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KİM 683
Electroanalytical
Chemistry
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Impedance ?
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Who won the Nobel Prize in Chemistry in 2019?
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Conversion and storage of energy
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Solar cells, DSSC, Photoelectrochemical Cells etc
315M421 Kuantum Parçacık Duyarlı Güneş Hücreleri İçin Yeni Nesil
Nanokompozit Elektrotlar
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111T488 Nanoboyutlu metal oksit fotoaktif elektrotların elektrokimyasal sentezi ve
fotoakım performanslarının belirlenmesi (1001)
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Electroplating
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Electrochemical synthesis of some industrial
materials eg. Al, Cl2, Na,..
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Organic compounds: adiponitrile
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Electroanalytical chemistry (e.g. the analysis of chemicals in blood to
determine the development of a certain diseases)
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Obtaining thermodynamic data about a reaction
Waste water purification and recycling
Corrosion protection and so on ………………….
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Why electrochemistry?
• Clean and green (not always true!)
• Often no or few byproducts
• Electrons the cheapest redox “reagents”
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What is a redox reaction?
Oxidation is the addition of oxygen to a substance and
Reduction is the removal of oxygen from a substance.
oxygen removed
reduction
oxygen added
oxidation
Reduction and oxidation always take place together.
Why is this type of reaction called a redox reaction?
redox = reduction and oxidation
Which substances are oxidized and reduced in this reaction?
lead oxide + carbon lead
carbon monoxide+
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Redox and electronsMagnesium burns in oxygen
to form magnesium oxide.
A redox reaction can also
be explained in terms of the
gain or loss of electrons.
What happens to the atoms and electrons in this reaction?
It is obvious that the
magnesium has been
oxidized, but what has
happened to the oxygen?
magnesium + oxygen magnesium oxide
2Mg(s) O2(g) 2MgO(s)+
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Oxidation and electron loss
When magnesium burns in oxygen to form magnesium oxide, what happens to
magnesium and its electrons?
Oxidation is the loss of electrons.
oxidized
(electrons lost)
⚫ The magnesium has been oxidized.
⚫ The Mg atom has lost 2 electrons to form a Mg2+ ion.
Mg Mg2+ O2-O+
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Reduction and electron gain
When magnesium burns in oxygen to form magnesium oxide,
what happens to oxygen and its electrons?
Reduction is the gain of electrons.
reduced
(electrons gained)
Mg Mg2+ O2-O+
⚫ The oxygen has been reduced.
⚫ The O atom has gained 2 electrons to form a O2- ion.
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An electrochemical system is not homogeneous but is heterogeneous.
Conductionoccurs via
migration ofelectrons .Solid state
Physics : energyband theory.
Material transport occursvia migration, diffusionand convection
Electronically conducting phase : metal, semiconductor,Conducting polymer material etc.
Ionically conductingmedium : electrolytesolution, molten salt,solid electrolyte,polymericelectrolyte, etc.
Electrode/Electrolyte İnterface
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Representation of reduction and oxidation process of a species
A (molecule) in solution
Reduction
Oxidation
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a) Potentials for possible reductions at a platinum electrode, initially at ~ 1 V vs.
NHE in a solution of 0.01 M each of Fe3+, Sn4+, and Ni2+ in 1M HCl (b) Potentials for
possible oxidation reactions at a gold electrode, initially at ~0.1V vs. NHE in a
solution of 0.01 M each of Sn2+ and Fe2+ in 1M HI.
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Potentials for possible reductions at a mercury electrode in 0.01 M Cr3+ and Zn2+ in
1M HCl. The arrows indicate the directions of potential change discussed in the
text.
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Current is always in
the opposite
directions from
electron flow.
The flow of
positive charge.
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Electrochemical Cells
Galvanic
Electrolytic
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Electrochemical processes are oxidation-reduction
reactions in which:
• The energy released by a spontaneous reaction
is converted to electricity or
• Electrical energy is used to cause a
nonspontaneous reaction to occur
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A galvanic electrochemical cell at open circuit.
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A galvanic cell doing work.
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Zn(s) | Zn2+(aq) ¦¦ Cu2+(aq) | Cu(s)
Daniel Cell
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ZnSO4 solution CuSO4 solution
Zn2+
Cu2+SO42–
SO42–
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Zincanode
Coppercathode
ZnSO4 solution CuSO4 solution
Zn2+
Cu2+
Salt bridge
e– e–
Cl– K+
Voltmeter
CottonplugsSO4
2–
SO42–
Salt bridge provides electrical neutrality by providing
negative anions to equal the positive cations being
created at the Zn anode during oxidation. And cations
ions (K+) to replace Cu2+ being used up at reduction.
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Zincanode
Coppercathode
ZnSO4 solution CuSO4 solution
Zn2+
Cu2+
Salt bridge
Zn
e– e–
Cl– K+
Voltmeter
CottonplugsSO4
2–
SO42–
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Zincanode
Coppercathode
ZnSO4 solution CuSO4 solution
Zn2+
Cu2+
Salt bridge
Zn is oxidizedto Zn2+ at anode.
Zn(s) Zn2+(aq) + 2e–
Zn Zn2+
2e–
e– e–
Cl– K+
Voltmeter
CottonplugsSO4
2–
SO42–
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Zincanode
Coppercathode
ZnSO4 solution CuSO4 solution
Zn2+
Cu2+
Salt bridge
Cu
Zn is oxidizedto Zn2+ at anode.
Zn(s) Zn2+(aq) + 2e–
Zn Zn2+
2e–
e– e–
Cl– K+
Voltmeter
CottonplugsSO4
2–
SO42–
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Zincanode
Coppercathode
ZnSO4 solution CuSO4 solution
Zn2+
Cu2+
Salt bridge
2e–
Cu2+
Cu
Zn is oxidizedto Zn2+ at anode.
Zn(s) Zn2+(aq) + 2e–
Zn Zn2+
2e–
e– e–
Cl– K+
Voltmeter
CottonplugsSO4
2–
SO42–
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Zincanode
Coppercathode
ZnSO4 solution CuSO4 solution
Zn2+
Cu2+
Salt bridge
2e–
Cu2+
Cu
2e– + Cu2+(aq) Cu(s)
Cu2+ is reducedto Cu at cathode.
Zn is oxidizedto Zn2+ at anode.
Zn(s) Zn2+(aq) + 2e–
Zn Zn2+
2e–
e– e–
Cl– K+
Voltmeter
CottonplugsSO4
2–
SO42–
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Zincanode
Coppercathode
ZnSO4 solution CuSO4 solution
Zn2+
Cu2+
Salt bridge
Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)
Net reaction
2e–
Cu2+
Cu
2e– + Cu2+(aq) Cu(s)
Cu2+ is reducedto Cu at cathode.
Zn is oxidizedto Zn2+ at anode.
Zn(s) Zn2+(aq) + 2e–
Zn Zn2+
2e–
e– e–
Cl– K+
Voltmeter
CottonplugsSO4
2–
SO42–
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Galvanic Cells
spontaneous
redox reaction
anode
oxidation
cathode
reduction
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Zincanode
Coppercathode
ZnSO4 solution CuSO4 solution
Zn2+ SO42–
Cu2+
Salt bridge
Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)
Net reaction
2e–
Cu2+
Cu
2e– + Cu2+(aq) Cu(s)
Cu2+ is reducedto Cu at cathode.
Zn is oxidizedto Zn2+ at anode.
Zn(s) Zn2+(aq) + 2e–
Zn Zn2+
2e–
e– e–
Cl– K+
Voltmeter
CottonplugsSO4
2–
Galvanic Cell
Flow of e- Flow of current
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Cell notation
Zn(s) Zn2+(aq) Cu2+(aq) Cu(s)
phase difference phase differenceIndicated by solid vertical line Indicated by solid vertical line
anode components cathode components
salt bridge
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An electrolytic cell.
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Electrolytic Cells
• Electrolytic Cell – a cell in which a nonspontaneous redox reaction is forced to occur; a combination of two electrodes, an electrolyte and an external power source.
– Electrolysis – the process of supplying electrical energy to force a nonspontaneous redox reaction to occur
– The external power source acts as an “electron pump”; the electric energy is used to do work on the electrons to cause an electron transfer
Electrons are pulled from the anode and pushed to
the cathode by the battery or power supply
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Comparing Electrochemical Cells:
Voltaic and Electrolytic
It is best to think of “positive” and “negative” for electrodes as labels, not charges.