Redox and ElectrochemistryBy-
Saurav K. Rawat
(Rawat DA Greatt)1
Redox Reactions and Electrochemistry
I. Redox Reactionsa) Oxidation Numberb) Oxidizing and Reducing Reagents
II. Galavanic or Voltaic Cellsa) Anode/Cathode/Salt Bridgeb) Cell Notationsc) Determining Cell Potential/Cell Voltage/Electromotive
force (emf)
III. Relating Cell Potential to K and G0
IV. Effect of Concentration on Cell Potential
Redox Reactions and Electrochemistry
V. Corrosion
VI. Batteries
VII. Fuel Cells
VIII.Electrolytic Cellsa) Calculating amounts of substances reduced or
oxidized
Electrochemistry: Interconversion of electrical and chemical energy using redox reactions
Redox (Oxidation-Reduction) Reaction: Type of electrontransfer reaction. One substance gives up electrons;
the other accepts electrons.
OIL RIG
•Oxidation Half-Reaction; Oxidation Involves Loss of electrons
•Reduction Half-Reaction; Reduction Involves Gain of electrons
gge
e
Net Redox Rxn; 2Mg + O2 -> 2 Mg+2 + 2 O-2
Oxidation numberThe charge the atom would have in a molecule (or anionic compound) if electrons were completely transferredto 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
Oxidation number,continuedThe charge the atom would have in a molecule (or anionic compound) if electrons were completely transferredto 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.
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
6. The oxidation number of fluorine is always –1. (unless fluorine is in elemental form, F2)
H2O ; H=+1, O= -2SO3; O = -2; S = +6
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.
HCO3-
O = -2 H = +1
3x(-2) + 1 + ? = -1
C = +4
Oxidation numbers of all the atoms in HCO3
- ?
4.4
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
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
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.
Spontaneous Redox ReactionZn(s) + Cu+2 (aq) -> Cu(s) + Zn+2(aq)
Cu+2
Zn
time Zn+2Cu
Gets Smaller -> <- Gets Larger
Voltaic Cell Animation
Anode; Site of OxidationCathode; Site of Reduction
AnOx or both vowelsRed 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
Cell Notation
1. Anode
2. Salt Bridge
3. Cathode
Anode | Salt Bridge | Cathode
| : symbol is used whenever there is a different phase
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
Zn (s)| Zn+2 (aq, 1M)| K(NO3) (saturated)|Cu+2(aq, 1M)|Cu(s)
anode cathodeSalt bridge
More detail..
K(NO3)
Zn (s) + 2 H+(aq) -> H2 (g) + Zn+2 (aq)
Zn(s)| Zn+2|KNO3|H+(aq)|H2(g)|Pt
Electrochemical Cells
The difference in electrical potential between the anode and cathode is called:
• cell voltage
• electromotive force (emf)
• cell potential
000reductionoxidationCell EEE
UNITS: Volts Volt (V) = Joule (J) Coulomb, C
Standard Electrode Potentials
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.
V
Standard hydrogen electrode (SHE)
eatm
Reduction Reaction
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 ormore likely to be reduced
• E0 is for the reaction as written
• The half-cell reactions are 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
Relating E0Cell to G0
ech
workECell arg
Unitswork, Joulecharge, CoulombEcell; Volts
charge = nFFaraday, F; charge on 1 mole e-F = 96485 C/mole
work = (charge)Ecell = -nFEcell
G = work (maximum)
G = -nFEcell
Relating CELL to the
Equilibrium Constant, KG0 = -RT ln K
G0 = -nFE0cell
-RT ln K = -nFE0cell
K
nF
RTECell ln0
0257.0
96485
29831.8
moleC
KmolK
J
F
RT
Kn
Kn
ECell log0592.0
ln0257.00
Effect of Concentration on Cell Potential
G =G0 + RTlnQ
G0 = -nFE0cell
-nFEcell= -nFE0cell + RTln Q
Ecell= E0cell - RTln Q
nF
Ecell= E0cell - 0.0257ln Q
nEcell= E0
cell – 0.0592log Q n
Corrosion – Deterioration of Metals by Electrochemical Process
Corrosion – Deterioration of Metals by Electrochemical Process
Corrosion – Deterioration of Metals by Electrochemical Process
Cathodic Protection
Abbreviated Standard Reduction Potential Table
Batteries
19.6
Leclanché cell
Dry cell
Zn (s) Zn2+ (aq) + 2e-Anode:
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)
Batteries
Zn(Hg) + 2OH- (aq) ZnO (s) + H2O (l) + 2e-Anode:
Cathode: HgO (s) + H2O (l) + 2e- Hg (l) + 2OH- (aq)
Zn(Hg) + HgO (s) ZnO (s) + Hg (l)
Mercury Battery
19.6
Batteries
19.6
Anode:
Cathode:
Lead storagebattery
PbO2 (s) + 4H+ (aq) + SO42- (aq) + 2e- PbSO4 (s) + 2H2O (l)
Pb (s) + SO42- (aq) PbSO4 (s) + 2e-
Pb (s) + PbO2 (s) + 4H+ (aq) + 2SO42- (aq) 2PbSO4 (s) + 2H2O (l)
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
Fuel Cell
A fuel cell is an electrochemical cell that requires a continuous supply of reactants to keep functioning
Anode:
Cathode: O2 (g) + 2H2O (l) + 4e- 4OH- (aq)
2H2 (g) + 4OH- (aq) 4H2O (l) + 4e-
2H2 (g) + O2 (g) 2H2O (l)
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.
Faraday’s Constant Redox Eqn
Molar Mass
Charge =(Current)(Time)
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