lecture 10: electrochemistry · voltaic cells are convenient energy sources • batteries is a...

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)


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)


Glass Electrodes: Electrode consist of a very thin membrane made of special type of glass that is permeable to H+

ions


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


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



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--

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

lecture 10: electrochemistry · voltaic cells are convenient energy sources • batteries is a...

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