LECTURE 10: ELECTROCHEMISTRY
ELECTRON TRANSFER REACTIONS
••
Electron transfer reactions are Electron transfer reactions are OXIDATIONOXIDATION--REDUCTIONREDUCTION
or or
REDOX reactions.REDOX reactions.
••
Results in the generation of an electric current (electricity) oResults in the generation of an electric current (electricity) or be r be
caused by imposing an electric current. caused by imposing an electric current.
••
Therefore, this field of chemistry is often called Therefore, this field of chemistry is often called
ELECTROCHEMISTRY.ELECTROCHEMISTRY.
ELECTROCHEMICAL CELLS
ANODE:
Zn (s) Zn+2(aq) + 2e-
CATHODE:
Cu+2(aq) + 2e- Cu(s)
ELECTROCHEMICAL CELL: CELL DIAGRAM
Zn(s)|ZnSO4(1.00M)||CuSO4(1.00 M)|Cu(s)
Vertical lines separates phase boundaryDouble vertical lines denotes salt bridgeAnode written first (left of the salt bridge)Concentration of solution, pressures of gases are indicated in cell
diagrams
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 Zn+2
and Cu+2
EMF of a cell is normally measured using potentiometers.
SINGLE ELECTRODE POTENTIALSPotential 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)
STANDARD REDUCTION POTENTIALS
Reduction potentials for
many electrodes have been measured
and tabulated.
STANDARD CELL POTENTIAL
The cell potential at standard conditions can be found through this equation:
oanode
ocathode
o EEE
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.
•
For the oxidation in this cell,
•
For the reduction,
Ered
= −0.76 V
Ered
= +0.34 V
CELL POTENTIAL
= +0.34 V −
(−0.76 V)= +1.10 V
CELL POTENTIAL
oanode
ocathode
o EEE
•
The strongest oxidizers have the most positive reduction potentials.
•
The strongest reducers have the most negative reduction potentials.
OXIDIZING AND REDUCING AGENTS
The greater the difference between the two, the greater the voltage of the cell.
OXIDIZING AND REDUCING AGENTS
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)
THERMODYNAMICS OF ELECTROCHEMICAL CELLS
FREE ENERGY:
reactioncellredox of constant mequilibriu the calculate can we E know we if thus
RTvFEKln
vFKlnRTE
vFG-E
:conditions standard Under
o
o
o
oro
Predict whether the following reaction would occur spontaneously under standard state conditions, calculate the equilibrium constant at 25oC.
Sn(s) + 2Ag+(aq) Sn+2(aq) + 2Ag(s)
SAMPLE PROBLEM:
Based on the following electrode potentials
Fe+2(aq) + 2e-
Fe(s) Eo
= -0.447 vFe+3(aq) + e-
Fe+2(aq) Eo
= 0.771 vCalculate the standard reduction potential for the half-
reaction:Fe+3(aq) + 3e-
Fe(s) Eo
= ???
SAMPLE PROBLEM:
•
Remember thatG = G
+ RT lnK
•
This means−vFE = −vFE
+ RT lnK
THE NERNST EQUATION:
Dividing both sides by −nF, we get the Nernst
equation:
Klnv
0.0257V-EE
conditions standard at ,KlnvFRTE
0E m,equilibriu at
KlnvFRTEE
o
o
o
THE NERNST EQUATION
•
Notice that the Nernst
equation implies that a cell could be created that has the same substance at both electrodes.
•
For such a cell, would be 0, but K would not.Ecell
•
Therefore, as long as the concentrations are different, E will not be 0.
CONCENTRATION CELLS
Predict whether the following reaction would proceed spontaneously as written at 298 K:
Cd(s) + Fe+2(aq) Cd+2(aq) + Fe(s)Given that [Cd+2] = 0.15M and [Fe+2] = 0.68M
SAMPLE PROBLEM:
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:
M+z(aq) + ze-
M(s)
Ex. Copper, Zinc, Ag,
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.
Pt|H2
(g)|H+(aq)
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)
TYPES OF ELECTRODES:
Glass Electrodes: Electrode consist of a very thin membrane made of special type of glass that is permeable to H+
ions
TYPES OF ELECTRODES:
ION SELECTIVE ELECTRODES: Specific for cations
such as Li+, Na+, K+
Ag+, and Cu+2
and for anions such as S2-
and CN-.
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.
TYPES OF ELECTROCHEMICAL CELLS:
FUELS CELLS: HYDROGEN FUEL CELLS
Hydrogen and oxygen are bubbled through an electrolyte solution (NaOH
or H2
SO4
) with inert electrodes also serving as catalysts.
Anode Reaction:
H2
(g) + 2OH-(aq) 2H2
O(l) + 2e-
Cathode Reaction:
1/2O2
(g) + ½O2
(l) + 2e 2OH-
Example:Pt|H2 (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:
APPLICATIONS OF EMF MEASUREMENTS
DETERMINATION OF ACTIVITY COEFFICIENTS
ClH
AgCl2/1
2H
AgClHo
aaln257.0EoEaf
aaaln257.0EE
APPLICATIONS OF EMF MEASUREMENTS
DETERMINATION OF ACTIVITY COEFFICIENTS
m. particular a at found be now can tcoefficien activity mean the ,1 0,m at since E get we 0, to m eextrapolat HCl, of mvs 0.514lnm) (E of Plot
ln514.0Emln0514.0Eln514.0mln514.0EE
)mln(257.0EE
mmaa
o
o
o
2o
2222ClH
APPLICATIONS OF EMF MEASUREMENTSpH Determination
Ag(s)|AgCl(s)|HCl(aq),NaCl(aq)||HCl(aq)||KCl(sat’d)|Hg2
Cl2
(s)|Hg(l)
The overall emf
E for this arrangement is:
0591.0EEpH
pH0591.0EEalog0591.0EE
ref
ref
Href
POTENTIOMETRIC REDOX TITRATIONS
2
2
2
o22
1
1
1
o11
2-
22
1-
11
21122112
OxRdln
v0.0257-EE
OxRdln
v0.0257-E E
Rd ev OxRd eV Ox
Ox v Rdv Rdv Oxv
)v(vEvEvE
point; eequivalenc at
21
o22
o11
eq
POTENTIOMETRIC REDOX TITRATIONS
V25.120.771)(1.72E
v 0.771E2 Fe e Fev1.72 E1 Ce e Ce
1vv ,Fe Ce Fe Ce
:example
eq
2-3
3-4
213324
POTENTIOMETRIC REDOX TITRATIONS
POTENTIOMETRIC REDOX TITRATIONS
From the standard emf’s
of the Fe2+|Fe3+
and Ce3+|Ce4+
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 Ce+4
solution to a 50 mL
of a 0.10 Fe+2
solution.
Applications of Oxidation- Reduction Reactions /
Electrochemistry
Batteries
Alkaline Batteries
HYDROGEN FUEL CELLS
Hydrogen fuel cells
CORROSION AND…
CORROSION PREVENTION
Corrosion
•
Rusting -
spontaneous oxidation.•
Most structural metals have reduction potentials that are less positive than O2
.
•
Fe Fe+2
+2e-
Eº= 0.44 V
•
O2
+ 2H2O + 4e-
4OH-
Eº= 0.40 V•
Fe+2
+ O2 + H2O Fe2O3 + H+
•
Reactions happens in two places.
WaterRust
Iron Dissolves-
Fe Fe+2
e-
Salt speeds up process by increasing conductivity
O2
+ 2H2
O +4e-
4OH-
Fe2+
+ O2
+ 2H2
O Fe2
O3
+ 8 H+
Fe2+
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.
Dry Cell BatteryDry Cell Battery
Anode (Anode (--))
Zn Zn ------> Zn> Zn2+2+
+ 2e+ 2e--
Cathode (+)Cathode (+)
2 NH2 NH44
++
+ 2e+ 2e--
------> > 2 NH2 NH33
+ H+ H22
Alkaline BatteryAlkaline BatteryNearly same reactions as in Nearly same reactions as in
common dry cell, but under common dry cell, but under basic conditions.basic conditions.
Anode (Anode (--): ): Zn + 2 OHZn + 2 OH--
------> > ZnOZnO
+ H+ H22
O + 2eO + 2e--Cathode (+): Cathode (+): 2 MnO2 MnO22
+ H+ H22
O + 2eO + 2e--
------> > MnMn22
OO33
+ 2 OH+ 2 OH--
Mercury BatteryMercury BatteryAnode:Anode:
Zn is reducing agent under basic conditionsZn is reducing agent under basic conditionsCathode:Cathode:
HgOHgO
+ H+ H22
O + 2eO + 2e--
------> Hg + 2 OH> Hg + 2 OH--
Lead Storage BatteryLead Storage Battery
Anode (Anode (--) ) EEoo
= +0.36 V= +0.36 VPbPb
+ HSO+ HSO44
--
------> PbSO> PbSO44
+ H+ H++
+ 2e+ 2e--Cathode (+) Cathode (+) EEoo
= +1.68 V= +1.68 VPbOPbO22
+ HSO+ HSO44
--
+ 3 H+ 3 H++
+ 2e+ 2e--
------> PbSO> PbSO44
+ 2 H+ 2 H22
OO
NiNi--Cad BatteryCad Battery
Anode (Anode (--))CdCd
+ 2 OH+ 2 OH--
------> Cd(OH)> Cd(OH)22
+ 2e+ 2e--Cathode (+) Cathode (+) NiO(OHNiO(OH) + H) + H22
O + eO + e--
------> Ni(OH)> Ni(OH)22
+ OH+ OH--
ELEC
Harnessing the Power of Voltaic Cells
Batteries and Corrosion
•
Voltaic Cells are convenient energy sources•
Batteries is a self-contained group of voltaic cells arranged in series.•
Advantage: Portable•
Disadvantage: Very Expensive (US$1.20 / 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).
COMMERCIAL VOLTAIC CELLS
Dry CellsInvented in the 1860’s 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 NH4 Cl, ZnCl2 , H2 O and starch powdered graphite improves conductivity. The inactive cathode is a graphite rod.
Anode (oxidation)Zn(s)
Zn2+(aq)
= 2e-Cathode (reduction).
The cathodic
half-reaction is complex and even today, is still being studied. MnO2(s)
is reduced to Mn2
O3(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-
Mn2
O3(s)
+ 2NH3(aq)
+ H2
O (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)
Zn(NH3
)2
Cl2(s)
Overall Cell reaction:Overall Cell reaction:2MnO2 (s)
+ 2NH4
Cl(aq)
+ Zn(s)
Zn(NH3
)2
Cl2(s)
+ H2
O (l)
+ Mn2
O3(s) Ecell
= 1.5 V
Uses:
common household items, such as portable radios, toys, flashlights,Advantage;Advantage;
Inexpensive, safe, available in many sizesDisadvantages: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.
DRY CELL OR LeClanche cell
Invented by George Leclanche, a French Chemist.Invented by George Leclanche, a French Chemist.
Acid version:Acid version:
Zinc inner case that acts as the anode and a carbon rod in conZinc inner case that acts as the anode and a carbon rod in contact with tact with a moist paste of solid MnOa moist paste of solid MnO22
, solid NH, solid NH44
Cl, and carbon that acts as the cathode. As battery Cl, and carbon that acts as the cathode. As battery wear down, Conc. of Znwear down, Conc. of Zn+2+2
and NHand NH3 (aq)3 (aq)
increases thereby decreasing the voltage.increases thereby decreasing the voltage.Half reactions:
E°Cell
= 1.5 VAnode: Zn(s)
Zn+2(aq)
+ 2e-
Cathode:
2NH4+
(aq)
+ MnO2(s)
+ 2e-
Mn2
O3(s)
+ 2NH3(aq)
+ H2
O(l)
Advantage:Inexpensive, safe, many sizes
Disadvantage:High current drain, NH3(g)
build up, short shelf life
DRY CELL OR LeClanche cell
Alkaline BatteryThe 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)
ZnO(s)
+ H2
O (l)
+ 2e-Cathode (reduction).2MnO2 (s)
+ 2H2
O (l)
+ 2e-
Mn(OH)2(s)
+ 2OH-(aq)
Overall Cell reaction:Overall Cell reaction:2MnO2 (s)
+ H2
O (l)
+ Zn(s)
ZnO(s)
+ Mn(OH)2(s)
Ecell
= 1.5 V
Uses:Uses:
Same as for dry cell.Advantages:Advantages:
No voltage drop and longer shell life than dry cell because of alkaline electrolyte; sale ,amu sizes.Disadvantages;Disadvantages;
More expensive than common dry cell.
Alkaline Battery
Leclanche Battery: Alkaline VersionLeclanche Battery: Alkaline VersionIn alkaline version; solid NHIn alkaline version; solid NH44
Cl is replaced with KOH or NaOH. This makes cell last Cl is replaced with KOH or NaOH. This makes cell last longer mainly because the zinc anode corrodes less rapidly underlonger mainly because the zinc anode corrodes less rapidly under
basic conditions versus basic conditions versus acidic conditions.acidic conditions.Half reactions:
E°Cell
= 1.5 VAnode: Zn(s)
+ 2OH-(aq)
ZnO(s)
+ H2
O(l)
+ 2e-
Cathode:
MnO2 (s)
+ H2
O(l)
+ 2e-
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
ALKALINE BATTERY
Alkaline Batteries
Mercury and Silver batteries are similar.Mercury and Silver batteries are similar.Like the alkaline dry cell, both of these batteries use zinc in Like the alkaline dry cell, both of these batteries use zinc in a basic medium as the anode. a basic medium as the anode. The solid reactants are each compressed with KOH, and moist papeThe solid reactants are each compressed with KOH, and moist paper acts as a salt bridge.r acts as a salt bridge.Half reactions:
E°Cell
= 1.6 VAnode: Zn(s)
+ 2OH-(aq)
ZnO(s)
+ H2
O(l)
+ 2e-
Cathode (Hg):
HgO (s)
+ 2H2
O(l)
+ 2e-
Hg(s)
+ 2OH-(aq)
Cathode (Ag):
Ag2
O (s)
+ H2
O(l)
+ 2e-
2Ag(s)
+ 2OH-(aq)
Advantage:Small, large potential, silver is nontoxic.
Disadvantage:Mercury is toxic, silver is expensive.
MERCURY BUTTON CELL
LeadLead--Acid Battery.Acid Battery.
A typical 12A typical 12--V leadV lead--acid battery has six cells acid battery has six cells connected in series, each of which delivers about 2 V. Each celconnected in series, each of which delivers about 2 V. Each cell l contains two lead grids packed with the electrode material: thecontains two lead grids packed with the electrode material: the anode is anode is spongy Pb, and the cathode is powered PbO2. The grids are immerspongy Pb, and the cathode is powered PbO2. The grids are immersed sed in an electrolyte solution of 4.5 M Hin an electrolyte solution of 4.5 M H22 SOSO44 . Fiberglass sheets between the . Fiberglass sheets between the grids prevents shorting by accidental physical contact. When thgrids prevents shorting by accidental physical contact. When the cell e cell discharges, it generates electrical energy as a voltaic cell.discharges, it generates electrical energy as a voltaic cell.
Half reactions: E°Cell
= 2.0 VAnode: Pb(s) + SO4
2-
PbSO4 (s) +2 e- E° = 0.356
Cathode (Hg): PbO2 (s) + SO42- + 4H+ + 2e-
PbSO4 (s) + 2 H2O E° = 1.685VNet: PbO2 (s) + Pb(s) + 2H2 SO4
PbSO4 (s) + 2 H2 O E°Cell
= 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)
LEAD STORAGE BATTERY
Battery for the Technological AgeBattery for the Technological AgeRechargeable, lightweight Rechargeable, lightweight ““nini--cadcad””
are used for variety of cordless appliances. Main advantage isare used for variety of cordless appliances. Main advantage is
that the oxidizing and reducing agent can be regenerated easily that the oxidizing and reducing agent can be regenerated easily when recharged. These produce when recharged. These produce constant potential.constant potential.Half reactions:
E°Cell
= 1.4 V
Anode: Cd(s) + 2OH-(aq)
Cd(OH)2 (s) + 2e-
Cathode:
2Ni(OH) (s) + 2H2 O(l) + 2e-
Ni(OH)2 (s) + 2 OH-(aq)
NICKEL CADMIUM BATTERY
FUEL CELLS
FUEL CELLS; BATTERIESFuel Cell also an electrochemical device for converting chemicalFuel Cell also an electrochemical device for converting chemical
energy energy into electricity.into electricity.In contrast to storage battery, fuel cell does not need to involIn contrast to storage battery, fuel cell does not need to involve a reversible reaction since ve a reversible reaction since the reactant are supplied to the cell as needed from an externalthe reactant are supplied to the cell as needed from an external
source. This technology source. This technology has been used in the Gemini, Apollo and Space Shuttle program.has been used in the Gemini, Apollo and Space Shuttle program.Half reactions:
E°Cell
= 0.9 V
Anode: 2H2 (g) + 4OH-(aq) 4H2 O(l) + 4e-
Cathode:
O2 (g) + 2H2 O(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.
Corrosion
Not all spontaneous redox reaction are beneficial.Not all spontaneous redox reaction are beneficial.Natural redox process that oxidizes metal to their oxides and suNatural redox process that oxidizes metal to their oxides and sulfides runs billions lfides runs billions of dollars annually. Rust for example is not the direct productof dollars annually. Rust for example is not the direct product
from reaction from reaction between iron and oxygen but arises through a complex electrochembetween iron and oxygen but arises through a complex electrochemical process.ical process.
Rust: FeRust: Fe22
OO3 3 ••
X HX H22
OOAnode:
Fe(s)
Fe+2
+ 2e-
E°
= 0.44 VCathode:
O2 (g)
+ 4H+
+ 4e-
2H2
O (l)
E°
= 1.23 VNet:Net:
FeFe+2+2
will further oxidized to Fewill further oxidized to Fe22
OO3 3 ••
X HX H22
OO
Conditions for Iron Oxidation:Iron will oxidize in acidic mediumIron will oxidize in acidic medium
SOSO22
HH22
SOSO44
HH++
+ HSO+ HSO44
++
Anions improve conductivity for oxidation.Anions improve conductivity for oxidation.ClCl--
from seawater or NaCl (snow melting) enhances rustingfrom seawater or NaCl (snow melting) enhances rusting
Conditions for Prevention:Conditions for Prevention:Iron will not rust in dry air; moisture must be presentIron will not rust in dry air; moisture must be presentIron will not rust in airIron will not rust in air--free water; oxygen must be presentfree water; oxygen must be presentIron rusts most rapidly in ionic solution and low pH (high HIron rusts most rapidly in ionic solution and low pH (high H++))The loss of iron and deposit of rust occur at different placm onThe loss of iron and deposit of rust occur at different placm on
objectobjectIron rust faster in contact with a less active metal (Cu)Iron rust faster in contact with a less active metal (Cu)Iron rust slower in contact with a more active metal (Zn)Iron rust slower in contact with a more active metal (Zn)
CONDITIONS FOR CORROSION
Iron will not rust in dry air; moisture must be present.Iron will not rust in air-free water; oxygen must be presentIron rusts most rapidly in ionic solutions and at low pH (High H+)
Most common and economically destructive form of Most common and economically destructive form of corrosion is the rusting of iron. Rust is not a direct corrosion is the rusting of iron. Rust is not a direct product of the reaction between iron and oxygen but product of the reaction between iron and oxygen but arises through complex electrochemical process. The arises through complex electrochemical process. The features of a voltaic cell can help explain this process.features of a voltaic cell can help explain this process.
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).
IRON CORROSION CHEMISTRY
CORROSION AND…
CORROSION PREVENTION
•
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.
PREVENTING CORROSION
Corrosion Prevention
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