electrochemistry
DESCRIPTION
Electrochemistry. Overview. Quiz A. Rules. ½ Cells. Quiz B. Oxidation Numbers. ½ - reactions & E 0 values. Quiz G. Corrosion. Electrochemistry. The Electrolytic Cell. The Galvanic Cell. Building a Galvanic Cell. Electrolysis in Industry. Quiz E. Quiz C. General - PowerPoint PPT PresentationTRANSCRIPT
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
2
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.
3
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
4
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.
5
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.
6
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+
7
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+
11
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 = ∞
12
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
13
Principle Components of a Battery
Anode Current CollectorCathode Current Collector
Anode active mass
Cathode active mass
Separator
Container
Terminals
Electrolyte
14
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
15
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
16
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
17
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)
19
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.
20
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
21
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
22
Chlor-Alkali Process
NaCl Na2CO3 (Soda Ash)
Cl2 + NaOH + H2
Cl2NaOH
Fuel Feedstock
Pulp & Paper
Soap
Bleach
Dyes
Textiles
Pulp & Paper
Plastics
Organochlorines
Bleach
23
Chlor-alkali industries chartSodium chloride
ElectrolyticProcess
Chlorine CausticSoda
Pulp & PaperPlasticsSanitationBleachHerbicides
SoapRayonDyesPaperRubberTextilesBleachingNeutralisation
Limestone & Fuel
Carbon Dioxide
Sodiumbicarbonate
DrugsBeveragesBaking Powder
Ammonia
Soda Ash
Soaps
GlassDrugsPaperTextilesMetallurgyPetroleumWater softening
24
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
25
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
26
Na+
O H-
Na+
O H-
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
28
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
29
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.
30
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).
31
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 :
32
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.
33
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.
34
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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