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Electrochemistry
Electrochemistry
Electrochemistry is a branch of chemistry that studies chemical reactions
which take place in a solutionat the interface of an electron conductor(the
electrode: a metal) and an ionic conductor (the electrolyte), and which
involve electron transfer between the electrode and the electrolyte.
It deals with the chemical reactions produced by passing electric current or
the production of electric current through chemical reactions.
Electrochemical Processes The chemical reactions involving electricity are called electrochemical
reaction
Electrochemical reactions are redox reactions where an oxidation and a
reductionreaction go side by side and are separated in space.
The reactions are two types
1. Induced electrolytic reactions- these non spontaneous reactions are
forced by the passage of electricity through the reactants
2. spontaneous redox reaction that can produce electricity
REDOX REACTIONS
Oxidation-reduction or redox reactionstake place in electrochemical cells.
The term redox comes from the two concepts of reduction and oxidation.
Oxidationdescribes the lossof electrons:
Zn (s) Zn2+
(aq) + 2 e -
Reductiondescribes the gain of electrons
Cu2+
(aq) + 2 e- Cu (s)
Each of the reaction is known as half- reaction or half-cell and both
must always go side by side
The net reaction: Zn (s) + Cu2+
(aq) Zn2+
(aq) + Cu (s)
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Electrochemical Cells
An electrochemical cell is a device capable of either deriving electricalenergy
from chemical reactions, or facilitating chemical reactions through the
introduction of electrical energy.
There are two types of electrochemical cells.
TYPES:
1. Electrochemical cell and Electrolytic Cells
2.Galvanic cell and concentration cell
3. Reversible and irreversible cell
1. Spontaneous reactions occur in galvanic (voltaic) cells;
2. Nonspontaneous reactions occur in electrolytic cells.
Both types of cells contain electrodes where the oxidation and reduction
reactions occur.
Oxidation occurs at the electrode is termed as the anode and reductionoccurs at the electrode is called as the cathode.
1. Galvanic or Voltaic cell: Produces energy by a spontaneous reaction which
produces electricity as a result of electron transfer
The cells used for the generation of electrical energy from chemical
reactions are called galvanic or voltaic cells.
The redox reaction in a galvanic cell is a spontaneous reaction.
For this reason, galvanic cells are commonly used as batteries. Galvanic
cell reactions supply energy which is used to perform work. The energy
is harnessed by situating the oxidation and reduction reactions in separate
containers, joined by an apparatus that allows electrons to flow.
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Ex. A common galvanic cell is the Daniell cell, batteries, corrosion, etc
2. Electrolytic Cells The redox reaction in an electrolytic cell is non-
spontaneous. Electrical energy is required to induce the electrolysis reaction.
The cells used for electrolysis are called electrolytic cells
ex. Charging of battery, electroplating ,etc
Electrodes & Charge
The anode of a galvanic cell is negatively charged, since the spontaneous
oxidation at the anode is the sourceof the cell's electrons or negative charge.
The cathode of a galvanic cell is its positive terminal.
The anodeof an electrolytic cell is positive (cathode is negative), since theanode attracts anions from the solution.
In both galvanic and electrolytic cells, oxidation takes place at the anode and
electrons flow from the anode to the cathode.
An electrochemical cell is obtained by coupling two half cells
1. Anode: oxidation half-cell reaction takes place
2.
Cathode: reduction half-cell reaction occurs
3. Salt bridge: A salt bridge is often employed to provide ionic contact
between two half-cells with different electrolytesto prevent the solutions
from mixing and causing unwanted side reactions.
Ex.: filter paper soaked in KNO3
Functions of salt bridge:
1. The salt bridge allows ions to move on either side and maintain the
electrical neutrality of the electrolyte on both sides.
2. It serves as a bridge to complete the electric circuit.
3.
It prevents the liquid junction potential between the two electrodes.
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Liquid junction potential occurs when two solutions of different
concentrations are in contact with each other. The more concentrated
solution will have a tendency to diffuse into the comparatively less
concentrated one. The rate of diffusion of each ion will be roughly
proportional to its speed in an electric field. If anion diffuses more
rapidly than the cation, it will diffuse ahead into the dilute solution
leaving the later negatively charged and the concentrated solution
positively charged. So an electrical double layer of positive and negative
charges will be produced at the junction of the two solutions. So at the
point of junction, a difference of potential will develop because of theionic transfer. This potential is called liquid junction potential or
diffusion potential. The magnitude of the potential depends on the
relative speeds of the ions.
4.
External circuit: These two half-cells joined together by wire through
which electrons flow.
5. Electrolyte: Internal pathway that allows ions to migrate between them so
as to preserve electro neutrality.
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Difference in electrolytic cell and galvanic cell
Electrolytic cell Galvanic cell
1. Conversion of electrical energy
into chemical energy
1. Chemical energy into electrical
energy
2. The anode carries positive
charge vice versa.
2. The anode carries negative
charge Vice versa.
3. Electrons are supplied to the cell
from the external power supply.
3. Electrons are drawn from the
cell.
4. Not a spontaneous reaction. Eo
cellis -ve, then the process is
nonspontaneous. E.g electroplating
4. Spontaneous reaction. Eo
cell is+ve, then the process is
spontaneous. eg. Corrosion
5. The extent of chemical reaction
occurring at the electrode depends
on the quantity of electricity passed
& is governed by Faradays law of
electrolysis.
5. The e.m.f of the cell depends
on the concentration of the
electrolyte and chemical
nature of the electrode (Nernst
Equation)
6. The amount of electricity passed
during electrolysis is measured
by Coulometer.
6. The e.m.f produced in the cell
is measured by potentiometer.
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ELECTRODE POTENTIAL (E)
When a metal (M) is placed in the solution of its own ion it may act as an
anode or cathode.
As an anode: Positive metal ions passing into solution (oxidation) leaving
behind e-s on the electode.
The anode attains negative charge due to accumulation of e-s which attracts
positively charged free ions (cations) from the solution. Due to the
attraction the positive ions remain close to the metal.
M Mn++ne-
As a cathode: Positive ions depositing on the metal electrode (reduction)
and it attracts negatively charged ions.
Mn+
+ne-----M
HELMHOLTZ ELECTRICAL DOUBLE LAYER
Thus a sort of electrical double layer (positive or negative ions) is formed
all around the metal. This layer is called Helmholtz electrical double layer.
This layer prevents further passing of or deposition of metal ions on the
metal. Consequently a difference in potential (Galvanic potential) set up
between the metal and its solution.
At equilibrium, the potential difference is constant, which is known as
Electrode potential of a metal.EMF: The diff. of potential between the two electrodes in a voltaic cell which
causes flow of current/electrons is called the electromotive force
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Electrode potential of the metal shows its tendency to under go loss or gain of
electrons
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Electrode potential (E):
The measure of tendency of an electrode to lose or gain electrons when it is
contact with a solution of its own ion is called electrode potential.
Standard electrode potential (E
o
):The measure of tendency of an electrode to lose or gain electrons, when it is
contact with a solution of its own salt of 1 molar concentration at 25oC is known
as standard electrode potential.
Oxidation electrode potential (Eoxid):
The measure of tendency of an electrode to lose electrons when it is contact with a
solution of its own ion is called oxidation electrode potential.
Reduction electrode potential (E red):The measure of tendency of an electrode to
gain electrons when it is contact with a solution of its own ion is called electrode
reduction potential.
Factors affecting electrode potential
1. Nature of the electrode metal
2. Temperature
3. Concentration of metal ions in solution
Measurement of Single electrode potential
Impossible to know the absolute value of single electrode potential.
To measure the ele. pot of one electrodes which is connected to reference
electrode to form a complete cell
The electrode whose electrode pot. is either exactly known or arbitrarily
fixed is called reference electrode
The pot. of one electrode is fixed arbitrarily then another electrode pot. can
be measured
EMF of the complete cell can be directly read from the potentiometer.
Ecell= Ecathode- E
anode
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Potential of the single electrode can be calculated by
Nernst Equation:
The final and the most fundamental form the Nernst equation is written as:E = E
o RT/nF ln ap/ar
where R is the universal gas constant, T is the absolute temperature in degrees
Kelvin, z is the charge number of the electrode reaction (which is the number of
moles of electrons involved in the reaction as written), and F is the Faraday
constant (96,500 C mole-1
). The notation aprepresents the chemical activitiesof
all of the species which appear on the product side of the electrode reaction and
the notation arrepresents the chemical activities of all of the species which appear
on the reactant side of the electrode reaction.
Derivation
Nernst equation was derived from Vant Hoff reaction isotherm for the
redox reaction, M2+
+ 2e---M
Vant Hoff Reaction Isotherm
G = - RT ln K +RT ln Q
Vant Hoff reaction isotherm is an equation for the change in free energy
during the chemical reaction.
it relates the G (free energy change) and K (equilibrium constant) for
the redox reaction as
G = - RT ln K +RT ln Q ---------------(1)
Q (Reaction Quotient) = aC. aD/ aA . aB = [P] /[R]
aC. aD = chemical activity of products
aA . aB = chemical activity of reactants
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The activities of gases are usually taken as their partial pressures and
the activities of solutes such as ions are usually taken as their molar
concentrations.
G = - RT ln K +RT ln [P] /[R] ---------------(1)
At standard conditions of T and P, at equilibrium [P] = [R]
Go = - RT ln K
So equation 1 becomes
G = Go +RT ln [P] /[R] ---------------------(2)
Electrical energy (nFE) arises from the expense of free energy of the system
(- G).
Let n faraday charge be taken out of a electrode or cell of emf E; then
work done by the cell will be calculated as:
Work = Charge x Potential = nFE
Work done by the cell is equal to decrease in free energy.
-G = nFE
-Go= nFE
o
G = Go+ RT ln [product] / [reactant]
-nFE = -nFEO + 2.303 RT log [product] / [reactant]
Ecell= Eocell
- 2.303 RT / nF log [P]/[R]
R = 8.314 J/K/mole
At 25oC, T =298 K
F= 96500 Coulombs
The oxidation potential of an electrode for the reaction
M----Mn+
+ ne-
Eoxid = Eooxid- 0.0592/n log [Mn+] / [M] ----- half cell
[M] = 1
Eoxid = Eo
oxid- 0.0592/n log [Mn+]
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ELECTROCTREMICAL SERIES (e.m.f series): The arrangement of metals in the
increasing order of their std. reduction electrode potential
Half-Reaction E0
Li++ e- Li
Na+ + e- Na
Mg2+ + e- Mg
Al3+
+ 3e- Al
Ti2+
+ 2e- Ti
Mn2+ +e- Mn
Zn2+
+ 2e- ZnFe
2++ 2e- Fe
Co2+
+ 2e- Co
Ni2+
+ 2e- Ni
Fe3+
+ 3e- Fe
2H++ 2e- H2 (g)
Sn4+
+ 2e- Sn2+
Cu2++ 2e- Cu
Fe3+
+ e- Fe2+
Ag++ e- Ag
Pt4+
+ e- Pt
Mn4+
+ 2e- Mn
2+
Cr6+
+ 3e- Cr
3+
Au+ + e- Au
Mn7++ 5e- Mn2+
Cr4+
+ 1e- Cr3+
Au3+
+ 3e- Au
F2+ 2e- 2F-
-3.05
-2.70
-2.40
-1.66
-1.63
-1.18
-0.76-0.44
-0.28
-0.25
-0.04
0.00
+.15
+0.34
+0.77
+0.80
+0.86
+1.23
+1.33
+1.50
+1.51
+1.60
+1.69
+2.87
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Application of electrochemical series
1.
Relative ease of oxidation or reduction can be predicted
The metals on the top with ve reduction pot. Can more easily undergooxidation & act as anode.
Ex. Zn = -0.76 V favours oxidation reaction
While the metals at the bottom with +ve reduction. Pot. has great
tendency to undergo reduction & act as cathode
Ex. Cu = +0.34 V favours reduction reaction.
2. Cell representation can be predicted
A galvanic cell formed by two half cells with 2 diff. metals.
From EMF series, the metal which is undergoing oxidation or reduction can
be predicted
The electrode undergoes oxidation (anode) is written at left
The electrode undergoes reduction (cathode) is written at right
Ex. For Danial cell with Zn and Cu electrodes
The cell representation
Zn / Zn2+
// Cu2+
/ Cu
The cell reaction can also be written
Anodic reaction:
Zn (s) Zn2+
(aq) + 2 e
Cathodic reaction:
Cu2+
(aq) + 2 e- Cu (s)
The net reaction:
Zn (s) + Cu2+
(aq) Zn2+
(aq) + Cu (s)
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3. Calculation of Std emf (Eocell) of the cell
Oxidn Half cell|| Redn half cell
Eocell= E
oright - E
oleft
Or E
ocell= E
ocathode (red.) E
oanode (red.)
Or Eocell= E
oref.) E
ounknown
Ex.
Std e.m.f (Eocell) of (Danial cell) Zn-Cu cell
Eocell= EoCu EoZn= +0.34 (-0.76) V = +1.1 V
4. Calculation of std. free energy (Go)
Go = - nFEo
and equilibrium constant (Keqm)
Go= - RT ln Keqm
- nFEo= - RT ln Keqm
Eo= __RT_ 2.303 log Keqm
nF
log Keqm= nEo/ 0.0592 at 25
oC
5. The spontaneity or feasibility of the cell reaction can be predicted
Spontaneity of the redox reaction can be predicted from the e.m.f value of
complete cell reaction.
If the value of E cell is + ve,
(std. free energy (Go) is negative, since G
o = - nFE
o) the reaction is
feasible.
If E cell is -ve, Go
is positive. Then the reaction is not feasible.
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6. Hydrogen liberating tendency of the metal can be predicted
The metal with low reduction potential will displace H2 from an acid
solution.
Zn+ H2SO4 ZnSO4+ H2The metal with +ve potential will not displace H2 from an acid solution
Ag + H2SO4 no reactn
7. Replacing tendency of a metal (M) by another M:
The metals with low reduction potential undergo oxidation and pass into the
solution and the M with high red. pot. Undergo reduction and get deposits on
electrode.
Zn, Ni undergo dissolution in CuSO4 solution and will displace Cu from solution.
8. Corrosion tendency of M:
The metals higher in the series are anodic or more active and they are more prone
to corrosion.
The metals lower in the series are noble metals and they are less prone to
corrosion.
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Reversible and irreversible cells
Daniell cell has the emf value 1.09 volt. If an opposing emf exactly equal to
1.09 volt is applied to the cell, the cell reaction stops
Zn + Cu2+
--> Cu + Zn2+
but if it is increased infinitesimally beyond 1.09 volt, the cell reaction is
reversed.
Cu + Zn2+
--> Zn + Cu2+
Such a cell is termed a reversible cell.
Thus, the following are the main characteristics of reversible cell:
(i) The chemical reaction of the cell stops when an exactly equal opposing
emf is applied.
(ii) The chemical reaction of the cell is reversed and the current flows in
opposite direction when the opposing emf is slightly greater than that of the
cell.
(iii) The cell produces a small emf if the opposing emf is infinity smaller
than that of the cell
Any other cell which does not obey the above two conditions is termed as
irreversible. A cell consisting of zinc and copper electrodes dipped into thesolution of sulphuric acid is irreversible. Similarly, the cell
Zn|H2S04(aq)|Ag
is also irreversible because when the external emf is greater than the emf of
the cell, the cell reaction,
Zn + 2H+--> Zn
2++ H2
is not reversed but the cell reaction becomes
2Ag + 2H+--> 2Ag
++ H2
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Concentration cell
A concentration cell is made up of two half cells having identical electrodes,
identical electrolyte, except that the concentration of the reactive ions at the
two electrodes are different.
Theelectrical energy in a concentration cell arises from the transfer of
electrons from the electrode in the lower concentration side to the electrode
in higher concentration side.
The two half - cells may be joined by a salt bridge.
When a metal (M) electrode is dipped in a solution containing its own ions
[Mn+], then the potential (E) is developed at the electrode at the electrode,
the value of which varies with conc. Of the ions in accordance with the
Nernsts equation:
E =Eo - 2.303 RT / nF log 1/ [Mn+
]
E =Eo + 2.303 RT / nF log C
Let us consider a general conc. cell represented as:M/M
n+(C1M) II M
n+(C2M) I M
+
Where C1 and C2 are the concentrations of active metals ions (Mn+
) in
contact with two electrodes respectively and C2> C1.
THE CELL REACTIONS ARE:
At left electrode (anode): M --Mn+
(C1.) + ne-
At right electrode (cathode): Mn+
(C2) + ne--M
THE NET CELL REACTION: Mn+
(C2) ---Mn+
(C1.)
E.M.Fcell. of cell = Eright - Eleft
= [Eo+ 2.303RT/nF log C2] [Eo+ 2.303RT/nF log C1]
= 2.303RT/nF log C2/C1
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Ecell = 0.0592/n log C2/C1
The e.m.f so developed is due to the mere transference of metal ions from
the solution of higher conc. (C2) to lower conc. (C1.).
Applications:
i. Determination of solubility of sparingly soluble salts:
ii.
Determination of the valency of an ion.
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REFERENCE ELECTRODS:
The electrode with exactly known standard potential or fixed as constants or zero,
with which we can compare the potentials of other electrodes
The Characteristics of reference electrodes1. The electrode potential should be fixed as constants or zero.
2. The temperature co-efficient should be very low.
3. It should be a reversible electrode. It can function both as an anode and a
cathode
4. It should be easy to handle and use in laboratory.
1. Primary reference electrode
Standard Hydrogen Electrode:
Construction
This is a gas electrode.
It consists of a thin rectangular platinum foil which is coated with fresh platinum
black to increase the adsorption capacity of H2gas.
The inner tube is enclosed in an outer jacket having an inlet tube for sending in H 2
gas and has a perforated wider base for the escape of excess of H2. This unit is
dipped in 1 M HCl taken in a beaker such that the metal foil remains in the
solution.
In the above system, when the H2gas at a pressure of 1atm is bubbled through 1M
HCl, the electrode (constructed) or formed is called STANDARD HYDROGEN
ELECTRODE (SHE) or Normal H2electrode (NHE).
Working
This is represented as Pt, H2(1 atm) / H
+
(IM) When pure and dry H2gas ispassed through the inlet tube, a part of the gas gets adsorbed and the excess
bubbles out through the perforations.
Depends on the other half cell connected to SHE, SHE can under go either
oxidation or reduction.
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The electrode is represented as:
Pt, H2(1 atm); H+(1M)
The electrod potential of SHE is fixed as zero.The either of the reaction takeplace
H+(aq) + e
--- H2(g)
H2(g) --H+(aq) + e-
Act both as anode and cathode based on another M connected, it is a
reversible electrode.
When M electrode with lower reduction potential than H2, is coupled with
SHE, M undergoes oxidation and act as anode. Ex. Zn and SHE acts as
cathode
If the M electrode possesses higher reduction potential than H2, is coupled
with SHE, undergoes reduction and act as cathode. Ex. Cu
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Determination of single electrode potential of a metal
1.
If the single electrode Zn in ZnSO4 is connected to SHE then the emf of the
complete cell determined by means of a potentiometer. Since emf of SHE is
zero, the observed e.m.f gives directly the e.m.f of half-cell containing thesolution under test.
Zn has the tendency to under go oxidation, act as anode.
Pot. of the Zn single electrode can be obtained by
E cell= Ec.-Ea
Ecell= Ecathode (red.) Eanode (red.)
Ecell= ESHE Eunknown(Zn2+
, Zn)
ESHE= 0
Eunknown(Zn2+
, Zn)= 0 Ecell
Eunknown(Zn2+
, Zn)= Ecell
2. If the single electrode Cu in CuSO4 is connected to SHE then the emf of the
complete cell determined by means of a potentiometer. Since emf of SHE is zero,
the observed e.m.f gives directly the e.m.f of half-cell containing the solution
under test.
Cu has the tendency to undergo reduction & act as cathode
Pot. of the Cu single electrode
Ecell= Eunknown(Cu2+
, Cu) - ESHE
ESHE= 0
Ecell= Eunknown(Cu2+
, Cu)
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Zn possess lower red. Pot. than H2. Zn undergoes oxidation and act as anode.
H2 undergoes reduction.
H+(aq) + e
--- H2(g) --- act as cathode
Cu possess higher red. Pot. than H2. Cu undergoes reduction and
act as cathode. H2undergoes oxidation,
H2(g) --H+(aq) + e
- --- act as anode
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The electrode potential for decinormal 0.1 N KCl is +0.3335 V for normal
KCl solution +0.2810 V and for saturated calomel electrode (SCE) is
+0.2422 V
The electrode is represented as: Hg, Hg2Cl2(s), KCl (sat.solution) / Pt
The electrode potential of SCE is arbitrarily fixed as + 0.2422 V.
The electrode is based on the redox reaction
Hg22+
(s)+ 2e
-
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Eunknown (Zn2+
, Zn) = +0.2422 Ecell
2. When Cu electrode is connected to SCE
Cu undergo reduction and act as cathode and SCE act as anodeAt SCE oxidation takes place
Hg Hg++ e
-
2 Hg+
+ 2Cl2 -Hg2Cl2
The single electrode pot. Can be calculated
Ecell = Eunknown(Cu2+
, Cu) - ESCE
ESCE = +0.2422 V
Eounknown(Cu
2+, Cu) = Ecell+ 0.2422
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