Electrochemistry

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Overview

Electrochemistry

½ - reactions & E0 values

The Galvanic CellThe Electrolytic Cell

Oxidation Numbers

Quiz A ½ Cells

Building aGalvanic

Cell

General Types of Batteries

CommonTypes of Batteries

Zn / C Ni / Cd Pb / Acid

Quiz B

Quiz CElectrolysis in

Industry

Quiz F Cl2

Quiz E

Diaphragm Membrane Mercury

Al

Principle Components of

a battery

Corrosion

Quiz G

Quiz D

Rules

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Rules for Assigning Oxidation Numbers

1. The oxidation number of an elemental substance is zero.

2. The oxidation number of a monoatomic ion (simple ion) is equal to the charge of the ion.

3. The sum of the oxidation numbers of all the atoms in a species is equal to the charge of the species.

4. Assign hydrogen an oxidation number of +1 when it is combined with a non-metal and -1 when it is combined with a metal.

5. Assign fluorine an oxidation number of -1 in all its compounds.

6. Assign oxygen an oxidation number of -2.

Note: these rules are in order of importance.

e.g. Rule 5 is paramount over Rule 6 if there is a potential conflict between the rules.

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Quiz A – Oxidation Numbers

1. Assign an oxidation number to each of the elements in the following species.

a) Oxygen gas, O2

b) Water, H2O

c) Hydrogen peroxide, H2O2

d) The perchlorate ion, ClO4-

e) Sulfurous acid, H2SO4

f) Oxygen difluoride, OF2

g) Ozone gas, O3

h) Sodium thiosulfate, Na2S2O3

i) Potassium tetrathionate, K2S4O6

j) Sodium hydride, NaH

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½ - reactions & E0 values

Written as reductions :

The more positive the Eo, the stronger the oxidizing agent on the LHS.

The more negative the Eo, the stronger the reducing agent on the RHS.

Ox + ne- Red

F2(g) + 2e- 2F-(aq)

Zero on the relative scale.

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Standard Hydrogen Electrode

1 M H+ (H3O+)T = 25 oC (298 K)

Pt electrode

PH2 (g) = 1 atm

H2

2H+ 2e-

2e-

The standard hydrogen electrode is given an Eo value (by convention) of0.00V. All Eo values are quoted relative to this zero.

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The two types of ½ - Cell1. Simple metal/ion redox couple

Ni electrode

1M Ni2+

2. Inert electrode with ionic redox couple

1M Fe3+

Pt electrode

1M Fe2+

e.g. Ni2+/Ni

e.g. Pt with Fe3+/Fe2+

Inert electrodes include Pt, Pd & Carbon

Fe2+

Fe3+

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Rules for Balancing Aqueous Redox Equations

1. Identify the elements changing oxidation numbers.

2. Separate the reaction into an oxidation ½ -reaction and reduction ½-reaction.

3. Balance the two ½-reactions thus:

i) Balance the element that is changing oxidation state.

ii) Insert the correct number of electrons into the ½-reaction.

iii) Balance the charge by adding H+(acid) or OH-(base).

iv) Balance the number of oxygen atoms by the addition of water.

v) Check the hydrogens are balanced.

4. Adjust the two ½-reactions to transfer the same number of electrons.

5. Combine the two ½-reactions to give a balanced redox reaction.

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Quiz B- Eo values & Balancing Redox Reactions

2. Balance the following aqueous redox reactions.

a) Cr2O72- + Fe2+ Fe3+ + Cr3+ (acid)

b) MnO4- + C2O4

2- Mn2+ + CO2 (acid)

c) F2 + NaOH OF2 + NaF

d) OCl- + I- IO3- + Cl2 (acid)

1 a) Which is the stronger oxidizing agent : Cl2(g) or MnO4- (aq).

b) Which is the stronger reducing agent : Fe(s) or Zn(s).

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The Galvanic Cell

e.g. Cu2+ (aq) + 2e- → Cu (s)

Zn2+ (aq) + 2e- → Zn (s) ?E0 = +0.34V

E0 = -0.76VZn (s) → Zn2+ (aq) + 2e-

∆E0 = Eored – E0

oxCu2+ + Zn → Zn2+ + Cu = +0.34V – (-0.76V)∆E0 = +1.10V

Reduction at Cathode & Oxidation at AnodeREDCat & AnOx

All galvanic cells have positive ΔE0 values

All spontaneous redox reactions have postive ΔE0 values

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Building a Galvanic Cell

SO2-4

SO2-4

SO2-4

SO2-4

SO2-4

Zn2+

Zn2+

Zn2+

Zn2+

Zn2+

SO2-4

SO2-4

SO2-4

SO2-4

SO2-4

Cu2+

Cu2+

Cu2+

ZnCu

0.00

Cu

Zn2+

Zn

+1.10

K+

Cl-Cl-

Cl-Cl-Cl-Cl-Cl-

Cl-Cl-

Cl-Cl-

Cl-Cl-

Cl-Cl-

Cl-Cl-Cl-Cl-Cl-

Cl-Cl-Cl-Cl-

Cl-Cl-Cl-Cl-Cl-

Cl-

Cl-Cl-

Cl-Cl-

Cl-

Cl-Cl-

Cl-Cl-

K+K+K+K+

K+K+K+

K+K+

K+K+K+

K+K+

K+K+K+ K+K+

K+K+

K+K+

K+K+

K+K+

K+K+

K+

K+

K+

K+

K+

K+

K+

Cl-Cl-

Cl-Cl-Cl-Cl-Cl-

Cl-Cl-

Cl-Cl-

Cl-Cl-

Cl-Cl-

Cl-Cl-Cl-Cl-Cl-

Cl-Cl-Cl-Cl-

Cl-Cl-Cl-Cl-Cl-

Cl-

Cl-Cl-

Cl-Cl-

Cl-

Cl-Cl-

Cl-Cl-

K+K+K+K+

K+K+K+

K+K+

K+K+K+

K+K+

K+K+K+ K+K+

K+K+

K+K+

K+K+

K+K+

K+K+

K+

K+

K+

K+

K+

K+

+0.95

Cu2+ (aq) + 2e- → Cu (s)Zn (s) → Zn2+ (aq) + 2e-

Cu2+

Cu2+

Cu2+ + Zn → Zn2+ + Cu

Qinit =[Zn2+]

[Cu2+]= 1

K =[Zn2+]

[Cu2+]= 1.6 x 1037

Cl-

K+

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Quiz C- Galvanic Cells

1. A galvanic cell involving the following two ½-cells : Mn2+/Mn (Eo = -1.18V) and Cr3+/Cr (Eo = -0.74V), is to be constructed.

a) Draw a diagram of the cell.

b) Show on the diagram the direction of the electron flow.

c) Show on the diagram the direction that the cations move in the salt bridge.

d) Write a balanced cell equation.

e) Calculate the Ecell at i) t = 0

ii) t = ∞

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General Types of Batteries

Primary Cell Non-rechargeable

Provides electricity until it dies (i.e. achieves equlibrium)

Disposable as redox couple is non-reversible

e.g. Zinc/Manganese battery (Dry cell)

Secondary Cell Rechargeable

Provides electricity until it goes flat

Connect to external power source to reverse redox reactions

e.g. Pb/PbSO4 battery

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Principle Components of a Battery

Anode Current CollectorCathode Current Collector

Anode active mass

Cathode active mass

Separator

Container

Terminals

Electrolyte

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Zn/C Dry Cell Battery

Zn(s) Zn2+(aq) + 2e-

2MnO2(s) + H2O(l) + 2e- Mn2O3(s) + 2OH-(aq)

Zn(s) + 2MnO2(s) + H2O(l) Zn2+(aq) + Mn2O3(s) + 2OH-(aq)

Electrolyte : NH4Cl / ZnCl2 / MnO2 / C Powder

Current collectors : Graphite & Zinc

Cathode

Anode

Zn cathode

Carbon anode

Carbon paste (carbon, electrolyte(NH4Cl) and MnO2)

Insulation

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Nickel / Cadmium Battery

Cd(s) + 2OH-(aq) Cd(OH)2(s) + 2e-

NiO(OH)(s) + H2O(l) + e- Ni(OH)2(s) + OH-

Electrolyte : KOH

Current collectors : Ni & Cd

2NiO(OH)(s) + Cd(s) + 2H2O(l) 2Ni(OH)2(s) + Cd(OH)2(s)

x2Cathode

Anode

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Lead / Acid Battery

Pb(s) + SO42-(aq) PbSO4(s) + 2e-

PbO2(s) + 4H+(aq) + SO42-(aq) + 2e- 2H2O(l) + PbSO4(s)

PbO2(s) + Pb(s) + 4H+(aq) + 2SO42-(aq) 2PbSO4(s) + 2H2O(l)

Electrolyte : H2SO4

Current collectors : Both Pb

Cathode

Anode

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Quiz D - Batteries

1. The lead storage battery is based upon the following ½ -cells:

PbSO4(s) + 2e- Pb(s) + SO42- (aq) E0 = -0.31 V

PbO2(s) + 4H+(aq) + SO42-(aq) + 2e- PbSO4 (s) + 2H2O(l) E0 = +1.70 V

a) Write down the overall redox reaction that occurs when the lead storage battery is being recharged.

b) Given a car battery is usually required to supply about 12 V, how many galvanic cells must be in series within a car battery?

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SO2-4

SO2-4

SO2-4

SO2-4

SO2-4

Zn2+Zn2+

Zn2+

Zn2+

Zn2+

Zn

SO2-4

SO2-4

SO2-4

SO2-4

SO2-4

Cu2+

Cu2+

Cu2+

Cu2+

Cu

The Electrolytic Cell

K+

Cl-Cl-

Cl-Cl-Cl-Cl-Cl-

Cl-Cl-

Cl-Cl-

Cl-Cl-

Cl-Cl-

Cl-Cl-Cl-Cl-Cl-

Cl-Cl-Cl-Cl-

Cl-Cl-Cl-Cl-Cl-

Cl-

Cl-Cl-

Cl-Cl-

Cl-

Cl-Cl-

Cl-Cl-

K+K+K+K+

K+K+K+

K+K+

K+K+K+

K+K+

K+K+K+ K+K+

K+K+

K+K+

K+K+

K+K+

K+K+

K+

K+

K+

K+

K+

K+

K+

Cl-Cl-

Cl-Cl-Cl-Cl-Cl-

Cl-Cl-

Cl-Cl-

Cl-Cl-

Cl-Cl-

Cl-Cl-Cl-Cl-Cl-

Cl-Cl-Cl-Cl-

Cl-Cl-Cl-Cl-Cl-

Cl-

Cl-Cl-

Cl-Cl-

Cl-

Cl-Cl-

Cl-Cl-

K+K+K+K+

K+K+K+

K+K+

K+K+K+

K+K+

K+K+K+ K+K+

K+K+

K+K+

K+K+

K+K+

K+K+

K+

K+

K+

K+

K+

K+

Cu2+

Zn2+ + Cu(s) Cu2+ + Zn(s)

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Quiz E- Electrolytic Cells

1. Design an electrochemical cell to produce Ce4+ and Au (gold).

Ce4+(aq) + e- Ce3+(aq) E0 = +1.61 V

Au3+(aq) + 3e- Au(s) E0 = +1.50 V

Your diagram should include the direction of the electron flow, the directionof anion movement in the salt bridge and the minimum potential required from thebattery to produce the desired products.

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Electrochemistry in Industry

Al2O3 dissolved in molten Na3AlF6 at 1000oC

C(s) + 2O2-→ CO2 + 4e-

Al3+ + 3e- → Al

+

-

x3

x4

3C(s) + 6O2- + 4Al3+ → 3CO2(g) + 4Al(s)

3C(s) + 2Al2O3(s)→ 3CO2(g) + 4Al(s)

Aluminium Production

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Electrochemistry in Industry

Al

Volatile hydrocarbons+ CO + CO2 + SO2

CoalTar

pitch

NaturalGas forheatingCoke

Carbon anode plant

Electrolysiscell

Filter to removeParticulate matter

HF + F2

removalAl2O3

purification

Bauxite frommine

CO2

SO2} to Air

Na3AlF6

Aluminium Production

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Chlor-Alkali Process

NaCl Na2CO3 (Soda Ash)

Cl2 + NaOH + H2

Cl2NaOH

Fuel Feedstock

Pulp & Paper

Soap

Bleach

Dyes

Textiles

Pulp & Paper

Plastics

Organochlorines

Bleach

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Chlor-alkali industries chartSodium chloride

ElectrolyticProcess

Chlorine CausticSoda

Pulp & PaperPlasticsSanitationBleachHerbicides

SoapRayonDyesPaperRubberTextilesBleachingNeutralisation

Limestone & Fuel

Carbon Dioxide

Sodiumbicarbonate

DrugsBeveragesBaking Powder

Ammonia

Soda Ash

Soaps

GlassDrugsPaperTextilesMetallurgyPetroleumWater softening

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Electrolytic Cells for Chlor-alkaliDiaphragm Cell

NaCl

Cl-

Na+

Cl-

Na+

Na+

Na+

Cl-

Cl-

Cl22e-

H2O

Permeable membrane. Both cations and anions can cross.

OH H

OH H

OH H

OH H

2e-

H2

O H-

O H-

30% NaOH containing Cl-

Needs Evaporation

Ste

el g

auze

cath

ode

Dim

ensi

onally

sta

ble

anode

4OH- → 2H2O + O2 + 4e-O2

O2

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Cl-Na+

Na+

O H-

Electrolytic Cells for Chlor-alkaliMembrane Cell

NaCl

Cl-

Na+

Cl-

Na+

Na+

Na+

Cl-

Cl-

Cl22e-

H2O

Cation permeable membrane. Cations can cross but not anions.

OH H

OH H

OH H

OH H

2e-

H2

O H-

O H-

50% NaOH

Ideally!

Ste

el g

auze

cath

ode

Dim

ensi

onally

sta

ble

anode

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Na+

O H-

Na+

O H-

http://www.ukzn.ac.za/index.asp

Electrolytic Cells for Chlor-alkaliMercury Cell

2Cl-(aq) → Cl2(g) + 2e-

Hg(l) + 2Na+(aq) + 2e- → 2HgNa(l)

2H2O(l) + 2HgNa(l) → 2NaOH(aq) + 2Hg(l) + H2(g)

Cell 17% NaCl

NaHg amalgam

H2(g)

50% NaOH

Recycle Hg

H2O

35% NaCl

Cl2(g)

Denuder

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Typical data for Chlor-alkali cells

MERCURY CELL

DIAPHRAGM CELL

MEMBRANE CELL

Cell voltage (V) -4.4 -3.45 -3.5

Current density (A cm-2) 1.0 0.2 0.45

Current efficiency for Cl2 (%) 97 96 93

Energy consumption (kWh /ton NaOH)

(a) Electrolysis only 3150 2550 2700

(b) Electrolysis + evaporation to 50% NaOH 3150 3260 2920

Purity Cl2(%) 99.2 98 99.3

Purity H2 (%) 99.9 99.9 99.9

O2 in Cl2(%) 0.1 1-2 0.3

Cl- in 50% NaOH (%) 0.003 1-1.2 0.005

Need for evaporation to 50% NaOH No Yes Currently

Hg Pollution problem maybe No No

Table1: Typical data for recent commercial chlor-alkali cells

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Quiz F- Chlor-alkali data

2. By consulting Table 1 answer the following questions.

a) Which cell is favoured by the cell voltage?

b) Which cell is favoured by the current density?

c) Which cell is favoured by the current efficiency?

d) Which cell is favoured by the energy consumption?

1 a) Write down balanced half-cell reactions for the processes happening at the anode and the cathode in the membrane (and diaphragm) cell. b) Write down a balanced cell reaction for the membrane cell.

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Corrosion

Fe can exist in three different oxidation states : 0, +2 and +3. 1. Fe Fe2+ + 2e- E = -0.44 V 2. Fe Fe3+ + 3e- E = -0.06 V

O2 + 4H+ + 4e- 2H2O E = +1.23 V Therefore, for Fe Fe2+ E = +1.23 – (-0.44)

= +1.67 V and for Fe Fe3+ E = +1.23 – (-0.06) = +1.29 V

The common oxidizing agent in the environment is, of course, O2 gas.

We use iron(Fe) as a building material. The trouble is it rusts (i.e. it oxidizes).

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2Fe2+ + ½O2 + (2 + n)H2O Fe2O3.nH2O + 4H+

Therefore Fe Fe2+ + 2e- will be the first thing to happen

then Fe2+ Fe3+ + e- will happen subsequently .

so 2Fe 2Fe2+ + 4e- (Anode)

O2 + 4H+ + 4e- 2H2O (Cathode)

2Fe + O2 + 4H+ 2Fe2+ + 2H2O

2 Fe(s) + O2(g) + nH2O(l) Fe2O3.nH2O(s)23

Therefore corrosion requires O2/H2O to occur and will be catalysed by H+.

CorrosionFe is first oxidized to Fe2+ and subsequently to Fe3+.

then

Overall :

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Fe

Corrosion

Cu2+ + 2e- Cu E0 = + 0.34V

Fe2+ + 2e- Fe E0 = - 0.44V

Fe + Cu2+ Fe2+ + Cu ΔE0 =+ 0.78V

Spontaneous reaction

= Corrosion of Fe

Cu

Fe corrodes when it behaves as an anode.

Fe behaves as an anode.

Cu behaves as a cathode.

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Fe

Corrosion

Zn2+ + 2e- Zn E0 = - 0.76 V

Fe2+ + 2e- Fe E0 = - 0.44 V

Zn + Fe2+ Zn2+ + Fe ΔE0 =+ 0.32 V = Corrosion of Zn

Zn

Fe will not corrode when it behaves as a cathode.

Fe behaves as a cathode.

Zn behaves as an anode.

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Quiz G- Corrosion

1. Which of the following metals are suitable for use as sacrificial anodes to protect against corrosion of an underground iron tank? Explain your choices.

i) Aluminium ii) Silver iii) Nickel iv) Sodium

2. Why do steel bridge-supports rust at the water line but not above or below?

3. In Scotland, the common occurrence of ice, in winter, requires the of use salt on the roads. This lowers the freezing point of water and melts the ice. Scottish cars get rusty very quickly in the winter….Why?

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