LECTURE 10: REDOX AND ELECTROCHEMISTRY

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LECTURE 10 REDOX AND ELECTROCHEMISTRY
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The redox process in compound formation.
Figure 4.9
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Sample Problem 4.7
Determining the Oxidation Number of an Element
(a) Zinc chloride
(b) Sulfur trioxide
(c) Nitric acid
PLAN
The O.N.s of the ions in a polyatomic ion add up
to the charge of the ion and the O.N.s of the
atoms or ions in a compound add up to zero.
SOLUTION
(a) ZnCl2. The O.N. for zinc is 2 and that for
chloride is -1.
(b) SO3. Each oxygen is an oxide with an O.N. of
-2. Therefore, the O.N. of sulfur must be 6.
(c) HNO3. H has an O.N. of 1 and each oxygen is
-2. Therefore, the N must have an O.N. of 5.
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Figure 4.10
Highest and lowest oxidation numbers of reactive
main-group elements.
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Sample Problem 4.8
Recognizing Oxidizing and Reducing Agents
PLAN
Assign an O.N. for each atom and see which atom
gained and which atom lost electrons in going
from reactants to products.
An increase in O.N. means the species was
oxidized (and is the reducing agent) and a
decrease in O.N. means the species was reduced
(is the oxidizing agent).
SOLUTION
The O.N. of Al increases Al is oxidized it is
the reducing agent.
The O.N. of H decreases H is reduced H2SO4 is
the oxidizing agent.
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Sample Problem 4.8
Recognizing Oxidizing and Reducing Agents
continued
The O.N. of C increases C is oxidized CO is the
reducing agent.
The O.N. of Pb decreases Pb is reduced PbO is
the oxidizing agent.
The O.N. of H increases it is oxidized H2 is
the reducing agent.
The O.N. of O decreases it is reduced O2 is the
oxidizing agent.
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Figure 4.11
A summary of terminology for oxidation-reduction
(redox) reactions.
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An active metal displacing hydrogen from water.
Figure 4.12
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Displacing one metal by another.
Figure 4.13
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Figure 4.14
The activity series of the metals.
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• TYPE OF REDOX REACTIONS
• Combination Reactions
• Decomposition Reactions
• Displacement reactions
• Combustion reactions

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Sample Problem 4.9
Identifying the Type of Redox Reaction
PROBLEM
Classify each of the following redox reactions as
a combination, decomposition, or displacement
reaction, write a balanced molecular equation for
each, as well as total and net ionic equations
for part (c), and identify the oxidizing and
reducing agents
PLAN
Combination reactions produce fewer products than
reactants.
Decomposition reactions produce more products
than reactants.
Displacement reactions have the same number of
products and reactants.
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Sample Problem 4.9
Identifying the Type of Redox Reaction
continued
(a) Combination
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Mg is the reducing agent N2 is the oxidizing
agent.
(b) Decomposition
or
H2O2 is the oxidizing and reducing agent.
(c) Displacement
Pb(NO3)2 is the oxidizing agent Al is the
reducing agent.
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REDOX TITRATIONS In redox titration, known
concentration of the oxidizing agent is used to
se the concentration of the reducing agent ( or
vice versa).
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• Redox Titration.

Redox titrations are used to find out information
about one reactant, using known information about
the other.
Worked example 1. Iron(II) sulphate can be
oxidised using acidified potassium permanganate
solution. Calculate the mass of iron(II)
sulphate which will completely react with 200 ml
of 0.25 mol l-1 acidified permanganate solution.
Write the redox equation- 2MnO4- 16H
5 Fe2 ? 2Mn2 8H2O 5Fe3
Calculate the number of moles of the known
reactant No. of moles C x V(litres)
0.25 x 0.2 0.05 mol
Use mole ratio in equation to calculate the
number of moles of the unkown reactant. 0.05
mol of MnO4- reacts with 5/2 x 0.05 0.125 mol
of Fe2
Use mass no. of mole x gfm to calculate the
mass of iron(II) sulphate Mass of FeSO4
number of moles x gfm
Mass of FeSO4 0.125 x 152 19 g
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• Calculations for you to try.
• Iron(II) ions react with acidified dichromate
  solution as shown below-
• 6Fe2 Cr2O72- 14H ?
  6Fe3 2Cr3 7H2O
• Calculate the number of moles of iron(II)
  ions which will completely
• react with 25cm3 of 0.4 mol l-1 dichromate
  solution.

Number of moles of dichromate C x
V(litres) 0.4 x 25/1000 0.01
From the mole ratio in the balanced equation
number of moles of iron(II) 6 x 0.01
0.06
Higher Grade Chemistry
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2. Hydrogen peroxide reacts with acidified
permanganate solution as shown below-
5H2O2 2MnO4- 6H ? 2Mn2 5O2
8H2O 100 cm3 of hydrogen peroxide
solution reacts with 10 cm3 of 0.2 mol l-1
permanganate solution. Calculate the
concentration of the hydrogen peroxide solution.
Number of moles of permanganate C x
V(litres) 0.2 x 10/1000
0.002
From the mole ratio in the balanced equation
number of moles of H2O2 5/2 x 0.002
0.005
Use Concentration number of moles/Volume
(litres) 0.005 / 0.1
0.05 mol l-1
Higher Grade Chemistry
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ELECTRON TRANSFER REACTIONS

• Electron transfer reactions are
  OXIDATION-REDUCTION or REDOX reactions.
• Results in the generation of an electric current
  (electricity) or be caused by imposing an
  electric current.
• Therefore, this field of chemistry is often
  called ELECTROCHEMISTRY.

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ELECTROCHEMICAL CELLS
ANODE Zn (s) Zn2(aq)
2e- CATHODE Cu2(aq) 2e- Cu(s)
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ELECTROCHEMICAL CELL CELL DIAGRAM

• Zn(s)ZnSO4(1.00M)CuSO4(1.00 M)Cu(s)
• Vertical lines separates phase boundary
• Double vertical lines denotes salt bridge
• Anode written first (left of the salt bridge)
• Concentration of solution, pressures of gases are
  indicated in cell diagrams

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ELECTROMOTIVE FORCE, EMF (E)

• The potential difference between the anode and
  cathode in a cell is called the electromotive
  force (emf).
• It is also called the cell potential, and is
  designated E.
• For the cell in the diagram E 1.104 v at 25oC
  and 1.0 molar solutions of Zn2 and Cu2
• EMF of a cell is normally measured using
  potentiometers.

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SINGLE ELECTRODE POTENTIALS

• Potential of all electrodes are measured in
  reference to a STANDARD HYDROGEN ELECTRODE.
• By definition, the reduction potential for
  hydrogen is 0 V
• 2 H (aq, 1M) 2 e- ??? H2 (g, 1 atm)

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STANDARD REDUCTION POTENTIALS
Reduction potentials for many electrodes have
been measured and tabulated.
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STANDARD CELL POTENTIAL
Because cell potential is based on the potential
energy per unit of charge, it is an intensive
property. The reverse of a half-cell reaction
will have the same Eo value but of opposite sign.
The cell potential at standard conditions can be
found through this equation
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CELL POTENTIAL

• For the oxidation in this cell,
• For the reduction,

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CELL POTENTIAL
0.34 V - (-0.76 V) 1.10 V
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OXIDIZING AND REDUCING AGENTS

• The strongest oxidizers have the most positive
  reduction potentials.
• The strongest reducers have the most negative
  reduction potentials.

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OXIDIZING AND REDUCING AGENTS

• The greater the difference between the two, the
  greater the voltage of the cell.

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THERMODYNAMICS OF ELECTROCHEMICAL CELLS

• ?G for a redox reaction can be found by using
  the equation
• ?G -vFE
• where v is the number of moles of electrons
  transferred, and F is a constant, the Faraday.
• 1 F 96,485 C/mol 96,485 J/V-mol (96,500)

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FREE ENERGY
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SAMPLE PROBLEM

• Predict whether the following reaction would
  occur spontaneously under standard state
  conditions, calculate the equilibrium constant at
  25oC.
• Sn(s) 2Ag(aq) Sn2(aq) 2Ag(s)

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SAMPLE PROBLEM

• Based on the following electrode potentials
• Fe2(aq) 2e- Fe(s) Eo
  -0.447 v
• Fe3(aq) e- Fe2(aq) Eo
  0.771 v
• Calculate the standard reduction potential for
  the half-reaction
• Fe3(aq) 3e- Fe(s) Eo
  ???

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THE NERNST EQUATION

• Remember that
• ?G ?G? RT lnK
• This means
• -vFE -vFE? RT lnK

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THE NERNST EQUATION

• Dividing both sides by -nF, we get the Nernst
  equation

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CONCENTRATION CELLS

• Notice that the Nernst equation implies that a
  cell could be created that has the same substance
  at both electrodes.

• Therefore, as long as the concentrations are
  different, E will not be 0.

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SAMPLE PROBLEM

• Predict whether the following reaction would
  proceed spontaneously as written at 298 K
• Cd(s) Fe2(aq) Cd2(aq)
  Fe(s)
• Given that Cd2 0.15M and Fe2 0.68M

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TYPES OF ELECTRODES
METAL ELECTRODES Piece of metal immersed in a
solution containing the cations of the
metal. Galvanic cells employ gas
electrodes. Electrode reaction Mz(aq) ze-
M(s) Ex. Copper, Zinc, Ag,
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TYPES OF ELECTRODES
GAS ELECTRODES A gas (1 atm) over a solution
of cation/anion of the gas in an inert metal
electrode. Example is the Standard Hydrogen
Electrode with Pt metal. PtH2(g)H(aq)
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TYPES OF ELECTRODES
METAL INSOLUBLE SALT ELCTRODES Coating a
piece of metal with an insoluble salt of the same
metal. Example is a Ag-AgCl electrode. Ag(s)AgCl(
s)Cl-(aq)
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TYPES OF ELECTRODES
Glass Electrodes Electrode consist of a very
thin membrane made of special type of glass that
is permeable to H ions
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TYPES OF ELECTRODES
ION SELECTIVE ELECTRODES Specific for cations
such as Li, Na, K Ag, and Cu2 and for anions
such as S2- and CN-.
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TYPES OF ELECTROCHEMICAL CELLS
CONCENTRATION CELLS Concentration cells
contain electrodes made of the same metal and
solutions containing the same ions but at
different concentrations.
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TYPES OF ELECTROCHEMICAL CELLS
FUELS CELLS HYDROGEN FUEL CELLS Hydrogen and
oxygen are bubbled through an electrolyte
solution (NaOH or H2SO4) with inert electrodes
also serving as catalysts. Anode Reaction H2(g)
2OH-(aq) 2H2O(l) 2e- Cathode
Reaction 1/2O2(g) ½O2(l) 2e 2OH-
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APPLICATIONS OF EMF MEASUREMENTS
DETERMINATION OF ACTIVITY COEFFICIENTS

• Example
• PtH2(1bar)HCl(1m)AgCl(s)Ag(s)
• Overall Reaction
• 1/2H2(g) AgCl(s) Ag(s) H(aq)
  Cl-(aq)
• And emf of the at 298 k is given as

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APPLICATIONS OF EMF MEASUREMENTS
DETERMINATION OF ACTIVITY COEFFICIENTS
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APPLICATIONS OF EMF MEASUREMENTS
pH Determination Ag(s)AgCl(s)HCl(aq),NaCl(aq)
HCl(aq)KCl(satd)Hg2Cl2(s)Hg(l) The overall
emf E for this arrangement is
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POTENTIOMETRIC REDOX TITRATIONS
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POTENTIOMETRIC REDOX TITRATIONS
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POTENTIOMETRIC REDOX TITRATIONS
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POTENTIOMETRIC REDOX TITRATIONS
From the standard emfs of the Fe2Fe3 and
Ce3Ce4 couples, calculate the equilibrium
constant for the following reaction at 298
K. The reaction in part (a) is employed in a
redox titration. Calculate the emf of the cell
after the addition of 10.0 mL of a 0.10 m Ce4
solution to a 50 mL of a 0.10 Fe2 solution.
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Applications of Oxidation-Reduction Reactions /
Electrochemistry
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Batteries
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Alkaline Batteries
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HYDROGEN FUEL CELLS
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Hydrogen fuel cells
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CORROSION AND
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CORROSION PREVENTION
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Corrosion

• Rusting - spontaneous oxidation.
• Most structural metals have reduction potentials
  that are less positive than O2 .
• Fe Fe2 2e- Eº 0.44 V
• O2 2H2O 4e- 4OH- Eº 0.40 V
• Fe2 O2 H2O Fe2O3 H
• Reactions happens in two places.

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Salt speeds up process by increasing conductivity
Water
Fe2 O2 2H2O Fe2O3 8 H
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Preventing Corrosion

• Coating to keep out air and water.
• Galvanizing - Putting on a zinc coat
• Has a lower reduction potential, so it is more
  easily oxidized.
• Alloying with metals that form oxide coats.
• Cathodic Protection - Attaching large pieces of
  an active metal like magnesium that get oxidized
  instead.

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Dry Cell Battery

• Anode (-)
• Zn ---gt Zn2 2e-
• Cathode ()
• 2 NH4 2e- ---gt 2 NH3 H2

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Alkaline Battery

• Nearly same reactions as in common dry cell, but
  under basic conditions.

Anode (-) Zn 2 OH- ---gt ZnO H2O
2e- Cathode () 2 MnO2 H2O 2e- ---gt
Mn2O3 2 OH-
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Mercury Battery

• Anode
• Zn is reducing agent under basic conditions
• Cathode
• HgO H2O 2e- ---gt Hg 2 OH-

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Lead Storage Battery

• Anode (-) Eo 0.36 V
• Pb HSO4- ---gt PbSO4 H 2e-
• Cathode () Eo 1.68 V
• PbO2 HSO4- 3 H 2e- ---gt PbSO4 2
  H2O

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Ni-Cad Battery

• Anode (-)
• Cd 2 OH- ---gt Cd(OH)2 2e-
• Cathode ()
• NiO(OH) H2O e- ---gt Ni(OH)2 OH-

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ELEC
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Harnessing the Power of Voltaic Cells
Batteries and Corrosion
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COMMERCIAL VOLTAIC CELLS

• Voltaic Cells are convenient energy sources
• Batteries is a self-contained group of voltaic
  cells arranged in series.
• Advantage Portable
• Disadvantage Very Expensive (US1.20 /
  Kwatt-h)
• Need cells in series to provide power

The Processes occurring during the discharge and
recharge of a lead-acid battery. When the
lead-acid battery is discharging (top) it behaves
like a voltaic cell the anode is negative
(electrode-1) and the cathode is positive
(electrode-2). When it is recharging (bottom),
it behaves like an electrolytic cell the anode
is positive (electrode-2) and the cathode is
negative (electrode-1).
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DRY CELL OR LeClanche cell
Dry Cells Invented in the 1860s the common dry
cell or LeClanche cell, has become a familiar
household item. An active zinc anode in the form
of a can house a mixture of MnO2 and an acidic
electrolytic paste, consisting of NH4Cl, ZnCl2,
H2O and starch powdered graphite improves
conductivity. The inactive cathode is a graphite
rod.
Anode (oxidation) Zn(s) g Zn2(aq)
2e- Cathode (reduction). The cathodic
half-reaction is complex and even today, is still
being studied. MnO2(s) is reduced to Mn2O3(s)
through a series of steps that may involve the
presence of Mn2 and an acid-base reaction
between NH4 and OH- 2MnO2 (s) 2NH4(aq)
2e- g Mn2O3(s) 2NH3(aq) H2O (l) The
ammonia, some of which may be gaseous, forms a
complex ion with Zn2, which crystallize in
contact Cl- ion Zn2(aq) 2NH3 (aq) 2Cl-(aq)
g Zn(NH3)2Cl2(s) Overall Cell reaction 2MnO2
(s) 2NH4Cl(aq) Zn(s) g Zn(NH3)2Cl2(s) H2O
(l) Mn2O3(s) Ecell 1.5 V Uses
common household items, such as portable radios,
toys, flashlights, Advantage Inexpensive, safe,
available in many sizes Disadvantages At high
current drain, NH3(g) builds up causing drop in
voltage, short shelf life because zinc anode
reacts with the acidic NH4 ions.
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DRY CELL OR LeClanche cell
Invented by George Leclanche, a French
Chemist. Acid version Zinc inner case that acts
as the anode and a carbon rod in contact with a
moist paste of solid MnO2 , solid NH4Cl, and
carbon that acts as the cathode. As battery wear
down, Conc. of Zn2 and NH3 (aq) increases
thereby decreasing the voltage. Half
reactions ECell 1.5 V Anode Zn(s) g
Zn2(aq) 2e- Cathode 2NH4(aq) MnO2(s)
2e- g Mn2O3(s) 2NH3(aq) H2O(l)
Advantage Inexpensive, safe, many
sizes Disadvantage High current drain, NH3(g)
build up, short shelf life
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Alkaline Battery
Alkaline Battery The alkaline battery is an
improved dry cell. The half-reactions are
similar, but the electrolyte is a basic KOH
paste, which eliminates the buildup of gases and
maintains the Zn electrode.
Anode (oxidation) Zn(s) 2OH- (aq) g ZnO(s)
H2O (l) 2e- Cathode (reduction). 2MnO2 (s)
2H2O (l) 2e- g Mn(OH)2(s)
2OH-(aq) Overall Cell reaction 2MnO2 (s) H2O
(l) Zn(s) g ZnO(s) Mn(OH)2(s) Ecell
1.5 V Uses Same as for dry
cell. Advantages No voltage drop and longer
shell life than dry cell because of alkaline
electrolyte sale ,amu sizes. Disadvantages More
expensive than common dry cell.
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ALKALINE BATTERY
Leclanche Battery Alkaline Version In alkaline
version solid NH4Cl is replaced with KOH or
NaOH. This makes cell last longer mainly because
the zinc anode corrodes less rapidly under basic
conditions versus acidic conditions. Half
reactions ECell 1.5 V Anode Zn(s)
2OH-(aq) g ZnO(s) H2O(l)
2e- Cathode MnO2 (s) H2O(l) 2e- g MnO3 (s)
2OH-(aq) Nernst equation E E -
(0.592/n)log Q, Q is constant !!
Advantage No voltage drop, longer shelf
life. Disadvantage More expensive
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Alkaline Batteries
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MERCURY BUTTON CELL
Mercury and Silver batteries are similar. Like
the alkaline dry cell, both of these batteries
use zinc in a basic medium as the anode. The
solid reactants are each compressed with KOH, and
moist paper acts as a salt bridge. Half
reactions ECell 1.6 V Anode Zn(s)
2OH-(aq) g ZnO(s) H2O(l) 2e- Cathode
(Hg) HgO (s) 2H2O(l) 2e- g Hg(s)
2OH-(aq) Cathode (Ag) Ag2O (s) H2O(l) 2e-
g 2Ag(s) 2OH-(aq)
Advantage Small, large potential, silver is
nontoxic. Disadvantage Mercury is toxic, silver
is expensive.
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LEAD STORAGE BATTERY

• Lead-Acid Battery. A typical 12-V lead-acid
  battery has six cells connected in series, each
  of which delivers about 2 V. Each cell contains
  two lead grids packed with the electrode
  material the anode is spongy Pb, and the
  cathode is powered PbO2. The grids are immersed
  in an electrolyte solution of 4.5 M H2SO4.
  Fiberglass sheets between the grids prevents
  shorting by accidental physical contact. When
  the cell discharges, it generates electrical
  energy as a voltaic cell.

Half reactions ECell 2.0 V Anode Pb(s)
SO42- ? PbSO4 (s) 2 e- E 0.356 Cathode
(Hg) PbO2 (s) SO42- 4H 2e- ? PbSO4
(s) 2 H2O E 1.685V Net PbO2 (s) Pb(s)
2H2SO4 ? PbSO4 (s) 2 H2O ECell 2.0 V
Note hat both half-reaction produce Pb2 ion, one
through oxidation of Pb, the other through
reduction of PbO2. At both electrodes, the Pb2
react with SO42- to form PbSO4(s)
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NICKEL CADMIUM BATTERY
Battery for the Technological Age Rechargeable,
lightweight ni-cad are used for variety of
cordless appliances. Main advantage is that the
oxidizing and reducing agent can be regenerated
easily when recharged. These produce constant
potential. Half reactions ECell 1.4 V Anode
Cd(s) 2OH-(aq) ? Cd(OH)2 (s) 2e-
Cathode 2Ni(OH) (s) 2H2O(l) 2e- ? Ni(OH)2
(s) 2 OH-(aq)
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FUEL CELLS
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FUEL CELLS BATTERIES
Fuel Cell also an electrochemical device for
converting chemical energy into electricity. In
contrast to storage battery, fuel cell does not
need to involve a reversible reaction since the
reactant are supplied to the cell as needed from
an external source. This technology has been
used in the Gemini, Apollo and Space Shuttle
program. Half reactions ECell 0.9 V Anode
2H2 (g) 4OH-(aq) ? 4H2O(l) 4e- Cathode
O2 (g) 2H2O(l) 4e- ? 4OH-(aq)
Advantage Clean, portable and product is water.
Efficient (75) contrast to 20-25 car, 35-40
from coal electrical plant Disadvantage Cannot
store electrical energy, needs continuous flow of
reactant, Electrodes are short lived and
expensive.
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Corrosion
Rust Fe2O3 X H2O Anode Fe(s) g Fe2
2e- E 0.44 V Cathode O2 (g) 4H 4e- g
2H2O (l) E 1.23 V Net Fe2 will further
oxidized to Fe2O3 X H2O
Not all spontaneous redox reaction are
beneficial. Natural redox process that oxidizes
metal to their oxides and sulfides runs billions
of dollars annually. Rust for example is not the
direct product from reaction between iron and
oxygen but arises through a complex
electrochemical process.
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CONDITIONS FOR CORROSION
Conditions for Iron Oxidation Iron will oxidize
in acidic medium SO2 g H2SO4 g H
HSO4 Anions improve conductivity for
oxidation. Cl- from seawater or NaCl (snow
melting) enhances rusting Conditions for
Prevention Iron will not rust in dry air
moisture must be present Iron will not rust in
air-free water oxygen must be present Iron rusts
most rapidly in ionic solution and low pH (high
H) The loss of iron and deposit of rust occur at
different placm on object Iron rust faster in
contact with a less active metal (Cu) Iron rust
slower in contact with a more active metal (Zn)
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IRON CORROSION CHEMISTRY
Most common and economically destructive form of
corrosion is the rusting of iron. Rust is not a
direct product of the reaction between iron and
oxygen but arises through complex electrochemical
process. The features of a voltaic cell can help
explain this process.
Iron will not rust in dry air moisture must be
present. Iron will not rust in air-free water
oxygen must be present Iron rusts most rapidly in
ionic solutions and at low pH (High H)
The loss of iron and the depositing of rust often
occur at different places on the same
object. Iron rust faster in contact with a less
active metal (such as Cu) and more slowly in
contact with a more active metal (such as Zn).
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CORROSION AND
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CORROSION PREVENTION
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PREVENTING CORROSION

• Coating to keep out air and water.
• Galvanizing - Putting on a zinc coat
• Has a lower reduction potential, so it is more
  easily oxidized.
• Alloying with metals that form oxide coats.
• Cathodic Protection - Attaching large pieces of
  an active metal like magnesium that get oxidized
  instead.

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Corrosion Prevention