential between the anode and
cathode is called:
• cell voltage
• electromotive force (emf)
• cell potential
0 0 0
Cell oxidation reduction E = E + E
UNITS: Volts Volt (V) = Joule (J)
Coulomb, C

25. Standard Electrode Potentials
19.3
Standard reduction potential (E0) is the voltage associated with a
reduction reaction at an electrode when all solutes are 1 M and
all gases are at 1 atm.
2e- + 2H+ (1 M) H2 (1 atm)
E0 = 0 V
Standard hydrogen electrode (SHE)
Reduction Reaction

27. Determining if Redox Reaction is Spontaneous
• + E°CELL ; spontaneous
reaction
• E°CELL = 0; equilibrium
• - E°CELL; nonspontaneous
reaction
More positive E°CELL ;
stronger oxidizing agent or
more likely to be reduced

29. • E0 is for the reaction as written
• The half-cell reactions are
19.3
reversible
• The sign of E0 changes when
the reaction is reversed
• Changing the stoichiometric
coefficients of a half-cell
reaction does not change the
value of E0
• The more positive E0 the
greater the tendency for the
substance to be reduced

30. Relating E0
Cell to DG0
E work Cell arg
ch e
=
Units
work, Joule
charge, Coulomb
Ecell; Volts
charge = nF
Faraday, F; charge on 1 mole e-
F = 96485 C/mole
work = (charge)Ecell = -nFEcell
DG = work (maximum)
DG = -nFEcell

31. Relating Eo
CELL to the
Equilibrium Constant, K
DG0 = -RT ln K
DG0 = -nFE0
cell
-RT ln K = -nFE0
cell
K
E RT Cell 0 = ln
nF
( )
0.0257
J
ö çè
8.31 298
96485
=
÷ø
æ
=
C
mole
K
molK
RT
F
K
ECell 0 = 0.0257 ln = 0.0592 log
n
K
n

34. Effect of Concentration on Cell Potential
DG =DG0 + RTlnQ
DG0 = -nFE0
cell
-nFEcell= -nFE0
cell + RTln Q
Ecell= E0
cell - RTln Q
nF
Ecell= E0
cell - 0.0257ln Q
n
Ecell= E0
cell – 0.0592log Q
n

35. Corrosion – Deterioration of Metals
by Electrochemical Process

36. Corrosion – Deterioration of Metals
by Electrochemical Process

37. Corrosion – Deterioration of Metals
by Electrochemical Process

39. Cathodic Protection

40. Abbreviated Standard Reduction
Potential Table

42. Batteries
19.6
Dry cell
Leclanché cell
Anode: Zn (s) Zn2+ (aq) + 2e-
Cathode: 2NH4
+ (aq) + 2MnO2 (s) + 2e- Mn2O3 (s) + 2NH3 (aq) + H2O (l)
Zn (s) + 2NH4 (aq) + 2MnO2 (s) Zn2+ (aq) + 2NH3 (aq) + H2O (l) + Mn2O3 (s)

43. Batteries
Mercury Battery
Zn(Hg) + 2OH- (aq) ZnO Anode: (s) + H2O (l) + 2e-
Cathode: HgO (s) + H2O (l) + 2e- Hg (l) + 2OH- (aq)
Zn(Hg) + HgO (s) ZnO (s) + Hg (l)
19.6

44. Batteries
2- (aq) PbSO4 (s) + 2e-
2- (aq) + 2e- PbSO4 (s) + 2H2O (l)
19.6
Lead storage
battery
Anode:
Cathode:
Pb (s) + SO4
PbO2 (s) + 4H+ (aq) + SO4
Pb (s) + PbO2 (s) + 4H+ (aq) + 2SO4
2- (aq) 2PbSO4 (s) + 2H2O (l)

45. Fuel Cell vs. Battery
• Battery; Energy storage device
– Reactant chemicals already in device
– Once Chemicals used up; discard (unless rechargeable)
• Fuel Cell; Energy conversion device
– Won’t work unless reactants supplied
– Reactants continuously supplied; products continuously
removed

46. Fuel Cell
A fuel cell is an
electrochemical cell
that requires a
continuous supply of
reactants to keep
functioning
2H2 (g) + 4OH- (aq) 4H2O (l) + 4e-
Anode:
Cathode: O2 (g) + 2H2O (l) + 4e- 4OH- (aq)
2H2 (g) + O2 (g) 2H2O (l)

47. Types of Electrochemical Cells
• Voltaic/Galvanic Cell; Energy released
from spontaneous redox reaction can be
transformed into electrical energy.
• Electrolytic Cell; Electrical energy is used
to drive a nonspontaneous redox reaction.

49. Faraday’s Constant Redox Eqn
Molar Mass
Charge =(Current)(Time)

50. Rawat’s Creation-rwtdgreat@
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