74 trang 35 lượt tải

Lecture 10 Electrochemistry - Chemistry Laboratory | Trường Đại học Quốc tế, Đại học Quốc gia Thành phố HCM

69

Lecture 10 Electrochemistry - Chemistry Laboratory | Trường Đại học Quốc tế, Đại học Quốc gia Thành phố HCM được sưu tầm và soạn thảo dưới dạng file PDF để gửi tới các bạn sinh viên cùng tham khảo, ôn tập đầy đủ kiến thức, chuẩn bị cho các buổi học thật tốt. Mời bạn đọc đón xem!

Môn: Chemistry Laboratory (CH012IU) 59 tài liệu

Trường: Trường Đại học Quốc tế, Đại học Quốc gia Thành phố Hồ Chí Minh 696 tài liệu

Tác giả:

Profile

Mai Nguyệt

1 năm trước
Danh sách Quiz
/74
Lecture'10'
Electrochemistry-and-Redox-
titration-
Instructor:'Nguyen'ThaoTrang
Analytical Chemistry 1
Outlines
• Introduction.to.Electrochemistry
• Applications.of.Standard.electrode.potentials
• Redox.titration
2
th iế đ ện c cự
Outlines
• Introduction.to.Electrochemistry
• Applications.of.Standard.electrode.potentials
• Redox.titration
3
Basic.concepts.of.electrochemistry
• Redox.reaction.:.reactions.with.electron.transfer
– oxidized.:.loses.electrons
– reduced.:.gains.electrons
– oxidizing.agent.(oxidant).:.takes.electrons.from.others.and.becomes.
reduced
– reducing.agent.(reductant). :.gives.electrons. to.another.substance.and.
is.oxidized
– Example
4
Fe
3+.
is.an.oxidizing.agent.as:.Fe
3+
+.e
-
à Fe
2+
V V
2+.
is.an.reducing.agent.as:.V
2+
à
3+
+.e
-
Basic.concepts.of.electrochemistry
• Redox.reaction.:
– Half.reactions:.A.common.approach.for.listing.species.that.undergo.
REDOX.is.by.using.half.reactions.
à Redox.reaction.is.the.product.of.two.half.reactions
– Each.half.reaction.is.associated.with.a.standard.reduction.potential,.
E
o
’
5
Reduction:.Fe
3+
+.e
-
à Fe
2+
Oxidation:.V
2+
à V
3+ -
+.e
B
ox
+.e
-
→ B
red
,.E
o
’
A
red
→ A
ox
+.e
-
,.E
o
’
oxidixing
agent
reducing agent
half rxn (reduction rxn)
half rxn (oxidation rxn)
Basic.concepts.of.electrochemistry
• Redox.reaction.:
– Listed.half.reactions.with.standard.reduction.potential:
6
Basic.concepts.of.electrochemistry
• Redox.reaction.:
–
Example:.Determine.the.balanced.equation.for.the.reaction.of.Fe
2+
with.
Cr
2
O
7
2-
in.an.acidic.solution.
– Solution:
• From.the.standard.reduction.potentials,.we.can.get:
Fe
3+
+.e
-
=.Fe
2+
Cr
2
O
7
2-
+.14H
+
+.6e
-
=.2Cr 7H
3+
+.
2
O
• We.know.that.one.of.the.half.reactions.must.be.reversed.since..oxidation.
and.reduction.reactions.must.both.occur
•
Based.on.the.question,.Fe
2+.
must.undergo.oxidation.reaction:
à
reverse.reaction.(1):.Fe
2+.
=.Fe
3+ -
+.e
– To.balance. the.equation,. the.number.of.electrons. in.both.equation. must.be.
the.same.à multiply.(1.).by.6:
6.x.(1):.....................6Fe
2+.
=.6Fe
3+
+.6e
-
+.(2):..Cr +.14H +.6e =.2Cr
2
O
7
2- + - 3+
+.7H
2
O.
6Fe
2+.
+.Cr
2 7
O
2-
+.14H
+
=.6Fe
3+
+.2Cr
3+
+.7H
2
O.
7
E
o
= 0.77
E
o
= 1.33
E >E
Cr2O7(2-)/Cr(3+)
0
Fe(3+)/Fe(2+)
=> Cr2O7
2-
gains e more easily than Fe
3+
Basic.concepts.of.electrochemistry
• Electric.charge.q:.
– Is.measured.in.[coulomb].or.[C]
–
Magnitude.of.the.charge.of.a.single.electron.=.1.602.x.10
-19
C.
–
A.mole.of.electrons.has.a.charge.of.(1.602.x.10
-19
C)x.(6.022.x 10
23
mol
-1
).=.9.649.x.10
4
C/mol,.which.is.called.the.Faraday'constant'(F).
– Relationship.between. charge.and.moles:
8
where.n'is.the.number.moles.of.electrons.transferred.
Basic.concepts.of.electrochemistry
• Electric.charge.q:.
9
Basic.concepts.of.electrochemistry
• Electric.current:
– The.quantity.of.charge.flowing.each.second.through.a.circuit.is.called.
the.current.-
– The.unit.of.current.is.the.ampere,-abbreviated.A.
– A.current.of.1.A..represents.a.charge.of.1.C.per.second.flowing.past.a.
point.in.a.circuit.
10
FIGURE-13-1-Electrons.flowing.into.a.
coil.of.Ptwire.at.which.Sn
4+.
ions.in.
solution.are.reduced.to.Sn
2+
.
Basic.concepts.of.electrochemistry
• Voltage,.work,.free.energy:
– Difference.in.electric'potential,.E.(measured.by.volt).between.two.
points.is.the.work.needed.when.moving.an.electric.charge.from.one.
point.to.the.other:
– Work:..Energy.in.joules
• one.joule.of.energy.= .between.points.
whose.potential.differ.by.one.volt.
– Example:.How.much.work.is.required.to.move.2.36.mmolof.electrons.
through.a.potential.difference.of.1.05.V.?
–
11
2.39 x 10
2
J
Work.done.on.surroundings..=.- ΔG
à
à
à
àà ΔG.=.-E.q =.- nEF
< 0=> E>0 => Del E (potential diff) >0
Basic.concepts.of.electrochemistry
• Types.of.electrochemical.cell:
1. Galvanic-cell-(Voltaic):-
A.cell.generates.electrical.energy.from.a.spontaneous..
chemical.reaction.
2. Electrolytic-cell:
Is.where.non-spontaneous.redox.reactions.can.be.made.to.
occur.by.the.addition.of.electrical.energy.(process.is.termed.
electrolysis).
12
Basic.concepts.of.electrochemistry
• Common.components.of.a.electrochemical.cell
– Electrodes:.conduct.electricity.between.cell.and.
surroundings
• Anode:.oxidation.occurs.
• Cathode:.reduction.occurs.
– Electrolytes:.mixture.of.ions.involved.in.reaction.or.
carrying.charge
– Salt.bridge:.completes.circuit.(provides.charge.balance)
13
Basic.concepts.of.electrochemistry
• Galvanic.cell:
– A.simple.Galvanic.cell:.2.electrodes.in.a.solution.of.CdCl
2
:
One.electrode.is.cadmium;.the.other.is.metallic.silver.
coated.with.solid.AgCl.
14
Anode:......Cd.(s).
↔ Cd
2+
(aq) +.2e
-
Cathode:...2AgCl.(s).+.2e
-
↔ 2Ag(s).+.2Cl
-
(aq).
Cd.(s).+.2AgCl.
(s).↔ Cd
2+
(aq) +.2Ag(s).+.2Cl
-
(aq)
à Oxidation.of.Cd.electrode.provides.
electrons.that.flow.through.the.circuit.to.
the.Ag.electrode.
ΔG.=.-150.kJ/molCd
E
o
= - 0.44
E >E
AgCl/Ag
0
Cd(2+)/Cd
=> reverse the rxn at the anode
AgCl(s)+e- Ag(s)+Cl-(aq) E = 0.22
↔
0
Basic.concepts.of.electrochemistry
• Galvanic.cell:
– A.simple.Galvanic.cell:.2.electrodes.Cd.and.Ag.in.a.solution.
of.Cd(NO
3
)
2
and.AgNO
3.
:
15
A.little.current.flows.through.the.circuit.as.Ag.
can.react.at.the.Cd.electrode..
Not.really.
work!
the rxn still occur but no flow e
e b grap b i Ag+ and NO3-ị ở
Basic.concepts.of.electrochemistry
• Galvanic.cell:.separation.of.two.half.reactions.by.the.
use.of.a.salt.bridge
16
Porous.glass.disks.that.allow.ions.
to.diffuse.but.minimize.mixing.of.
solutions.inside.and.outside.the.
bridge.
Ions.migrate.in.and.out.to.offset.
the.charge.buildup
Basic.concepts.of.electrochemistry
• Cell.notation:.line.diagram.is.used.to.describe.
electrochemical.cells
17
Standard.reduction.potential
•
Standard.reduction.potential,.E
o
’.for.each.half.is.measured.
using.the.standard.hydrogen.electrode.(S.H.E.)
•
Example:. Ag
+
(aq).+.e
-
↔ Ag.(s);.E
o
’=0.799.V
à
Set.the.potential.of.the.S.H.E.at.25
o
C.=.0
à A.standard.reduction.potential.is.really.a.potential.difference'between.the.
potential.of.the.reaction.of.interest.and.the.potential.of.S.H.E.(set.to.0)
18
Standard.reduction.potential
• A.positive.standard.reduction.potential.means.electron.flows.
from.the.S.H.E..
–
Example:. Ag
+
(aq).+.e
-
↔ Ag.(s);.E
o
’=0.799.V.:.e
-
flow.from.the.Pt
electrode.to.the.Ag.electrode
• A.negative.standard.reduction.potential.means.electron.flows.
to.the.S.H.E..
–
Example:.Cd
2+
(aq).+.2e
-
↔ Cd.(s);.E
o
’=.-0.402.V.:.e
-
flow.from.the.Cd.
electrode.to.the.Ptelectrode
19
Standard.reduction.potential
20
Standard.reduction.potential
•
Relative.electrode.potential,.E
o
(or.reduction.potential).
describes.the.haft.reactionwritten.as.reduction.
21
Cell.potential
• Cell.potential.E
cell
:.is.the.potential.difference.between.two.
electrodes.of.the.cell.
• Cell.potential.is.related.to.free.energy.of.the.reaction:
∆𝐺 =−𝑛𝐹𝐸
()**
At.standard.condition:
∆𝐺
+
=−𝑛𝐹𝐸
+
()**
=−𝑅𝑇𝑙𝑛𝐾
)0
22
Cell.potential
• The.signs.of.the.electrodes.are.assigned.as.a.Galvanic.cell.is.
developed
∆𝐺 =−𝑛𝐹𝐸
()**
<.0
à 𝐸
()**
>.0
• Cell.potential.E
cell
:.is.the.potential.difference.between.two.
electrodes.of.the.cell.
𝐸
()**
=𝐸
123456(8)
−𝐸
*):5(;)
à𝐸
123456(8)
>.𝐸
*):5(;)
(+).cathode:.reduction
(-).anode:.oxidation
23
= E - E
cathode anode
not at any condition
E
o
= E - E
o
cathode
o
anode
at standard condition
Cell.potential
• Example:.
Cathode.(+):.Ag
+
( Eaq).+e
-
↔ Ag(s);.
o
=.0.799.(V)
Anode.(
-):.....Cu.(s).↔ Cu
2+.
(aq).+.2e
-
;.E
o
=.0.399.(V)
Overall.reaction:.2
Ag
+
(aq).+.Cu.(s).↔ 2.Ag(s).+.Cu
2+.
(aq).
à Measured6𝐸
()**
= 0.400.V
à Reaction.is.spontaneous!
24
Right.(+)Left.(-)
Cell.potential
• Example:.
Cathode.(+):.Cu
2+
(aq).+.2e
-
↔ Cu(s);.E
o
=.0.399.(V)
Anode.(
-):.....Ag.(s).↔ Ag
+.
(aq).+.e
-
;.E
o
=.0.799.(V)
Overall.
reaction:.2 Ag.(s)..+.Cu
2+
(aq).↔ 2.Ag
+
(aq aq).+.Cu( ).
à Measured6𝐸
()**
= - 0.400.V
à Reaction.is.not.spontaneous!
However,.
2 Ag
+
(aq)..+.Cu.(s).↔ 2.Ag(s).+.Cu
2+.
(aq).is.always.
spontaneous.no.matter.how.the.cell.diagram.is.written!
25
Right.(+)Left.(-)
Nernst.equation.for.half.reaction.
• Consider.the.half.reaction:
– The.electrode.potential:
26
At.T.(K)
At.298.(K)
Solute'A'
à
[A]'='molar'
concentration
Gas'A
à
[A]'='p
A
partial'
pressure
Pure'liquid,'solid,'or'solvent,'
no'term'for'A
Nernst.equation.for.half.reaction.
• Multiply.a.half.reaction.by.a.number:
–
E
o
does.not.change
– n.and.the.log.term.do.change
27
Nernst.equation.for.a.complete.reaction.
• For.a.complete.reaction:
𝐸 =𝐸
()**
=𝐸
6(8)
−𝐸
6(;)
• Procedure.for.writing.a.net.cell.reaction.and.finding.its.
potential:
–
Write.reduction'half-reactions.for.both.half-cells,.look.up.for.E
o
values,.
then.balance.the.number.of.electron.in.both.reactions.
– Write.the.Nernst.equation.for.the.right.half.cell,.E
(+)
,.for.the.left.half.
cell.E
(-)
.
– Calculate.E
cell
– Write.the.net.cell.reaction.by.subtracting.the.left.half-reaction.from.
the.right.half.one.
28
Potential.of.the.electrode.
attached.to.the.+.of.the.voltmeter
Potential.of.the.electrode.
attached.to.the.- of.the.voltmeter
<0: reverse to the direction
Nernst.equation.for.a.complete.reaction.
29
Outlines
• Introduction.to.Electrochemistry
• Applications.of.Standard.electrode.potentials
• Redox.titration
30
Reference.electrode
• Set.up.a.Galvanic.cell.to.measure.the.relative.amount.of.
Fe
2+
/Fe
3+.
31
Half.reactions.is.written.as.reductions:
Nernst.equations.for.half.reactions:
Cell.potential:
fixed
Depends'on'[Fe
2+
]/[Fe
3+
]
The.half-cell.on.the.left.acts.
as.a.reference'electrodeà
provides.a.constant'
potential'to.the.left.side.of.
the.potentiometer.
Reference.electrode
• Set.up.a.Galvanic.cell.to.measure.the.relative.amount.of.
Fe
2+
/Fe
3+.
32
The.half-cell.on.the.left.
acts.as.a.reference'
electrode.
Indicator.
electrode
The.Ptwire.is.the.indicator.electrode,.whose.
potential.responds.to.[Fe
2+
]/[Fe
3+
].
à Inert.metals.used.as.indicator.electrodes
Reference.electrode
• A.standard.hydrogen.electrode.is.difficult.to.use.because.H
2
gas.and.a.freshly.prepared.catalytic.Ptsurface.are.required..
• Common.reference.electrodes:
Ag/AgCl
(Silver-silver.chloride)
Calomel.electrode
Reference.electrode
• Voltage.conversions.between.different.reference.scales
• Example:
– If.an.electrode.(A).has.a.potential.of.- 0.461.V.with.respect.to.a.S.C.E,.
what.is.the.potential.with.respect.to.S.H.E.electrode?.If.B,.whose.
potential.is.0.033.V.from.silver-silver.chloride
34
standard calomel electron
Calculation.of.potentials.of.electrochemical.cells
• The potential of an electrochemical cell is the difference
between the elctrode potential of the cathode and anode:
𝐸 =𝐸
()**
=𝐸
6(8)
−𝐸
6(;)
35
Potential.of.the.electrode.
attached.to.the.+.of.the.voltmeter
Potential.of.the.electrode.
attached.to.the.- of.the.voltmeter
Calculation.of.potentials.of.electrochemical.cells
• Procedure.for.writing.a.net.cell.reaction.and.finding.
its.potential:
– Write.reduction'half-reactions.for.both.half-cells,.look.up.
for.
E
o
values,.then.balance.the.number.of.electron.in.both.
reactions.
– Write.the.Nernst.equation.for.the.right.half.cell,.E
(+)
,.for.
the.left.half.cell.E
(-)
.
– Calculate.E
cell
– Write.the.net.cell.reaction.by.subtracting.the.left.half-
reaction.from.the.right.half.one.
36

Tài liệu khác của Trường Đại học Quốc tế, Đại học Quốc gia Thành phố Hồ Chí Minh

Preview text:

Analytical Chemistry 1 Lecture'10' Electrochemistry-and-Redox- titration- Instructor:'Nguyen'ThaoTrang Outlines
• Introduction.to.Electrochemistry
• Applications.of.Standard.electrode.potentials thế điện cực • Redox.titration 2 Outlines
• Introduction.to.Electrochemistry
• Applications.of.Standard.electrode.potentials • Redox.titration 3
Basic.concepts.of.electrochemistry
• Redox.reaction.:.reactions.with.electron.transfer
– oxidized.:.loses. electrons – reduced.:.gains.electrons
– oxidizing.agent.(oxidant).:.takes.electrons.from.others. and.becomes. reduced
– reducing.agent.(reductant). :.gives.electrons. to.another.substance.and. is.oxidized – Example
Fe3+.is.an.oxidizing.agent.as:.Fe3++.e- à Fe2+
V2+.is.an.reducing.agent.as:.V2+ à V3++.e- 4
Basic.concepts.of.electrochemistry • Redox.reaction.:
– Half.reactions:. A.common.approach.for.listing.species.that.undergo.
REDOX.is.by.using.half.reactions. oxidixing agent
Reduction:.Fe3++.e- à Fe2+ reducing agent half rxn (reduction rxn) Oxidation:.V2+ à V3+ - +.e half rxn (oxidation rxn)
à Redox.reaction.is.the.product.of.two.half.reactions
– Each.half.reaction.is.associated.with.a.standard.reduction.potential,. Eo’ Box+.e- → Bred,.Eo’ Ared → Aox+.e-,.Eo’ 5
Basic.concepts.of.electrochemistry • Redox.reaction.:
– Listed.half.reactions.with.standard.reduction.potential: 6
Basic.concepts.of.electrochemistry • Redox.reaction.:
– Example:.Determine.the.balanced.equation.for.the.reaction.of.Fe2+ with.Cr 2- 2O7 in.an.acidic.solution. – Solution: E 0
Cr2O7(2-)/Cr(3+)>E Fe(3+)/Fe(2+)=> Cr2O7 2-gains e more easily than Fe3+
• From.the.standard.reduction.potentials,.we.can.get: Fe3++.e- =.Fe2+ Eo= 0.77 Cr 2- 3+ 2O7 +.14H+ +.6e- =.2Cr +.7H2O Eo= 1.33
• We.know.that.one.of.the.half.reactions.must.be.reversed.since. oxidation.
and.reduction.reactions.must.both.occur
• Based.on.the.question,.Fe2+.must.undergo.oxidation.reaction: Fe2+ ez lose its e
à reverse.reaction.(1):.Fe2+.=.Fe3+ - +.e
– To.balance. the.equation,. the.number.of.electrons. in.both.equation. must.be.
the.same.à multiply.(1.).by.6:
6.x.(1):.....................6Fe2+.=.6Fe3++.6e- +.(2):..Cr 2- + - 3+ 2O7 +.14H +.6e =.2Cr +.7H2O. 6Fe2+.+.Cr2O 2- 7 +.14H+ =.6Fe3++.2Cr3++.7H2O. 7
Basic.concepts.of.electrochemistry • Electric.charge.q:.
– Is.measured.in.[coulomb].or.[C]
– Magnitude.of.the.charge.of.a.single.electron.=.1.602.x.10-19C.
– A.mole.of.electrons.has.a.charge.of.(1.602.x.10-19C)x.(6.022.x 1023
mol-1).=.9.649.x.104 C/mol,.which.is.called.the.Faraday'constant'(F).
– Relationship.between. charge.and.moles:
where.n'is.the.number.moles.of.electrons.transferred. 8
Basic.concepts.of.electrochemistry • Electric.charge.q:. 9
Basic.concepts.of.electrochemistry • Electric.current:
– The.quantity.of.charge.flowing.each.second.through.a.circuit.is.called. the.current.-
– The.unit.of.current.is.the.ampere,-abbreviated.A.
– A.current.of.1.A..represents.a.charge.of.1.C.per.second.flowing.past.a. point.in.a.circuit.
FIGURE-13-1-Electrons.flowing.into.a.
coil.of.Ptwire.at.which.Sn4+.ions.in. solution.are.reduced.to.Sn2+ . 10
Basic.concepts.of.electrochemistry
• Voltage,.work,.free.energy:
– Difference.in.electric'potential,.E.(measured.by.volt).between.two.
points.is.the.work.needed.when.moving.an.electric.charge.from.one. point.to.the.other: – Work:..Energy.in.joules • one.joule.of.energy.= .between.points.
whose.potential.differ.by.one.volt.
– Example:.How.much.work.is.required.to.move.2.36.mmolof.electrons.
through.a.potential.difference.of.1.05.V.? 2.39 x 102 J –
Work.done.on.surroundings..=.- ΔG 11
à ΔG.=.-E.q =.- nEF < 0=> E>0 => Del E (potential diff) >0
Basic.concepts.of.electrochemistry
• Types.of.electrochemical.cell: 1. Galvanic-cell-(Voltaic):-
A.cell.generates.electrical.energy.from.a.spontaneous. chemical.reaction. 2. Electrolytic-cell:
Is.where.non-spontaneous.redox.reactions.can.be.made.to.
occur.by.the.addition.of.electrical.energy.(process.is.termed. electrolysis). 12
Basic.concepts.of.electrochemistry
• Common.components.of.a.electrochemical.cell
– Electrodes:.conduct.electricity.between.cell.and. surroundings • Anode:.oxidation.occurs. • Cathode:.reduction.occurs.
– Electrolytes:.mixture.of.ions.involved.in.reaction.or. carrying.charge
– Salt.bridge:.completes.circuit.(provides.charge.balance) 13
Basic.concepts.of.electrochemistry • Galvanic.cell:
– A.simple.Galvanic.cell:.2.electrodes.in.a.solution.of.CdCl2:
One.electrode.is.cadmium;.the.other.is.metallic.silver. coated.with.solid.AgCl.E 0
AgCl/Ag>E Cd(2+)/Cd=> reverse the rxn at the anode
Anode:..... Cd.(s).↔ Cd2+(aq) +.2e-Eo= - 0.44
Cathode:...2AgCl.(s).+.2e- ↔ 2Ag(s).+.2Cl-(aq).
Cd.(s).+.2AgCl.(s).↔ Cd2+(aq) +.2Ag(s).+.2Cl-(aq) ΔG.=.-150.kJ/molCd
à Oxidation.of.Cd.electrode.provides.
electrons. that.flow.through.the.circuit.to. the.Ag.electrode.
AgCl(s)+e- ↔ Ag(s)+Cl-(aq) E0= 0.22 14
Basic.concepts.of.electrochemistry • Galvanic.cell:
– A.simple.Galvanic.cell:.2.electrodes.Cd.and.Ag.in.a.solution. of.Cd(NO3)2 and.AgNO3.: Not.really. work!
A.little.current.flows.through.the.circuit.as.Ag.
can.react.at.the.Cd.electrode..
the rxn still occur but no flow e e bị grap bởi Ag+ and NO3- 15
Basic.concepts.of.electrochemistry
• Galvanic.cell:.separation.of.two.half.reactions.by.the. use.of.a.salt.bridge
Porous.glass.disks.that.allow.ions.
to.diffuse.but.minimize.mixing.of.
solutions.inside.and.outside.the. bridge.
Ions.migrate.in.and.out.to.offset. the.charge.buildup 16
Basic.concepts.of.electrochemistry
• Cell.notation:.line.diagram.is.used.to.describe. electrochemical.cells 17 Standard.reduction.potential
• Standard.reduction.potential,.Eo’.for.each.half.is.measured.
using.the.standard.hydrogen.electrode.(S.H.E.)
• Example:. Ag+(aq).+.e- ↔ Ag.(s);.Eo’=0.799.V
à Set.the.potential.of.the.S.H.E.at.25oC.=.0
à A.standard.reduction.potential.is.really.a.potential.difference'between.the.
potential.of.the.reaction.of.interest.and.the.potential.of.S.H.E.(set.to.0) 18 Standard.reduction.potential
• A.positive.standard.reduction.potential.means.electron.flows. from.the.S.H.E..
– Example:. Ag+(aq).+.e- ↔ Ag.(s);.Eo’=0.799.V.:.e- flow.from.the.Pt electrode.to.the.Ag.electrode
• A.negative.standard.reduction.potential.means.electron.flows. to.the.S.H.E..
– Example:.Cd2+(aq).+.2e- ↔ Cd.(s);.Eo’=.-0.402.V.:.e- flow.from.the.Cd. electrode.to.the.Ptelectrode 19 Standard.reduction.potential 20 Standard.reduction.potential
• Relative.electrode.potential,.Eo (or.reduction.potential).
describes.the.haft.reactionwritten.as.reduction. 21 Cell.potential
• Cell.potential.Ecell:.is.the.potential.difference.between.two. electrodes.of.the.cell.
• Cell.potential.is.related.to.free.energy.of.the.reaction: ∆𝐺 =−𝑛𝐹𝐸()** At.standard.condition:
∆𝐺+ =−𝑛𝐹𝐸+()**=−𝑅𝑇𝑙𝑛𝐾)0 22 Cell.potential
• The.signs.of.the.electrodes.are.assigned.as.a.Galvanic.cell.is. developed
∆𝐺 =−𝑛𝐹𝐸()**<.0 à 𝐸()**>.0
• Cell.potential.Ecell:.is.the.potential.difference.between.two. electrodes.of.the.cell. Eo = Eo o cathode - E anode at standard condition
𝐸()**=𝐸123456(8)−𝐸*):5(;) = Ecathode - Eanode
à𝐸123456(8)>.𝐸*):5(;) not at any condition (+).cathode:.reduction (-).anode:.oxidation 23 Cell.potential • Example:. Left.(-) Right.(+)
Cathode.(+):.Ag+ (aq).+e- ↔ Ag(s);.Eo =.0.799.(V)
Anode.(-):... .Cu.(s).↔ Cu2+.(aq).+.2e- ;.Eo =.0.399.(V)
Overall.reaction:.2Ag+ (aq).+.Cu.(s).↔ 2.Ag(s).+.Cu2+.(aq). à Measured6𝐸()**= 0.400.V à Reaction.is.spontaneous! 24 Cell.potential • Example:. Left.(-) Right.(+)
Cathode.(+):.Cu2+ (aq).+.2e- ↔ Cu(s);.Eo =.0.399.(V)
Anode.(-):... .Ag.(s).↔ Ag+.(aq).+.e- ;.Eo =.0.799.(V)
Overall.reaction:.2 Ag.(s)..+.Cu2+ (aq).↔ 2.Ag+(aq).+.Cu(aq).
à Measured6𝐸()**= - 0.400.V
à Reaction.is.not.spontaneous!
However,.2 Ag+ (aq)..+.Cu.(s).↔ 2.Ag(s).+.Cu2+.(aq).is.always.
spontaneous.no.matter.how.the.cell.diagram.is.written! 25
Nernst.equation.for.half.reaction.
• Consider.the.half.reaction: – The.electrode.potential: Solute'A'à [A]'='molar' concentration At.T.(K) Gas'Aà [A]'='pA partial' pressure At.298.(K)
Pure'liquid,'solid,'or'solvent,' no'term'for'A 26
Nernst.equation.for.half.reaction.
• Multiply.a.half.reaction.by.a.number: – Eo does.not.change
– n.and.the.log.term.do.change 27
Nernst.equation.for.a.complete.reaction. • For.a.complete.reaction:
𝐸 =𝐸()**=𝐸6(8)−𝐸6(;)
<0: reverse to the direction Potential.of.the.electrode. Potential.of.the.electrode.
attached.to.the.+.of.the.voltmeter
attached.to.the.- of.the.voltmeter
• Procedure.for.writing.a.net.cell.reaction.and.finding.its. potential:
– Write.reduction'half-reactions.for.both.half-cells,.look.up.for.Eo values,.
then.balance.the.number.of.electron.in.both.reactions.
– Write.the.Nernst.equation.for.the.right.half.cell,.E(+),.for.the.left.half. cell.E(-). – Calculate.Ecell
– Write.the.net.cell.reaction.by.subtracting.the.left.half-reaction.from. the.right.half.one. 28
Nernst.equation.for.a.complete.reaction. 29 Outlines
• Introduction.to.Electrochemistry
• Applications.of.Standard.electrode.potentials • Redox.titration 30 Reference.electrode
• Set.up.a.Galvanic.cell.to.measure.the.relative.amount.of. Fe2+/Fe3+.
Half.reactions.is.written.as.reductions:
Nernst.equations.for.half.reactions: Cell.potential:
The.half-cell.on.the.left.acts. as.a.reference'electrodeà Depends'on'[Fe2+]/[Fe3+] fixed provides.a.constant' potential'to.the.left.side.of. 31 the.potentiometer. Reference.electrode
• Set.up.a.Galvanic.cell.to.measure.the.relative.amount.of. Fe2+/Fe3+. Indicator. electrode
The.Ptwire.is.the.indicator.electrode,.whose.
potential.responds.to.[Fe2+]/[Fe3+]. The.half-cell.on.the.left. acts.as.a.reference'
à Inert.metals.used.as.indicator.electrodes electrode. 32 Reference.electrode
• A.standard.hydrogen.electrode.is.difficult.to.use.because.H2
gas.and.a.freshly.prepared.catalytic.Ptsurface.are.required..
• Common.reference.electrodes:
Ag/AgCl(Silver-silver.chloride) Calomel.electrode Reference.electrode
• Voltage.conversions.between.different.reference.scales • Example: standard calomel electron
– If.an.electrode.(A).has.a.potential.of.- 0.461.V.with.respect.to.a.S.C.E,.
what.is.the.potential.with.respect.to.S.H.E.electrode?.If.B,.whose.
potential.is.0.033.V.from.silver-silver.chloride 34
Calculation.of.potentials.of.electrochemical.cells
• The potential of an electrochemical cell is the difference
between the elctrode potential of the cathode and anode:
𝐸 =𝐸()**=𝐸6(8)−𝐸6(;) Potential.of.the.electrode. Potential.of.the.electrode.
attached.to.the.+.of.the.voltmeter
attached.to.the.- of.the.voltmeter 35
Calculation.of.potentials.of.electrochemical.cells
• Procedure.for.writing.a.net.cell.reaction.and.finding. its.potential:
– Write.reduction'half-reactions.for.both.half-cells,.look.up.
for.Eo values,.then.balance.the.number.of.electron.in.both. reactions.
– Write.the.Nernst.equation.for.the.right.half.cell,.E(+),.for. the.left.half.cell.E(-). – Calculate.Ecell
– Write.the.net.cell.reaction.by.subtracting.the.left.half-
reaction.from.the.right.half.one. 36

Tài liệu liên quan: