Transcript Redox Reactions

Redox Reactions and Electrochemistry
I.
Redox Reactions
a)
b)
II.
Oxidation Number
Oxidizing and Reducing Reagents
Galavanic or Voltaic Cells
a)
b)
c)
Anode/Cathode/Salt Bridge
Cell Notations
Determining Cell Potential/Cell Voltage/Electromotive force (emf)
Relating Cell Potential to K and DG0
IV. Effect of Concentration on Cell Potential
V. Corrosion
VI. Batteries
VII. Fuel Cells
VIII.Electrolytic Cells
III.
a)
Calculating amounts of substances reduced or oxidized
REDOX REACTION
DEFINITIONS


Oxidation: Loss of electrons.
Reduction: Gain of electrons.
LEO says GER
Oxidation: Gain of oxygen
 Reduction: Lost of oxygen

Oxidation: Increasing of oxidation number
 Reduction: Reducing of oxidation number

CO  CO2 (CO oxidized) why?
CH3COOH  CH3CHO (CH3COOH reduced) why?
H2SO3  H2SO4 ??
HNO3  HNO2 ??
Ca2+  Ca (Ca2+ reduced) why?
Na  Na+ (Na oxidized) why?
Fe3+  Fe2+
Mn2+  MnO4– (Mn2+ oxidized) why?
Cl2  2 Cl– (Cl2 reduced) why?
H2O2  H2O ??
NaH  H2 ??
LEO says GER :
Lose Electrons = Oxidation
Na0  Na+1 + 1e-
Sodium is oxidized
Gain Electrons = Reduction
Cl0 + 1e–  Cl–
Chlorine is reduced
Rules for Assignment
of Oxidation Number (ON)
1) The ON of all pure elements is zero.
2) The ON of H is +1, except in hydrides, where it is -1.
3) The ON of O is -2, except in peroxides, where it is -1.
4) The algebraic sum of ON must equal zero for a
neutral molecule or the charge on an ion.
Variable Oxidation Number of Elements











Sulfur: SO42-(+6), SO32-(+4), S(0), FeS2(-1), H2S(-2)
Carbon: CO2(+4), C(0), CH4(-4)
Nitrogen: NO3-(+5), NO2-(+3), NO(+2), N2O(+1), N2(0), NH3(-3)
Iron: Fe2O3(+3), FeO(+2), Fe(0)
Manganese: MnO4-(+7), MnO2(+4), Mn2O3(+3), MnO(+2), Mn(0)
Copper: CuO(+2), Cu2O(+1), Cu(0)
Tin: SnO2(+4), Sn2+(+2), Sn(0)
Uranium: UO22+(+6), UO2(+4), U(0)
Arsenic: H3AsO40(+5), H3AsO30(+3), As(0), AsH3(-1)
Chromium: CrO42-(+6), Cr2O3(+3), Cr(0)
Gold: AuCl4-(+3), Au(CN)2-(+1), Au(0)
BALANCING OVERALL REDOX REACTIONS
Example balance the redox reaction below:
Fe + Cl2  Fe3+ + ClStep 1: Assign oxidation number,
Fe0 + Cl20  Fe3+ + ClStep 2: Determine number of electrons lost or gained by
reactants.
Fe0 + Cl20  Fe3+ + Cl

3e2eStep 3: Cross multiply.
2Fe + 3Cl20  2Fe3+ + 6Cl-
may be written as the sum of two half-cell reactions:
2Fe  2Fe3+ + 6e- (oxidation)
3Cl20 + 6e-  6Cl- (reduction)
All overall redox reactions can be expressed as the sum of
two half-cell reactions, one a reduction and one an
oxidation.
The overall reaction:
2Fe + 3Cl20  2Fe3+ + 6Cl-
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Another example - balance the redox reaction:
FeS2 + O2  Fe(OH)3 + SO42-
Fe+2S20 + O20  Fe+3(OH)3 + S+6O42

15e4e4FeS2 + 15O2  Fe(OH)3 + SO424FeS2 + 15O2  4Fe(OH)3 + 8SO424FeS2 + 15O2 +14H2O  4Fe(OH)3 + 8SO42- + 16H+
This reaction is the main cause of acid generation in
drainage from sulfide ore deposits. Note that we get 4
moles of H+ for every mole of pyrite oxidized!
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Electrochemistry: Interconversion of electrical and
chemical energy using redox reactions
Oxidation Half-Reaction: Oxidation Involves Loss
of electrons
2Mg
2Mg2+ + 4e-
Reduction Half-Reaction: Reduction Involves Gain of
electrons
O2 + 4e2O2-
Net Redox Reaction: 2Mg + O2  2 Mg+2 + 2 O-2
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 formonoatomic
ions in ionic compound
CaBr2; Ca = +2, Br = -1
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
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.
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.
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. The oxidation number of oxygen is usually –2. In
H2O2 and O22- it is –1, in elemental form (O2 or O3)
it is 0.
H2O ; H=+1, O= -2
SO3; O = -2; S = +6
HCO3-
IF7
O = -2
H = +1
F = -1
7x(-1) + ? = 0
I = +7
3x(-2) + 1 + C = -1
C = +4
NaIO3
Na = +1
O = -2
3x(-2) + 1 + ? = 0
I = +5
Determination of Oxidizing and Reducing Agents
Determine oxidation number for all atoms in both the reactants
and products.
Look at same atom in reactants and products and see if
oxidation number increased or decreased.
If oxidation number decreased:substance reduced  Oxidizing Agent
If oxidation number increased; substance oxidized  Reducing Agent
Spontaneous Redox Reaction
Zn(s) + Cu+2 (aq) -> Cu(s) + Zn+2(aq)
Zn
Cu+2
time
Zn+2
Cu
Gets Smaller
Gets Larger
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.
Voltaic Cell
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
Cell Notation
1.
2.
3.
Anode
Salt Bridge
Cathode
Anode | Salt Bridge | Cathode
| : symbol is used whenever there is a different phase
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)
cathode
anode
More detail..
Zn (s)| Zn+2 (aq, 1M)| K(NO3) (saturated)|Cu+2(aq, 1M)|Cu(s)
anode
Salt bridge
cathode
Zn (s) + 2 H+(aq) -> H2 (g) + Zn+2 (aq)
K(NO3)
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
0
0
0
ECell
 Eoxidation
 Ereduction
E 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.
Reduction Reaction:
2e- + 2H+ (1 M)
E0 = 0 V
Standard hydrogen electrode (SHE)
H2 (1 atm)
Determining if Redox Reaction is Spontaneous
+ E°CELL spontaneous reaction
0 E°CELL equilibrium
- E°CELL nonspontaneous reaction
More positive E°CELL stronger
oxidizing agent or more 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
ECell
work

ch arg e
Charge = nF
Units
Work: Joule
Charge (Q): Coulomb
Ecell : Volts
Faraday (F): charge on 1 mole eF = 96485 C/mole
Work = (charge)Ecell = -nFEcell
G = work (maximum)
G = -nFEcell
Relating EoCELL to the Equilibrium Constant, K
G0 = -RT ln K
G0 = -nFE0cell
E
0
Cell
RT

ln K
nF
0
Cell
E
-RT ln K = -nFE0cell
J 

8
.
31

298K 
RT 
molK 

 0.0257
C
F
96485
mole
0.0257
0.0592

ln K 
log K
n
n
Effect of Concentration on Cell Potential
G = G0 + RTlnQ
G0 = -nFE0cell
Ecell= E0cell - 0.0257ln Q
n
-nFEcell = -nFE0cell + RTln Q
Ecell = E0cell - RTln Q
nF
Ecell= E0cell – 0.0592log Q
n
Corrosion – Deterioration of Metals
by Electrochemical Process
Cathodic Protection
Batteries
Dry cell
A: Zn (s)
Zn2+ (aq) + 2eC: 2NH4+ (aq) + 2MnO2 (s) + 2eZn (s) + 2NH4 (aq) + 2MnO2 (s)
Mn2O3 (s) + 2NH3 (aq) + H2O (l)
Zn2+ (aq) + 2NH3 (aq) + H2O (l) + Mn2O3 (s)
Batteries
Mercury Battery
Anode:
Zn(Hg) + 2OH- (aq)
Cathode:
HgO (s) + H2O (l) + 2e-
Zn(Hg) + HgO (s)
ZnO (s) + H2O (l) + 2eHg (l) + 2OH- (aq)
ZnO (s) + Hg (l)
Batteries
Lead storage
battery
Anode:
Cathode:
Pb (s) + SO42- (aq)
PbSO4 (s) + 2e-
PbO2 (s) + 4H+ (aq) + SO42- (aq) + 2e-
Pb (s) + PbO2 (s) + 4H+ (aq) + 2SO42- (aq)
PbSO4 (s) + 2H2O (l)
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:
2H2 (g) + 4OH- (aq)
Cathode: O2 (g) + 2H2O (l) + 4e2H2 (g) + O2 (g)
4H2O (l) + 4e4OH- (aq)
2H2O (l)
Charge =(Current)(Time)
Molar Mass
Faraday’s Constant
Redox Eqn