Electrochemistry

Applications of Redox

ReviewOxidation reduction reactions involve a

transfer of electrons.OIL- RIGOxidation Involves LossReduction Involves Gain

Solid lead(II) sulfide reacts with oxygen in the air at high temperatures to form lead(II) oxide and sulfur dioxide. Which substance is a reductant (reducing agent) and which is an oxidant (oxidizing agent)?  

A. PbS, reductant; O2, oxidant 

B. PbS, reductant; SO2, oxidant 

C. Pb2+, reductant; S2- oxidant 

D. PbS, reductant; no oxidant 

E. PbS, oxidant; SO2, reductant

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.

e-e- e-

e-e-

H+

MnO4-

Fe+2

Galvanic Cell

Salt Bridge allows current to flow

H+

MnO4-

Fe+2e-

Electricity travels in a complete circuit

H+

MnO4-

Fe+2

Porous Disk

Instead of a salt bridge

Reducing Agent

Oxidizing Agent

e-

e-

e- e-

e-

e-

Anode Cathode

Cell PotentialOxidizing agent pulls the electron.Reducing agent pushes 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.

Eºcell = EºZn Zn+2 + EºCu+2 Cu

We can look up reduction potentials in a table.

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

Cell Potential Determine 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

Reduction potentialMore negative Eº

– more easily electron is added– More easily reduced– Better oxidizing agent

More positive Eº – more easily electron is lost– More easily oxidized– Better reducing agent

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.

Cu2+ Fe+2

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

Under standard conditions, which of the following is the net reaction that occurs in the cell?

Cd|Cd2+ || Cu2+|Cu  

a.  Cu2+ + Cd → Cu + Cd2+  

b.  Cu + Cd → Cu2+ + Cd2+  

c.  Cu2+ + Cd2+ → Cu + Cd  

d.  Cu + Cd 2+ → Cd + Cu2+ 

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

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

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

w = -qE = -nFE = G

Potential, Work and G Gº = -nFEº if Eº > 0, then Gº < 0 spontaneous if Eº< 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)

nF2Al(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)

nFEº = RTln(K)

nF

nF Eº = ln(K) RT

Batteries are Galvanic CellsCar batteries are lead storage batteries.Pb +PbO2 +H2SO4 PbSO4(s) +H2O

Batteries are Galvanic CellsDry Cell

Zn + NH4+ +MnO2

Zn+2 + NH3 + H2O + Mn2O3

Batteries are Galvanic CellsAlkaline

Zn +MnO2 ZnO+ Mn2O3 (in base)

Batteries are Galvanic CellsNiCad 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 Fe2O3 + H+ Reactions happens in two places.

Water

Rust

Iron Dissolves-

Fe Fe+2

e-

Salt speeds up process by increasing conductivity

O2 + 2H2O +4e- 4OH-

Fe2+ + O2 + 2H2O Fe2O3 + 8 H+

Fe2+

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+

Calculating plating1. Current x time = charge

2. Charge ∕Faraday = mole of e-

3. Mol of e- to mole of element or compound

4. Mole to grams of compound

Or the reverse if you want time to plate

Calculate the mass of copper which can be deposited by the passage of 12.0 A for 25.0 min through a solution of copper(II) sulfate.

How long would it take to plate 5.00 g Fe from an aqueous solution of Fe(NO3)3 at a current of 2.00 A?

Other usesElectrolysis of water.Separating 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.

RedoxKnow the table

2. Recognized by change in oxidation state.

3. “Added acid”

4. Use the reduction potential table on the front cover.

5. Redox can replace. (single replacement)

6. Combination Oxidizing agent of one element will react with the reducing agent of the same element to produce the free element.

I- + IO3- + H+ I2 + H2O

7. Decomposition.a) peroxides to oxidesb) Chlorates to chloridesc) Electrolysis into elements.d) carbonates to oxides

A way to rememberAn Ox – anode is where oxidation occursRed Cat – Reduction occurs at cathodeGalvanic cell- spontaneous- anode is

negativeElectrolytic cell- voltage applied to make

anode positive