Electrochemistry

Applications of Redox

ReviewOxidation reduction reactions involve a

transfer of electrons.OIL- RIGOxidation Involves GainReduction Involves LossLEO-GER Lose Electrons OxidationGain Electrons Reduction

ApplicationsMoving electrons is electric current.8H++MnO4

-+ 5Fe+2 +5e- Mn+2 + 5Fe+3 +4H2O

Helps to break the reactions into half reactions.

8H++MnO4-+5e- Mn+2 +4H2O

5(Fe+2 Fe+3 + e- ) In the same mixture it happens without

doing useful work, but if separate

H+

MnO4-

Fe+2

Connected this way the reaction startsStops immediately because charge builds

up.

H+

MnO4-

Fe+2

Galvanic Cell

Salt Bridge allows current to flow

H+

MnO4-

Fe+2

e-

Electricity travels in a complete circuit Instead of a salt bridge

H+

MnO4-

Fe+2

Porous Disk

Reducing Agent

Oxidizing Agent

e-

e-

e- e-

e-

e-

Anode Cathode

Cell PotentialOxidizing agent pushes the electron.Reducing agent pulls the electron. The push or pull (“driving force”) is called

the cell potential Ecell

Also called the electromotive force (emf) Unit is the volt(V) = 1 joule of work/coulomb of chargeMeasured with a voltmeter

Zn+2 SO4-

2

1 M HCl

Anode

0.76

1 M ZnSO4

H+

Cl-

H2 in

Cathode

1 M HCl

H+

Cl-

H2 in

Standard Hydrogen ElectrodeThis is the reference

all other oxidations are compared to

Eº = 0 º indicates standard

states of 25ºC, 1 atm, 1 M solutions.

Cell PotentialZn(s) + Cu+2 (aq) Zn+2(aq) + Cu(s)The total cell potential is the sum of the

potential at each electrode.

Eº cell = EºZn Zn+2 + Eº Cu+2 Cu

We can look up reduction potentials in a table.

One of the reactions must be reversed, so change it sign.

Cell PotentialDetermine the cell potential for a galvanic

cell based on the redox reaction.Cu(s) + Fe+3(aq) Cu+2(aq) + Fe+2(aq)

Fe+3(aq) + e- Fe+2(aq) Eº = 0.77 V

Cu+2(aq)+2e- Cu(s) Eº = 0.34 V

Cu(s) Cu+2(aq)+2e- Eº = -0.34 V

2Fe+3(aq) + 2e- 2Fe+2(aq) Eº = 0.77 V

Line Notation solidAqueousAqueoussolidAnode on the leftCathode on the rightSingle line different phases.Double line porous disk or salt bridge. If all the substances on one side are

aqueous, a platinum electrode is indicated.

For the last reaction Cu(s)Cu+2(aq)Fe+2(aq),Fe+3(aq)Pt(s)

Galvanic Cell The reaction always runs

spontaneously in the direction that produced a positive cell potential.

Four things for a complete description.

1) Cell Potential

2) Direction of flow

3) Designation of anode and cathode

4) Nature of all the components- electrodes and ions

PracticeCompletely describe the galvanic cell

based on the following half-reactions under standard conditions.

MnO4- + 8 H+ +5e- Mn+2 + 4H2O

Eº=1.51Fe+3 +3e- Fe(s)

Eº=0.036V

Potential, Work and Gemf = potential (V) = work (J) / Charge(C)E = work done by system / chargeE = -w/qCharge is measured in coulombs. -w = qE Faraday = 96,485 C/mol e-

q = nF = moles of e- x charge/mole e-

w = -qE = -nFE = G

Potential, Work and G Gº = -nFE º if E º < 0, then Gº > 0 spontaneous if E º > 0, then Gº < 0 nonspontaneous In fact, reverse is spontaneous. Calculate Gº for the following reaction: Cu+2(aq)+ Fe(s) Cu(s)+ Fe+2(aq)

Fe+2(aq) + e-Fe(s) Eº = 0.44 V

Cu+2(aq)+2e- Cu(s) Eº = 0.34 V

Cell Potential and Concentration

Qualitatively - Can predict direction of change in E from LeChâtelier.

2Al(s) + 3Mn+2(aq) 2Al+3(aq) + 3Mn(s)

Predict if Ecell will be greater or less than

Eºcell if [Al+3] = 1.5 M and [Mn+2] = 1.0 M if [Al+3] = 1.0 M and [Mn+2] = 1.5M if [Al+3] = 1.5 M and [Mn+2] = 1.5 M

The Nernst EquationG = Gº +RTln(Q) -nFE = -nFEº + RTln(Q)

E = Eº - RTln(Q) nF

2Al(s) + 3Mn+2(aq) 2Al+3(aq) + 3Mn(s) Eº = 0.48 V

Always have to figure out n by balancing. If concentration can gives voltage, then

from voltage we can tell concentration.

The Nernst Equation As reactions proceed concentrations of

products increase and reactants decrease.

Reach equilibrium where Q = K and Ecell = 0

0 = Eº - RTln(K) nF

Eº = RTln(K) nF

nFEº = ln(K) RT

Batteries are Galvanic CellsCar batteries are lead storage batteries.Pb +PbO2 +H2SO4 PbSO4(s) +H2ODry Cell

Zn + NH4+ +MnO2 Zn+2 + NH3 + H2O

Alkaline Zn +MnO2 ZnO+ Mn2O3 (in base)

NiCad NiO2 + Cd + 2H2O Cd(OH)2 +Ni(OH)2

CorrosionRusting - 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 Fe2 O3 + H+ Reaction happens in two places.

Water

Rust

Iron Dissolves- Fe Fe+2

e-

Salt speeds up process by increasing conductivity

Preventing CorrosionCoating to keep out air and water.Galvanizing - Putting on a zinc coatHas 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.

Running a galvanic cell backwards.Put a voltage bigger than the potential

and reverse the direction of the redox reaction.

Used for electroplating.

Electrolysis

1.0 M

Zn+2

e- e-

Anode Cathode

1.10

Zn Cu1.0 M

Cu+2

1.0 M

Zn+2

e- e-

AnodeCathode

A battery >1.10V

Zn Cu1.0 M

Cu+2

Calculating platingHave to count charge.Measure current I (in amperes)1 amp = 1 coulomb of charge per secondq = I x tq/nF = moles of metalMass of plated metalHow long must 5.00 amp current be

applied to produce 15.5 g of Ag from Ag+

Other usesElectroysis of water.Seperating mixtures of ions.More positive reduction potential means

the reaction proceeds forward. We want the reverse.Most negative reduction potential is

easiest to plate out of solution.