UNIT - I
ELECTRO CHEMISTRY
ELECTROLYTIC AND NON-ELECTROLYTIC CONDUCTORS
Introduction
Electrochemistry deals with the interactions of electrical energy with chemical species. It is broadly
divided into two categories, namely
(i)
(ii)
production of chemical change by electrical energy (phenomenon of electrolysis)
(ii) Conversation of chemical energy into electrical energy, i.e., generation of electricity by
spontaneous red ox reactions.
In this chapter both of these aspects will be described. All electrochemical reactions involve transfer of
electrons and are, therefore, oxidation-reduction (red ox) reactions.
Substances which allow the passage of electric current through then are called electrical conductors or
simply conductors. Those which do not allow the flow of electric current through them are termed
insulators.
Electrical conductors are of two types:
(i)
Metallic or electronic conductors:
Conductors which transfer electric current by transfer of electrons, without transfer of any matter, are
known as metallic or electronic conductors.
Eg: Metals such as copper, silver, aluminum, etc., non-metals like carbon (graphite - an allotropic form
of carbon) and various alloys belong to this class.
These materials contain electrons which are relatively free to move. The passage of current though these
materials have no observable effect other than a rise in their temperature.
(ii) Electrolytic conductors:
Conductors like aqueous solutions of acids, bases and salts in which the flow of electric current is
accompanied by chemical decomposition are known as electrolytic conductors.
The substances whose aqueous solutions allow the passage of electric current and are chemically
decomposed, are termed electrolytes.
The substances whose aqueous solutions do not conduct electric current are called non-electrolytes.
Eg: Solutions of cane sugar, glycerine, alcohol, etc., are examples of non-electrolytes.
In order to pass the current through an electrolytic conductor (aqueous solution or fused electrolyte), two
rods or plates (metallic conductors) are always needed which are connected with the terminals of a
battery. These rods or plates are known as electrodes.
The electrode through which the current enters the electrolytic solution is called the anode (positive
electrode) with the electrode through which the current leaves the electrolytic solution is known as
cathode (negative electrode).
The electrolytic solution conducts electricity not by virtue of the electrolytic as in metallic conductors
but as a result of movement of charged particles called ions towards the respective oppositely charged
electrodes.
The ions which carry positive charge and move towards cathode are termed cations while ions carrying
negative charge which move towards anode are called anions.
When these ions reach the boundary between a metallic and an electrolytic conductor, electrons are
being either attached to or removed from the ions.
Removal of electrons is termed oxidation (de-electronation) which occurs at anode
while addition of electrons is called reduction (electronation) that takes place at cathode.
Hence, flow of electrons through the outer circuit from anode to cathode across the boundary is
accompanied by oxidation and reduction.
CONDUCTANCE AND CONDUCTORS
ARRHENIUS THEORY OF ELECTOLYTIC DISSOCIATION
In order to explain the properties of electrolytic solutions, Arrhenius put forth, in 1884, a comprehensive
theory which is known as theory of electrolytic dissociation or ionic theory. The main points of the
theory are:
(i) An electrolyte, when dissolved in water, breaks up into two types of charged particles, one carrying
a positive charge and the other a negative charge. These charged particles are called ions. Positively
charged ions are termed cations and negatively charged as anions.
AB
A+ + B-
NaCl
Na+ + CL2K++ SO42-
K2SO4
Electrolyte
Ions
In its modern form, the theory assumes that solid electrolytes are composed of ions which are held
together by electrostatic forces of attraction. When an electrolyte is issolved in a solvent, these forces are
weakened and the electrolyte undergoes dissociation into ions. The ions are solvated.
A+Bor
A+B-+ aq
A+ + BA+(aq)+B- (aq)
(ii) The process of splitting of the molecules into ions of an electrolyte is called ionization. The
fraction of the total number of molecules present in solution as ions is known as degree of ionization or
degree of dissociation. It is denoted by
α= (Number of molecules dissociated into ions)/(Total number of molecules)
It has been observed that all electrolytes do not ionize to the same extent. Some are almost completely
ionized while others are feebly ionized. The degree of ionization depends on a number of factors (see
12.6).
(iii) Ions present in solution constantly re-unite to form neutral molecules and, thus, there is a state of
dynamic equilibrium between the ionized the ionized and non-ionised molecules, i.e.,
AB
A+ + B-
Applying the law of mass action to above equilibrium
[A+ ][B- ] [AB] =K
K is known as ionization constant. The electrolytes having high value of K are termed strong
electrolytes and those having low value of K as weak electrolytes.
(iv) When an electric current is passed through the electrolytic solution, the positive ions (cations)
move towards cathode and the negative ions (anions) move towards anode and get discharged, i.e.,
electrolysis occurs.
The ions are discharged always in equivalent amounts, no matter what their relative speeds are.
OHM’S LAW
According to ohm’s law the current ‘l’ flowing through a conductor is given by the relation E/R
Where E is electromotive force the difference of potential at the two ends of the conductor and ‘R’ is
resistance
l=E/R
Te strength of the current is usually measured in “amperes”
The quantity of current passing through a conductor is the product of current strength and time &
express as coulomb.
Definition of coulomb: The quality of current which passes in one second with a current strength of one
ampere.
When current is measured in ampere and potential difference in volts the resistance offered is a fixed
value called “ohm” (Symbol Ώ)
Ohm = volt / Ampere
Conductance, Specific conductance and Molar conductance
Conductance: Reciprocal of resistance is called as conductance.
C = 1 / R.
The unit of conductance is ohm-1 or mho.
Specific resistance:
The resistance of a conductor is directly proportional to the length(l) and inversely proportional to the
area of cross section of the conductor (or) area of cross section is denoted by S (.-.a=S)
R α l/a
R = ρ x l/a
Where “ρ” is a constant called specific resistance now if l = 1cm; a=1cm2.
Then R = ρ
Specific resistance is defines as the resistance in 1ohm of a substance of 1cm in length and 1cm2 in area
of cross section.
Units:
R = ρ x l/a
ρ = Ra/l
ρ = ohm cm2/cm = ohm.cm
ρ = ohm . cm
Specific Conductance (K):
The reciprocal of specific resistance is called “Specific Conductance”.
It is represented by K (Kppa).
K = 1/ρ
= 1/ohm-cm
= ohm-1cm-1
Units of specific conductance = ohm-1cm-1 ( Or ) Mho cm-1
Where a=s;
K=1/ρ
K = R x S/ L
(L/S = cell constant)
K = 1/R x L/S
(R = resistance)
Specific Conductance = Cell constant/Resistance
Cell constant = Specific conductance x Resistance
Equivalent Conductance (Λeq)
It is defined as the conductance of all the ions produced by the dissociation of 1gm equivalent of an
electrolyte dissolved in a certain volume (V) of the solvent at a constant temperature.
Λv = KV
Where v= volume in c.c containing an electrolyte
Where K = 1/R x 1/S.
Λv = 1/R x 1/S x V
= 1/ohm x cm/cm2 x cm3/equi V
Λv = ohm-1cm2equiv-1.
Relationship between Specific conductance and equivalent conductance
Λv = kv x 1000/CV
Where V=1000/N
Where V = volume in electrolyte of 1gm equivalent of 1000 cc
N = Normality
Molar Conductance(µv)
The molar conductance is defined as the conductance of all the ions produced by ionization of 1 g mole
of an electrolyte when present in V mL of solution. It is denoted by .
Molar conductance
μ = k ×V
.... (vii)
where V is the volume in mL containing 1 g mole of the electrolyte. If c is the concentration of the
solution in g mole per litre, then
μ = k × 1000/c
It units are ohm-1 cm2 mol-1.
Equivalent conductance = (MOlar conductance)/n
where
n = (Molecular mass)/(Equivalent mass)
APPLICATIONS OF CONDUCTANCE
CONDUCTOMETRIC TITRATIONS
DEFINATION: The conductometric titrations are defined as the titrations in which equivalent
point is determined with the help of conductivity measurements
PRINCIPLE: The electrical conductance of an electrolytic solution is proportional to (a) The ionic
concentration i.e. numbers of ions present in the solution and (b) The ionic mobilities. Conductometric
measurements based on the fact of measuring the ease or speed of the ion. Thus more the mobility of the
ion more will be the conductance
ADVANTAGES:
These titrations are more advantageously used in case of
1.
2.
3.
4.
Colored solutions where the selection of indicator is difficult
Acid –base and mixture of acid titration
In case of precipitation titrations
No need of any indicator
5. No special care is required near the end point
DISADVANTAGES: These titrations are not applicable in case of high concentration solutions with
high impurities
TYPES OF CONDUCTOMETRIC TITRATIONS :
1. TITRATION OF STRONG ACID AND STRONG BASE:
PRINCIPLE:
The ionic mobilities in aqueous solutions are generally of the order of 6.0 x 10-4 cm2, sec-1, volt-1. But
the ionic mobilities of H+ and OH- are abnormally high, H+ = 36.3 x 10-4 and OH- = 20.5 x 10-4 cm2, sec1
, volt-1. Thus the conductance of an electrolyte is quite sensitive to the concentration of H+ and OHions.
In conductometric acid – base titrations, the course of the titration as base is added to the acid or acid
to the base is followed from the gradual change in the conductivity of the titrant.
When HCl is completely ionised, the mobility of H+ ions is high; hence the conductance of HCl solution
will also be high. As NaOH is added to HCl, the fast moving H+ ions are removed by OH- ions of the
base as water.
( H + + Cl- ) + (Na+ + OH-)
NaCl + H2O
In solution water is little ionised and Na+ added are slow moving compared to H+ ions of the acid.
Hence as the titration proceeds, the conductance of the titrant HCl decreases gradually till the equivalent
point is reached.. At the equivalence point the conductance of the solution (Na+, Cl- and H2O) will be the
minimum. Addition of the base beyond end point adds Na+ and fast moving OH- ions, resulting in a
gradual increase in conductance.
PROCEDURE







The burette is rinsed and filled with 0.1M NaOH.
40ml of the given acid is placed in a 100ml beaker.
The cell washed with conductivity water is placed in HCl solution and is seen that the electrodes
of the cell are completely immersed in the solution..
The conductance of HCl solution is noted.
Add 1ml of NaOH from the burette into the HCl solution.
After each addition, the solution is stirred gently with a glass rod and the conductance of HCl is
noted.
The titration is continued till 20 ml of NaOH is added.
GRAPH:
TITRATION OF WEAK ACID (CH3COOH)BY STRONG BASE(NaOH):




Weak acid is charectacterized through its insignificant dissociation .
Therfore at the beginning of titration acid solution has a low conductance.
Addition of alkali to the weak acid solution initially results in a small fall in conductance , owing
to the replacement of H+ ions by acetate ions.
After the addition of a few drops of alkali , the conductance increases in a regular fashon ,due to
the actate ions formed and increase in population of acetate ions in the solution.
CH3COOH + Na+ + OHCH3COO- + Na++ H2O
TITRATION OF MIXTURE OF ACIDS WITH STRONG BASE:




The concentration of acids in a mixture can be determined in one titration
Due to the presence of striong acid ,the dissosiation of a weak acid is completely supressed due
to commenion effect
Therfore the addition of alkali will first results in neutralization of the strong acid
The weak acid start reacting only after neutralization of the strong acid .
PRECIPITATIVE TITRATIONS:




Concider titration of Ag+ ions with KCl ,on addition of any chloride ions the silver ions fall out
of aqueous solution forming an insoluble silver chloride ppt
AgNO3 + KC
AgCl (ppt) + K+ + NO3
Since mobility of Ag+ ions &K+n ions are similar to each other ,therefore no change in
conductance by the addition of KCl toAgNO3.
After neutralization of all Ag+ ions by K+ ,there is raise in conductance representes the presence
of K+ ions in exess.
But if AgNO3 is titrated against NaCl,The size and mobility of Na+ ion is different from Ag+
ion ,therefore decrease in conductance observes until all Ag+ ions are replaced by Na+ ions.
ELECTROCHEMICAL CELL
Electrochemical cell can be defined as a single arrangement consisting of two electrodes and capable of
producing electricity due of chemical reaction and vise versa
Electrochemical cell is a system or arrangement in which two electrodes are fitted in the same
electrolyte or in two different electrolytes which are joined by a salt bridge. Electrochemical cells are of
two types:
(a)
Electrolytic cell
(b)
Galvanic or voltaic cell
Electrolytic cell
It is a device in which electrolysis (chemical reaction involving oxidation and reduction) is carried out
by using electricity or in which conversion of electrical energy into chemical energy is done.
Galvanic or voltaic cell



It is a device in which a redox reaction is used to convert chemical energy into electrical energy,
i.e., electricity can be obtained with the help of oxidation and reduction reaction. The chemical
reaction responsible for production of electricity takes place in two separate compartments.
Each compartment consists of a suitable electrolyte solution and a metallic conductor. The
metallic conductor acts as an electrode.


The compartments containing the electrode and the solution of the electrolyte are called halfcells.
When the two compartments are connected by a salt bridge and electrodes are joined by a wire
through galvanometer the electricity begins to flow. This is the simple form of voltaic cell.
DANIELL CELL
It is designed to make use of the spontaneous redox reaction between zinc and cupric ions to produce an
electric current (Fig.12.7). It consists of two half-cells. The half-cells on the left contains a zinc metal
electrode dipped in ZnSO4solution.
The half-cell on the right consists of copper metal electrode in a solution CuSO4. The half-cells are
joined by a salt bridge that prevents the mechanical mixing of the solution.
When the zinc and copper electrodes are joined by wire, the following observations are made:
(i) There is a flow of electric current through the external circuit.
(ii) The zinc rod loses its mass while the copper rod gains in mass.
Zn
Cu+2 +2e-
Zn+2 + 2eCu
(iii) The concentration of ZnSO4 solution increases while the concentration of copper sulphate solution
decreases.
(iv) The solutions in both the compartments remain electrically neutral.
During the passage if electric current through external circuit, electrons flow from the zinc electrode to
the copper electrode. At the zinc electrode, the zinc metal is oxidized to zinc ions which go into the
solution. The electrons released at the electrode travel through the external circuit to the copper
electrode where they are used in the reduction of Cu2+ ions to metallic copper which is deposited on the
electrode. Thus, the overall redox reaction is:
Zn(s) + Cu2+
Cu(s) + Zn2+(aq)
Thus, indirect redox reaction leads to the production of electrical energy. At the zinc rod, oxidation
occurs. It is the anode of the cell and is negatively charged while at copper electrode, reduction takes,
place; it is the cathode of the cell and is positively charged.
Thus, the above points can be summed up as:
(i) Voltaic or Galvanic cell consists of two half-cells. The reactions occurring in half-cells are called
half-cell reactions. The half-cell in which oxidation taking place in it is called oxidation half-cell and the
reaction taking place in it is called oxidation half-cell reaction. Similarly, the half-cell occurs is called
reduction half-cell and the reaction taking place in it is called reduction half-cell reaction.
(ii) The electrode where oxidation occurs is called anode and theelectrode where reduction occurs
is termed cathode.
(iii)
Electrons flow from anode to cathode in the external circuit.
REPRESENTATION OF AN ELECTROCHEMICAL CELL (GALVANIC CELL)
It is very difficult to always sketch the electro chemical cell . Thus a symbolic representation is often
used to describe a cell .This representation is called a cell diagram.
The following universally accepted conventions are followed in representing an electrochemical cell:
(i) The anode (negative electrode) is written on the left hand side and cathode (positive electrode) on the
right hand side.
(ii) A vertical line or semicolon (;) indicates a contact between two phases. The anode of the cell is
represented by writing metal first and then the metal ion present in the electrolytic solution. Both are
separated by a vertical line or a semicolon. For example,
Zn|Zn2+ or Zn;Zn2+
The molar concentration or activity of the solution is written in brackets after the formula of the ion. For
example:
Zn|Zn2+(1 M) or Zn | Zn2+(0.1 M)
(iii) The cathode of the cell is represented by writing the cation of the electrolyte first and then metal.
Both are separated by a vertical line or semicolon. For example,
Cu2+|Cu or Cu2+;Cu or Cu2+(1 M)|Cu
(iv) The salt bridge which separates the two half-cells is indicated by two parallel vertical lines.
(v) Sometimes negative and positive signs are also put on the electrodes.
The Daniell cell can be represented as:
Zn|ZnS04(aq)||CuS04(aq)|C+u
Anode
Salt bridge
Oxidation half-cell
Cathode
Reduction half-cell
or Zn|Zn2+||Cu2+|Cu
or
Zn|Zn2+(1 M)||Cu2+(1 M)|Cu
ELECTRODE POTENTIAL
DEFINATION:A tendency of an electrode to loss or gain electrons when it is in contact with its own
ionic solution is called electrode potential
When a metal is placed in a solution of its ions, the metal acquires either a positive or negative charge
with respect to the solution.
On account of this, a definite potential difference is developed between the metal and the solution. This
potential difference is called electrode potential.
For example, when a plate of zinc is placed in a solution having Zn2+ ions, it becomes negatively
charged with respect to solution and thus a potential difference is set up between zinc plate and the
solution. This potential difference is termed the electrode potential of zinc.
Zn
Zn+2 + 2e-
(oxidation potential)
Similarly, when copper is placed in a solution having Cu2+ ions, it becomes positively charged with
respect to solution. A potential difference is set up between the copper plate and the solution. The
potential difference thus developed is termed as electrode potential of copper.
Cu+2 +2e-
(a)
Cu
Oxidation:
Metal ions pass from the electrode into solution leaving an excess of electrons and thus a negative
charge on the electrode.
(b)
Reduction:
Metal ions in solution gain electrons from the electrode leaving a positive charge on the electrode.
Interface of metal and the solution. The development of negative charge (as on zinc plate) or positive
charge (as on copper plate) can be explained in the following manner. When a metal rod is dipped in its
salt solution, two changes occur:
(i) The conversion of metal atoms into metal ions by the attractive force of polar water molecules.
M --> Mn + ne-
The metal ions go into the solution and the electrons remain on the metal making it negatively charged.
The tendency of the metal to change into ions is known as electrolytic solution pressure.
(ii) Metal ions start depositing on the metal surface leading to a positive charge on the metal.
Mn+ + ne- --> M
This tendency of the ions is termed osmotic pressure.
Single Electrode Potential (E)
It is a measure of the tendency of a metallic electrode to loose or gain
electrons when it is dipped in its own salt solution.
Standard Electrode Potential (Eo) (SEP)
It is a measure of the tendency of the metallic electrode to loose or gain
electrons when it is dipped in its own salt solution of unit concentration
(1M), at 25oC and atmospheric pressure
EMF OF A GALVANIC CELL
Every galvanic or voltaic cell is made up of two half-cells, the oxidation half-cell (anode) and the
reduction half-cell (cathode). The potentials of these half-cells are always different. On account of this
difference in electrode potentials, the electric current moves from the electrode at higher potential to the
electrode at lower potential, i.e., from cathode to anode. The direction of the flow of electrons is from
anode to cathode.
Flow of electrons
Anode <==============> Cathode
Flow of current
The difference in potentials of the two half-cells is known as the electromotive force (emf) of the cell or
cell potential.
The emf of the cell or cell potential can be calculated from the values of electrode potentials of the two
half-cells constituting the cell. The following three methods are in use:
(i) When oxidation potential of anode and reduction potential of cathode are taken into account:
ECello = Oxidation potential of anode
+ Reduction potential of cathode
= Eoxo (anode) + Eredo (cathode)
(ii) When reduction potentials of both electrodes are taken into account:
ECello = Reduction potential of cathode - Reduction potential of anode
= ECathodeo - EAnodeo
= Erighto - Elefto
(iii) When oxidation potentials of both electrodes are taken into account:
= Oxidation potential of anode - Oxidation potential of cathode
ECello = Eoxo (anode) - Eredo (cathode)
NERNST EQUATION
ELECTRODE AND CELL POTENTIALS /NERNST EQUATION
The electrode potential and the emf of the cell depend upon the nature of the electrode, temperature
and the activities (concentrations) of the ions in solution. The variation of electrode and cell potentials
with concentration of ions in solution can be obtained from thermodynamic considerations. For a general
reaction such as
M1A + M2B .....
n1X + n2Y +.... .......(i)
occurring in the cell, the Gibbs free energy change is given by the equation
G = ∆Go + 2.303RT log10 (axn1 × ayn2)/(aAm1 × aBm2) ....... (ii)
where 'a' represents the activities of reactants and products under a given set of conditions and
∆Go refers to free energy change for the reaction when the various reactants and products are present at
standard conditions.
The free energy change of a cell reaction is related to the electrical work that can be obtained from the
cell, i.e., ∆Go = -nFEcell and ∆Go = -nFEo. On substituting these values in Eq. (ii) we get
-nFEcell = -nFEo + 2.30eRT log10 (axn1 × ayn2)/(aAm1 × aBm2) ....... (iii)
or Ecell = Ecello - 2.303RT/nF log10 (axn1 × ayn2)/(aAm1 × aBm2) ....... (iv)
This equation is known as Nearnst equation.
Putting the values of R=8.314 JK-1 mol-1, T = 298 K and F=96500 C, Eq. (iv) reduces to
E = Eo - 0.0591/n log10 (axn1 × ayn2)/(aAm1 × aBm2) ....... (v)
= Eo - 0.0591/n log10 ([Products])/([Reactants]) ....... (vi)
ELECTRO-CHEMICAL SERIES:
The electrode potentials potentials of all electrodes which are determined with potentiometer are
different from each other . some of them are having negative electrode potential values compared to the
electrode potential of hydrogen electrode
Eg: Zn2+/Zn. (-0.762)Cd2+(aq)/Cd(-0.403) , Ni++(aq)/ Ni(-0.250)
Some of them are having positive electrode potential values compared to the electrode potential of
hydrogen electrode
Eg : Cu2+(aq)/Cu(=0.337); Ag+(aq)/Ag(=0.799); Au3+(aq)/Au(+1.50)
Based on these measured values of the standared potentials of various electrodes , the electrode systems
are arranged in a series in the ascending order of their reduction potentials. This series is known as
electro-chemical series
The higher up Metal in the series have greater tendency to be oxidized
The metals high up in the series are strong reducing agents and their ions are stable and near the bottom
of the series are inactive
Any metal above hydrogen will displace hydrogen from dilute acid solution
Eg: Na reacts with water to liberate hydrogen the oxidization from any metal below the hydrogen in the
series will be reduced by hydrogen
The metal above the hydrogen in the series can be easily oxidized hence they undergo corrosion easily
Li+(aq)/Li with higheast negative potential value is placed at the top of the series and the electrode with
highest positive value pt,F2/2F-(aq) is placed at the bottom most position of the series.
Electrode
Li/Li
Half cell reaction
Li
Li++ e
E0 Volts
-3.045
K/K+
K
K+ + e
-2.925
Ca/Ca++
Ca
Ca++ + 2e-
-2.87
Na/Na+
Na
Na+ = e
-2.714
Mg2+/Mg
Mg
Mg2+ + 2e
-2.37
Zn/Zn2+
Zn
Zn2+ + 2e-
-0.763
Fe/fe2+
Fe
fe2= + 2e
-0.440
Cd/cd2+
Cd
Cd2+ + 2e
-0.403
Pb/pb2+
Pb
Pb2+ + 2e
-0.126
Pt/H2(g)/H+
H2
2H+ + 2e
0.000
Cu/cu2+
Cu
Cu2+ + 2e
+0.337
Ag/Ag2+
Ag
Ag+ + e
+0.799
Pt/cl-cl2
2cl-
cl2 + 2e
+1.36
Au/Au3+
Au
Au3 + 3e
+1.50
Application / Significance of electrochemical series
(i)
Relative ease of oxidation or reduction
• The metals which lie above hydrogen in the series undergo
spontaneous oxidation and the metals which lie below SHE undergo
reduction spontaneously ( ie. Acts as Anodes and Cathodes
respectively)
Active anode
Noble Cathode
• The metals which lie above hydrogen are good reducing agents and
which lies below hydrogen will act as good oxidizing agents
(ii) Replacement tendency

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The metal lying above in emf series displaces the metal lying below it
from an electrolyte of the later.
Example 1: Ni spatula cannot be used to stir copper sulphate solution
due to the following reaction
Ni(s) + Cu2+ (aq) ___ Ni 2+ (aq) + Cu (s)
Example 2: when zinc is dipped in copper sulphate solution copper gets
deposited (displaced)
Zn (s) + CuSO4 (aq)
Zn SO4 (aq) + Cu (s)
(iii) Liberation of Hydrogen
• The metal with negative reduction potential will displace H2 from an
acid solution
Zn (s) + 2 HCl (aq)


Zn Cl2 (aq) + H2 -
Hence acids cannot be stored in galvanized steel containers.
For exactly the same reason galvanized steels are not used to store
food stuffs containing vinegar. (Vinegar is used as food preservative ; vinegar is acetic acid)
(v) Calculation of Standard emf of the cell
E cell = E cathode − E anode
( if both reduction potentials are considered)
E cell = E cathode + E anode
( if oxidation potential of anode and the reduction potential of cathode
are considered)
(vi) Corrosion
• The metals higher in the series are anodic and are more prone to
corrosion.
• The metals lower in the series are noble metals (cathodic) and they are
less prone to corrosion.
(vii) Predicting the spontaneity of cell reaction
• Spontaneity of the redox reaction can be predicted from the emf value
of complete cell reaction.
If the value of Ecell is positive, the reaction is feasible. as _G will be
negative ( i.e. it is an electrochemical cell)
If the value of Ecell is negative, the reaction is not feasible. as _G will be
positive ( i.e. it is an electrolytic cell)
.
MEASUREMENT OF SEP:
SEP cannot be measured directly. The electrode is coupled with a reference
electrode
Examples: Standard Hydrogen electrode (SHE)
Saturated Calomel Electrode (SCE)
REFERENCE ELECTRODES
The electrode of standard potential with which we can compare the
potentials of other electrodes is called a reference electrode.
The potential of the electrode remains constant at all temperatures.
It undergoes specific reduction or oxidation but the potential will be same
only sign will be different.
TYPES OF REFERENCE ELECTRODES:
Standard Hydrogen Electrode (SHE) /Normal Hydrogen Electrode (NHE)




Type of electrode: Gas electrode (Primary Reference Electrode)
Components: Electrode component: Pt-H2; Electrolyte component: HCl (1M)
Electrode representation: Pt, H2 (1atm) / H+ (1M)
Construction: Hydrogen electrode consists of a Platinum foil connected to a platinum wire
sealed in a glass tube. The electrode is in contact with 1M HCl and hydrogen
gas (1 atmosphere) is constantly bubbled.
 Limitations: 1.
2. It requires large volume of test solution
3. The potential of the electrode is dependent on atmospheric pressure
Reactions
As anode
H2
2 H+ + 2e-
As cathode
2 H+ + 2eSince a standard hydrogen electrode is
maintain, it is usually replaced by other
which are known as secondary reference
convenient to handle and are prepared
secondary reference electrodes are described
Eo = 0 V
H2
difficult to prepare and
reference
electrodes,
electrodes. These are
easily. Two important
here
METAL- METAL SALT ELECTRODE:
Saturated Calomel Electrode (SCE):
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Type/ class: Metal- metal insoluble salt electrode (Secondary Reference
Electrode)
Components: Electrode component: Pt – Hg ; Electrolyte component: Hg2Cl2(s) / KCl
Electrode representation: Hg, Hg2Cl2(s) - KCl (sat. solution)
Construction: Calomel electrode consists of a glass tube containing mercury at the bottom
over which mercurous chloride paste (calomel) is placed. The tube is filled
with saturated KCl solution. A platinum wire is fused into the layer of
mercury to provide electrical contact. The electrode potential differs with the
concentration of KCl.
Reactions:
As anode:
2Hg (l)
Hg22+ +Cl-
Hg22+ + 2e-
Hg2Cl2
As cathode
Hg22+ (2Cl -) + 2eCl
2Hg (l) +2
The net reaction can be
represented as
2Hg + 2Cl
Hg2Cl2
E is 0.3335 V for 0.1N KCl
(DNCE)
Applying Nernst equation for the net
reaction
E is 0.2810 V for 1N KCl (NCE)
Ecell = Eocell + 0.0591/n . log
[Hg]2[Cl−]2 / [Hg2Cl2]
E is 0.2422V for Sat. KCl (SCE)
The [Hg] and [Hg2Cl2] are unity .Therefore the Ecell depends on the
concentration of chloride ions , hence the electrode is said to be reversible
wrt chloride. (Another electrode which is reversible wrt to Chloride ions is Ag/ AgCl)
The concentration of chloride ions
ses during oxidation and _ses during
reduction.
E is 0.3335 V for 0.1N KCl (DNCE)
E is 0.2810 V for 1N KCl (NCE)
E is 0.2422V for Sat. KCl (SCE)
The electrode potential of any other electrode on hydrogen scale can be measured when it is combined
with calomel electrode. The emf of such a cell is measured. From the value of electrode potential of
calomel electrode, the electrode potential of the other electrode can be evaluated.
REDOX ELECTRODE:
If any electrolyte contains different oxidation states ,to get stability ions may loss or gain electrons . In
such cases the electrolyte acts as redox electrode.
Eg:1 If solution of iron contains both Fe2+ and Fe3+ states ,the Pt wire which is immersed in the
solution can have potential or can acquire electricity
Fe3+ + 2e
Fe 2+
Ce4+ + 2e
Ce 3+
Sn 4+ + 2e
Sn 2+
Eg: 2 QUIN HYDRONE ELECTRODES: It is 1:1 mixture of quinine and hydroquinone. It is reversible electrode with respect to H+ ion
Q + 2H+ + 2e
-
QH2
C6H4o2 + 2H+ + 2e
C6H4(OH)2
The Potential ‘E’ developed when an inner electrode Eg: Platinum immersed in this system is given by
Nernst reduction equation
E = E0 + 2.303
log10[QH2]
nf
[Q][H+]2
E = E 0 + 2.303RT log10[QH2] – 2.303RT log10[H+]2
2F
[Q]
2F
The ratio [QH2]/[Q] is maintained constant unity by saturating the solution with the substance
“Quinhydrone” Which is a 1:1 motor compound of quinine & hydeoquinous.
Since log101 = 0 & log10[H+]2 = 2log10[H+]
E = E0 – 2.303RT/2F x 2log10[H+]
E0 – 2.303RT/F log10[H+]
E = E0 – 0.0591 pH
Since electrode potential of quinhydronated depends upon the concentration of hydrogen ions.
It can be used for the determination of Potential value of H+ ions just like Hydrogen electrode
Standard oxidation potential of quinhydrone electrode at 250C is -0.6994 volts
Construction of electrode : adds a pinch of quinhydrone to the solution under examination and immerses
a clean platinum electrode in it.
Advantages and limitations of quinhydrone electrode




The electrode is very easy to setup.
PH value is accurate
Small qualities of solution are sufficient for the measurement
Due to instability of quinine in strong alkaline medium this hydrogen ion indicater is suitable
only up to a pH of 8
Ion selective electrode (ISE) or specific ion electrode(SIE)


Ion selective electrodes (ISE) use a membrane which is sensitive to a particular species
ISE develops potential by measuring of the concentration of the species of interest
Measurement of Concentration of Cations



The cell constructed for the measurement of concentration of cations [Mn+] is given below
Reference electrode solution to be analyzed || Internal standard solution
In ion selective electrodes the use of membrane is must generally the following four types of
membranes are used
a. Glass Membrane
b. Solid state membrane
c. Heterogenous membranes
d. Liquid membrane

The liquid membrane is usually a large organic molecule capable of interacting with anions and
cations and is absorbed
Membrane potential is given by the equation

EMn+ = RT/nf logeC2/C1

Suppose reference electrode is anode and membrane electrode is cathode then the cell potential
Ecell = Emn+ - ΔEref

Here the ΔEref is zero if identical electrodes were used as the concentration of Mn+ in the internal
reference (C2) is constant the Ecell is given by
Ecell = constant – RT/nf logeC1
Applications
1. To determine the concentration of gas levels like level in blood samples
2. To determine the concentration of a number of cation like H+,Li,Na+,K+,Ag+ ------3. To find the concentration of anions like f-,No3- etc
4. To determine the pH of solution
MEASUREMENT OF pH BY GLASS ELECTRODE:






It is an indicator electrode for H+ions
The potential of glass electrode depends up on the activity (the apparent concentration) of H +
ion in solution
The glass electrode consists of an electrically conducting glass membrane in the form of thin
bulb ,which is fused to a glass tube.
The bulb contains HCl solution and internal reference electrode such as Ag – AgCl electrode or
Pt wire in the solution of constant pH
The electrode may represented as
Ag /AgCl(s) , HCl(0.1M)/glass
Working of glass electrode: The glass electrode combined with a reference electrode and
electrode potential is measured by a potentiometer
The value of Ecell is the algebraic sum of the glass electrode potential and the potential of the
reference electrode.
Ecell = Eise --





Eref
The glass develops a potential Eg at the surface of the glass membrane ,when it is dipped in to
an aqueous solution having different pH compared to that of the solution within the bulb.
The potential related to the transfer of H+ ions through the glass membrane b/w the solution of
hydrogen ion activity (a1) and (a2)
Therfore Eg = -0.0591 log(a1/a2)
Where a1 = activity of H+ ions of external solution
a2 = activity of H+ ions in internal solution
Even when the activity of H+ ions of solutions on either side opf the membrane is are same , the
emf is not zero . due to difference in the response to pH by inner and outer surface of the
membrane . This is known as Asymmetry potential
The Glass electrode also sensitive to ions such as Na+in addition to H+ ions
A concentration cell is a galvanic cell in which the electric energy is produced by the transfer of material
from a system of higher concentration to lower concentration.
There are two types of concentration cells:
1. Electrode concentration cells.
2. Electrolyte concentration cells
i)
Electrode concentration cells:


In these cells, the potential difference is developed between two like electrodes at different
concentrations dipped in the same solution of the electrolyte.
For example, two hydrogen electrodes at different has pressure in the same solution of hydrogen
ions constitute a cell of this type.
(Pt,H2 (Pressure p1))/Anode |H+ | (H2 (Pressure p2)Pt)/Cathode


If p1, p2 oxidation occurs at L.H.S. electrode and reduction occurs at R.H.S. electrode.
Two hydrogen electrodes at unequal pressure P1 and P2 immersed in the same solution of H+ ions
the following reaction takes place
Oxidation reaction
H2(P1)
2H+ + 2e
Reduction reaction
Total reaction


2H+ + 2e-
H2(P1)
H2cp2
H2(P2)
It is clear that in this process there is overall chemical change and there is only transfer of H2 gas
from the electrode with pressure P1 to the electrode with pressure P2.
The emf depends only on the two pressures and independent of the concentration of electrolyte
solution hydrogen ions in which the electrodes are dipped
E = RT/2f loge P1/ P2 ;
Ecell = 0.0591/2 log(p1/p2) at 25o C
Eg:2 In amalgam cells two amalgams of same metal at two different concentrations are immersed in the
same solution of metal ions
Eg : Two unequal concentrations of zinc amalgams dipped in a solution of Znso4
Zn(Hg)C1 | Zn++|Zn (Hg) C2
Overall reaction is Zn(Hg) C1
Zn(Hg)C2
E = Rt/2F loge C1/C2
(ii)
Electrolyte concentration cells:





In these cells, electrodes are identical but these are immersed in solutions of the same electrolyte
of different concentrations.
The source of electrical energy in the cell is the tendency of the electrolyte to diffuse from a
solution of higher concentration to that of lower concentration.
With the expiry of time, the two concentrations tend to become equal.
Thus, at the start the emf of the cell is maximum and it gradually falls to zero.
Such a cell is represented in the following manner:
(C2 is greater than C1).
M|Mn+(C1)||Mn+(C2)|M
(Zn|Zn2+ (C1))/Anode || (Zn2+ (C2 )|Zn)/Cathode
or

The emf of the cell is given by the following expression:
Ecell = 0.0591/n log C(2(R.H.S.))/C(1(L.H.S.)) at 25o C

The concentration cells are used to determine the solubility of sparingly soluble salts, valency
of the cation of the electrolyte and transition point of the two allotropic forms of a metal used
as electrodes, etc.
Eg:
when the zinc electrodes dipped in two solutions of Znso4 with different concentration C1
&C2joined through a salt bridge can be represented as follows
(+) cathode Zn|Zn+2 (C1) ||Zn 2+(C2)| Zn- Ano
The electrode reaction on the left side is oxidation
Zn++(C2) + 2e-(anode) oxidation.
Zn
Zn++(C1) + 2e
Zn(+) (cathode) reduction.
The emf of such cell can be calculated by the following expression based on Nernst equation
E = 2.303Rt/nf log10C2/C1
Substituting the value of R = (8.313 x 107) T at 250C (213+25 = 298K) and f = 96500
columbs can calculate the emf of cell as follows
E = 0.0591/n log10C2/C1
Applications of Concentration cells:
1.
2.
3.
4.
To determine the solubility of a sparingly soluble salt
To calculate the valence of cations.
To determine transition point.
To calculate the extent of corrosion in metals
Potentiometric Titrations:
A titration in which the equivalent or end point of a reaction is determined with the help of the
potential of the reaction mixture is known as potentiometric titrations
Acid base titration:
Principle:



During the titration, against an alkali, H+ ion concentration in the half-cell containing QH2
will decrease. Correspondingly the Ecell decreases.
The emf will change slowly in the beginning, however, at the end point relatively higher
potential difference will be obtained and then again it will change slowly.
After the end point 5 or 6 ml more of the base are added quickly and, emf is recorded.
Procedure:







The acid solution strength has to be determinate is taken in beaker
The hydrogen electrode or QH2 electrode and colomel electrode were dipped in the acid
solution
The electrodes were connected to the potentiometer & emf is measured
A standard known volume of alkali solution is taken in burette
Subsequently Alkali solution added to acid solution with the help of burette ,stirred
thoroughly & emf of cell is recorded
Like the same procedure 10-15 readings to be taken
Repetition of the titration is done by adding 0.2ml of NaOH at a time near the end point.A
GRAPH:

graph is plotted between emf and volume of base added. The point of inflexion where the
curve changes its direction is the end point.

(average of two successive volumes of the base)
gives a curve whose peak indicates equivalence point.
Redox titrations:
Oxidation reduction type it is same as acid-base titration
Only difference is that hydrogen electrode is replaced by bright platinum electrode
The emf of the electrode is determined by activating the ratio of oxidized or reduced
Eg: fe2+ is titrated against K2Cr2o7
Fe2+ is treated with H2so4 & it should be taken in beaker and platinum and calomel electrodes are dipped
in beaker electrodes are connected to potentiometer
K2cr2o7 is taken in burette
Graph is drawn by plotting with emf reading and volume of K2cr2o7
The step rise is end point of the titration
Precipitation Titrations In this titration electrode is reversible
Silver electrode is used along with calomel electrode
Eg: Titration of silver nitrate with sodium chloride where silver chloride precipitates out
Silver nitrate is take in micro burette and sodium chloride in beaker containing electrodes
Emf of the cell is measured and plotted against volume of silver nitrate added.
Step rise in the curve show the end point of the titration
AgNo3 + Nacl
Agcl + NaNo3
Determination of PH by emf method
The emf of a solution depends on the concentration of H+ ions or PH of the solution
A hydrogen electrode containing solution of unknown PH is paired with a standard catomel electrode
The complete cell may be represented as Pt,H2(1atm) |H+(unknown)||Kcl(salt sol)|Hg2cl2Hg
Emf of above cell measured by potentiometer and PH of unknown solution can be calculated as follows
Ecell = Eright - Eleft
Ecell = 0.2415 – (-0.0591 x PH)
Ecell = 0.2415 + 0.0591 PH
Advantages of potentiometric titration
1. Coloured solutions where the use of indicator is impossible are estimated by this titrations
2. Solutions containing more than one halide can be analyzed in a single titration against silver
nitrate
3. Poly basic acids can be titrated in steps correspond to different steps of neutralization
4. By this titration there is no problem with the indicators bases on PH value of the solutions
BATTERIES
Battery is an electrochemical cell
A device which converts chemical energy to electrical energy is called battery
Battery applied to a group of two or more electric cells connected together electrically in series
Advantages of Batteries
1.
2.
3.
4.
Batteries can act as source of electric energy
Electronic equipments in the form of handsets have been made possible by batteries
Batteries for automotive and aircrafts stand by batteries
A variety of electronic gadgets have been made more use and popular with the introduction of
rechargeable storage batteries
The following requirements should be possessed by the batteries
High energy efficiency which is calculated as
% of efficiency = energy released on discharge x 100
Energy required for voltage
Reliability is another important criteria
Tolerance to different service conditions such as variation in temperature, vibration, shock etc.
We can broadly classify the batteries into three categories
a. Primary cell
b. Secondary cell
c. Fuel cells
1. The primary cells: The cell in which cell reaction is not reversible
When the reactants converts into products and electricity is produced and the cell becomes dead
cell cannot be used after that
These batteries are used for power od DC
Eg: Dry or teclanch cell
A cell without fluid component is primary cell
The anode of the cell is Zinc or containing electrolyte of NH4cl,Zncl2 and Mno2
To which satarch is added to prevent leakage
A carbon (graphite) rod serves as cathode which is immersed in the centre of cell
Net reaction is as follows
Zn(s)+2NH42+ +2cl-(ag) + 2Mno2
Mn2o3 + [Zn(NH3)2]cl2 + 2l
Dry cell find its applications in flash lights transmission radius, calculators etc
Lithium Cells Lithium cells belong primary cells.
s
The cells having lithium anodes are called lithium cells
Lithium cells can be classified as follows
a. Lithium cells with solid cathodes
b. Lithium cells with liquid cathodes
Lithium cells with solid cathodes: These systems may have solid or liquid electrolyte
Lithium manganese dioxide is widely used as a solid cathode lithium cell
Cathode is Mno2 & anode is lithium
Mno2 should be heated above 3000C to remove water
This i very important to improve the efficiency of the cell
Anodic and cathodic reactions are as follows
Anodic reaction: Li
Li+ + e-
Cahodic reaction: Li+ + xe- + Mno2
Net reaction: Li + Mno2
LiMno2
LiMno2
The electrolyte is a mixture of propylene carbonate and 1,2 – dimethoxy ethane
Applications
1. The cylindrical cells are used in fully automatic cameras
2. The coin cells are widely used in electronic device such as calculators and watches
Lithium cell with liquid cathode
Eg: Lithium sulphur dioxide cell
The co-solvents used are either acrylonitile propylene carbonate or mixture with of the two or 50% by
volume of so2
Cell reaction represents as follows
2Li + 2So2
LiS2o4
The second type of this cathode is Lithium Thionyl Chloride cells
It consists of a high – surface area carbon cathode
Thionyl chloride acts as the electrolyte solvent and the active cathode the reactions of cell represented as
follows
Anodic reaction: Li
Li + e
Cathodic reaction: 4Li + 4e- + 2 Socl2
Net reaction: 4Li + 2Socl2
4Licl + So2 + S
4Licl + So2 + S
Due to the nature of reaction Li – thionyl chloride cells possess very high energy density
Reactions involve two electrons per thionyl chloride
The So2 liberated as product is a liquid under the internal pressure of the cell
No- co solvent is required for the solution as thinly chloride is a liquid having moderate vapour pressure
Applications
1. Li-Socl2 cells are used on electronic circuit bonds for supplying fixed voltage for memory
prediction
2. These cells are also used for military and space applications
3. These cells are also used in medical device such as neuro – stimulators, drug delivery systems
Secondary cell or Storage cells:
The cells in which the cell reaction is reversed by passing direct current in opposite
direction.These cells can be used through a large number of cycles of charging and discharging
Eg: Lead acid storage cell
Nickel – cadmium cell
Lead acid storage cells
Storage cells can operate both as voltaic cell and as an electrolytic cell
When operating it act as voltaic cell it supplies electrical energy
At the time of recharging it acts as electrolyte cell
This storage cell has advantages of working both as an electrolytic cell and voltaic cell
Lead acid storage cell consists of lead anode and lead dioxide (Pho2) is cathode
Cathode is made of paste which passed in a grid made of lead
A number of (-ve plates) are connected in parallel the number of (+ve plates) are also connected in
parallel
The lead plates fit in between the lead dioxide plates
The plates are adjacent one by insulators like strips of wood, rubber or glass fiber
The entire combination is immersed in 20-21% dil H2so4
At anode oxidation reaction takes place
Pb
Pb+2 + 2e- (at lead anode)
The Pb+2 combine with So42- ion to produce Pbso4
Pb2+ + So42-
pbso4
The 2e- released at anode flows to pbo2 electrode and causes reduction of pbo2 to produce pb2+ which
finally combine with so42- to produce pbso4ca+(anode)
Pbo2 + 2H+ + 2ePb2+ + so42-
pb2+ + 2H2o
pbSo4
A lead accumulator for car consists of six lead aci storage cells in seris nd it is capable of
delivering 12V
The net reaction is
Pb + pbo2 + 4H+ + 2so42-
2pbso4 + 2H2o + energy
RE CHARGING:
During charging /recharging of the battery an external emf greater than 2 Volts so the cell reactions
are reversed as follows
Pbso4 + 2ePbso4 + sH2o + 2e
The net reaction is
pb + so42- (reaction at cathode)
pbo2 (reaction at anode)
2pbso4 + 2H2o + energy
pb + pbo2 + 4H+ + 2so42-
The positive pole of generator is attached to +ve pole of the battery
The negative pole of generator is attached to –ve pole of the battery
Battery containers are made of hard rubber and plastic
Eg: Six lead acid storage cells are in series are used
Applications:
It is used to supply current for electrical vehicles, gas engine ignition, mines, electrical trainee,
laboratories, hospitals, power stations
Advantages
1. It has relatively constant potential i.e. 12 volts
2. It is portable and inexpensive
3. It has features of both voltaic and electrolytic
Nickel cadmium cell/ Ni-cd cell
This battery consists of a cadmium anode and a cathode composed of a paste of Nio(oH)(s)
The cell reaction is as follows
Anode: cd(s) + 2oH
Cathode: 2Nio(oH) + 2H2o + 2eNet reaction: 2Ni(oH) + cd + 2H2o
cd(oH)2 + 2e2Ni(oH)2 + -oH
cd(oH)2 + 2Ni(oH)2
Recharging Ni-cd battery
The reaction can be easily reversed because the reaction products Ni(oH)2 and cd(oH)2 where to the
electrode surface
Ni-cd cell give a voltage 1.4 volts
There are two types of Ni-cd batteries
1. Pocket plate Ni-cd Battery: The materials used for cathode and anode are filled into pocket plates
(nickel coated steel strips) and connected in series
These cells long life > 20years with capacities between 10 to 1000 amperes
Voltage is 1.4 – 1.45 volts
Sintered – Plate Ni-cd battery
The anode materials fine Ni sintered in a mould around a nickel screen impregnated with
nickel nitrate and processed to produce nickel hydrate in pores
For cathode the molds are impregnated with cadmium salt to get hydrated oxides inside the pores
Electrolytes are assembled with suitably placed separators suh as porous polymer membrane
Electrolyte solution is KoH with1.24 to 1.3 specific gravity
Oxygen liberated during electrolysis, liberated through the porous membrane to cd cathode to
produce cd(oH)2 which helps in charging
These are made in different designs, plate’s horizontal disc & cylindrical cell
Advantages of Ni-cd batteries
1. The water consumption and float charge currents are exteremly low
2. These batteries are suitable to very rate discharge and low temperature operations
3. They have long life without any maintenance
Applications of Ni-cd batteries
The Ni-cd batteries are used for air – craft and diesel engine staring, lightening of trains, emergency
power and many military applications
Fuel cells
A fuel cell is an electrochemical cell coverts chemical energy containing fuel oxident systems into
electrical energy
The only one difference compared with other cells is chemical energy is provided by a fuel and an
oxidant stored outside the cell.
Fuel cell block diagram
Fuel and oxidising agents undergo the reactions and supplied separately to electrodes
Fuels are characterised by
1.
2.
3.
4.
High efficiency
Low noise levels
Free from vibration, heat transfer & thermal pollution
Built in wide range of power requirements
Hydrogen – Oxygen fuel cell
Best example for fuel cell is hydrogen oxygen cell
The cell consists of two inert porous electrodes
Electrolyte solution is 2.5% KoH
Electrodes are made up by graphite with finely divided platinium or 75/25 alloy of pb with Ag or Ni
Through aode hydrogen gas is bubbled
Through cathode oxygen gas is bubbled
sThe following cell reactions takes place
At anode: 2H2(g) + 4oH-
4H2o + 4e-
At cathode: o2(g) + 2H2o + 4eNet reaction: 2H2 + o2
4-oH(aq)
2H2o
The product discharged is water and the standard emf of the cell E0 = 1.23 volts
Number of such fuels is stacked together in series
Applications
They are used as auxiliary energy source in space vehicle, submarines or other military vehicles
Because of light weight these fuels are used for space crafts and H2o is valuable fresh water ie
astronauts.
Advantages
Energy conversion is high (75-82%)
Product H2o is used as drinking water for astronauts
Noise and thermal pollution is low
Fuel cells offer excellent method for fossil fuels
Maintenance is low
Limitations
Life times of fuel cells are not known
Initial cost is high
Distribution of hydrogen is not proper.
Fill in the blanks
1) Pure water does not conduct electricity because; it is___________
2) Calomel electrode is reversible with respect to ____________
3) A storage cell is a device that can operate ___________.
4) The cathode of ni-cd battery is composed of ____________.
5) A fuel cell converts a chemical energy of fuel directly to ________
6) A galvanic cell converts chemical energy into ____________.
7) Calomel electrode is constructed using a solution of ___________.
8) When storage cell is operating as voltaic cell it is said to be___________.
9) The resistance of a metallic conductor ______________ as the temperature is increased.
10) Graphite is a _________ conductor.
11) On dilution the specific conductivity of an electrode _____________
12) The units of equivalent conductance is ____________.
13) the total conductance of one gram equivalent of an electrolyte at a given dilution is called
____________.
14) When hydrogen is used as fuel in H-O fuel cell ,the electrode are made of
____________
15) In electro chemical series the elements are arranged in the __________ order of standard reduction
potentials
16) Ionization of an electrolyte in aqueous solution is due to ________
17) The specific conductance of a solution increases with _____________
18) NH4OH is a __________ electrolyte.
19) Speed ratio of the cat ion and anion is given by ___________
20) The standard electrode potential of saturated calomel electrode at 25c is _________________.
21) Nernst equation for electrode reaction is _________
22) __________ of the cells which do not store energy.
23) The substance which conduct electricity without decomposition is called ____________
24) The ionization of a strong electrolyte increases when the solution is diluted, the relation is given by
________
25) Acidic is a weak electrolyte because _____________
Multiple choice Questions
1) An ionizing solvent has
a)low value of dielectric constant b)a dielectric constant equal to one c) a high melting point d)none
2) Which of the following is a weak electrolyte?
a) NH4OH b) NaOH C) HCl d) NaCl
3) Calomel is
a) Mercuric sulphide b) mercurous sulphate c) an electrical cell d) none
4) Calomel electrode is constructed using a solution of
a) Sat KCl b) sat CaCl2 c) SatNH4Cl d)sat NaCl
5) Solid NaCl is a bad conductor of electricity because
a) It contains only molecules b) ions are not free to move c) it does not posses ions d)it does not contain
free molecules
6) The potential of std hydrogen electrode dipped in a solution of 1M concentration and hydrogen gas is
passed at 1atm pressure
a)1volt b)10 volt c)0 volt 100 volt
7) An electrochemical cell connected in series, that can be used as a source at a constant voltage is called
a) Battery b) voltaic cell c) electrolytic cell d) metal conductor
8) The standard reduction potential at 298K FOR Zn, Cr, H, Fe are 0.76V -0.74 -0.0V 0.77 respectively
the strongest reducing agent among there is
a) H b) Cr c) Zn d) Fe
9) The molar conductivities of a solution of any electrolyte is product obtained by multiplying
a) Specific conductivity with molecular weight
b) Specific conductivity with the volume of the solution containing 1gm of electrolyte
c) Reciprocal of conductivity with volume
d) Specific conductivity with equivalent weight
10) In lead-acid storage cell during discharging operation the concentration of H2SO4
a) Increases
b) decreases
SUBJECTIVE – QUESTIONS
1) What is a metallic conductor?
2) What are the units of specific conductivity?
3) What is specific conductance?
4) What is equivalent conductance?
5) Define molar conductivity
6) What is the unit of molar conductivity?
7) What is the unit of equivalent conductivity?
8) How do you measure equivalent conductivity?
9) What are conduct metric titrations? Give their significance?
10) State and explain Kohlrausch’s law of ionic mobilities.What is application of the law?
11) What are single electrode potentials? How do you determine the electrode potential Zn/ZnSO4?
12) Explain the standard electrode potential by taking calomel electrode as an example?
13) What are reference electodes? Explain the application of quinhydrone electrode?
14) What are batteries? How are they classified?
15) Explain the relative speeds of ions during electrolysis?
PRIMARY VOLTAIC CELL
PRIMARY VOLTAIC CELL (THE DRY CELL)
In this cell, once the chemicals have been consumed, further reaction is not possible. It
cannot be regenerated by reversing the current flow through the cell using an external direct
current source of electrical energy. The most common example of this type is dry cell.
The container of the dry cell is made of zinc which also serves as one of the electrodes. The
other electrode is a carbon rod in the centre of the cell. The zinc container is lined with a
porous paper. A moist mixture of ammonium chloride, manganese dioxide, zinc chloride and
a porous inert filler occupy the space between the paper lined zinc container and the carbon
rod. The cell is sealed with a material like wax.
As the cell operates, the zinc is oxidised to Zn2+
Zn ---> Zn2+ + 2e-
(Anode reaction)
The electrons are utilized at carbon rod (cathode) as the ammonium ions are reduced.
2NH4++2e- --> 2NH3 + H2
(Cathode reaction)
The cell reaction is
Zn+ 2 NH4+ ---> Zn2+ + 2NH3 + H2
Hydrogen is oxidized by MnO2 in the cell.
2MnO2 + H2 ---> 2MnO(OH)
Ammonia produced at cathode combines with zinc ions to form complex ion.
Zn2+ + 4NH3 ---> [Zn(NH3)4]2+
Ecell is 1.6 volt
Alkaline dry cell is similar to ordinary dry cell. It contains potassium hydroxide. The reaction
in alkaline dry cell are:
Zn + 2OH- ---> Zn(OH)2 + 2e-
(Anode reaction)
2MnO2 + 2H2O + 2e- ---> 2MnO(OH) + 2OH-
(Cathode reaction)
Zn + 2MnO2 + 2H2O ---> Zn(OH)2 + 2MnO(OH)
(Overall)
Ecell is 1.5 volt.
SECODARY VOLTAIC CELL (LEAD STORAGE BATTERY)
The cell in which original reactants are regenerated by passing direct current from external
source, i.e., it is re-charged, is called secondary cell. Lead storage battery is the example of
this type.
It consists of a group of lead plates bearing compressed spongy lead, alternating with a group
of lead plates bearing leaf dioxide, PbO2. These plates are immersed in a solution of about
30% H2SO4. When the cell discharge; it operates as a voltaic cell. The spongy lead is
oxidized to Pb2+ ions and lead plates acquire a negative charge.
Pb --> Pb2+ + 2e-
(Anode reaction)
Pb2+ ions combine with sulphate ions to form insoluble lead sulphate, PbSO4, which begins
to coat lead electrode.
Pb2+ + SO42- ---> PbSO4
(Precipitation)
The electrons are utilized at PbO2 electrode.
PbO2 + 4H+ + 2e- ---> Pb2+ 2H2O
Pb2+ + SO42- ---> PbSO4
(Cathode reaction)
(Precipitation)
Overall cell reaction is:
Pb + PbO2 + 4H+ + 2 SO42- ---> 2PbSO4 + 2H2O
Ecell is 2.041 volt.
When a potential slightly greater than the potential of battery is applied, the battery can be recharged.
2PbSO4 + 2H2O ---> Pb + PbO2 + 2H2SO4
After many repeated charge-discharge cycles, some of the lead sulphate falls to the bottom of
the container, the sulphuric acid concentration remains low and the battery cannot be
recharged fully.
FUEL CELL
Fuel cells are another means by which chemical energy may be converted into electrical
energy. The main disadvantage of a primary cell is that it can deliver current for a short
period only. This is due to the fact that the quantity of oxidising agent and reducing agent is
limited. But the energy can be obtained indefinitely from a fuel cell as long as the outside
supply of fuel is maintained. One of the examples is the hydrogen-oxygen fuel cell. The cell
consists of three compartments separated by a porous electrode. Hydrogen gas is introduced
into one compartment and oxygen gas is fed into another compartment. These gases then
diffuse slowly through the electrodes and react with an electrolyte that is in the central
compartment. The electrodes are made of porous carbon and the electrolyte is a resin
containing concentrated aqueous sodium hydroxide solution. Hydrogen is oxidised at anode
and oxygen is reduced at cathode. The overall cell reaction produces water. The reactions
which occur are:
Anode [H2(g) + 2OH-(aq)
Cathode
--->
2H2O(l) + 2e-] × 2
O2(g) + 2H2O(l) + 4e- ---> 4OH-(aq)
------------------------------------------------------------Overall 2H2(g) + O2(g) ---> 2H20(l)
This type of cells are used in space-crafts. Fuel cells are efficient and pollution free.
Concentration Cells
CONCENTRATIONS CELLS
If two plates of the same metal are dipped separately into two solutions of the same
electrolyte and are connected with a salt bridge, the whole arrangement is found to act as a
galvanic cell. In general, there are two types of concentration cells:
(
Example 43. A cell contains two hydrogen electrodes. The negative electrode is in contact
with a solution of 10-6 M hydrogen ions. The emf of the cell is 0.118volt at 25° C. Calculate
the concentration of hydrogen ions at the positive electrode.
Solution:
The cell may be represented as
Pt|H2(1 atm)|H+||H+|H2(1 atm)|Pt
10-6 M CM
Anode
Cathode
(-ve)
(+ve)
H2 ---> 2H+ + 2e-
2H+ + 2 ---> H2
Ecell = 0.0591/2 log([H+ ]cathode2)/[10-6 ]2
0.081 = (0.0591) log ([H+])/10-6
log[H+ ]cathode/10-6 =0.118/0.0591=2
[H+ ]cathode/10-6 = 102
[H+]cathode = 10-6 = 10-4 M
Solved Examples On Electrochemistry
Example 1.
Solution:
Find the charge in coulomb on 1 g-ion of
Charge on one ion of N3= 3 × 1.6 × 10-19 coulomb
Thus, charge on one g-ion of N3= 3 × 1.6 10-19 × 6.02 × 1023
= 2.89 × 105 coulomb
Example 2.
How much charge is required to reduce (a) 1 mole of Al3+ to Al and (b)1
mole of to Mn2+?
Solution:
(a) The reduction reaction is
Al3+
+ 3e- --> Al
1 mole
3 mole
Thus, 3 mole of electrons are needed to reduce 1 mole of Al3+.
Q=3×F
= 3 × 96500 = 289500 coulomb
(b) The reduction is
Mn4- + 8H+ 5e- --> MN2+ + 4H2O
1 mole
5 mole
Q=5×F
= 5 × 96500 = 48500 coulomb
Example 3.
How much electric charge is required to oxidise (a) 1 mole of H2O to
O2 and (b)1 mole of FeO to Fe2O3?
Solution:
(a) The oxidation reaction is
H2O --> 1/2 O2 + 2H+ + 2e1 mole
2 mole
Q=2×F
= 2 × 96500=193000 coulomb
(b) The oxidation reaction is
FeO + 1/2 H2O --> 1/2 Fe2O3 + H++eQ = F = 96500 coulomb