Chemistry unit 2

2. CELLS
Electrochemistry: It is a branch of chemistry which deals with the study of
transformation of chemical energy into electrical energy and vice versa”
Electrochemical cell and Classification with examples.
An electrochemical cell is a device, which is used to convert chemical energy
into electrical energy and vice versa.
These electrochemical cells are classified into two types as follows’
1) Galvanic or Voltaic cells: These are the electrochemical cells, which
converts chemical energy into electrical energy.
Ex. Daniel cell, Dry cell, etc
2)Electrolytic cells-are devices which convert electrical energy into
chemical energy.
Example: Electrolysis of molten NaCl, Recharge process of lead acid battery

3. The Nernst Equation
E cell = E 0
cell - (RT/nF) lnQ
Ecell = cell potential under nonstandard conditions (V)
Ecell = cell potential under nonstandard conditions (V)
E0
cell = cell potential under standard conditions
R = gas constant, which is 8.31 (volt-coulomb)/(mol-K)
T = temperature (K)
n = number of moles of electrons exchanged in the
electrochemical reaction (mol)
F = Faraday's constant, 96500 coulombs/mol
Q = reaction quotient, which is the equilibrium expression with
initial concentrations rather than equilibrium concentrations

4. Types of Electrodes
 Reference electrodes
 Indicator electrode

5. Reference Electrodes:
• Reference electrode: “Reference electrode are the electrode with
reference to those, the electrode potential of any electrode can be
measured.” It can acts both as an anode or cathode depending upon the
nature of other electrode.
• Reference electrode is defined as the electrode which has stable and
reproducible potential and complete the cell acting as half cell. The criteria
for an electrode to act as reference electrode
• 1. The potential of such electrode should be known
• 2. The potential show minimum variation
The Reference Electrodes can be classified in to two types
i) Primary reference electrodes Ex: Standard hydrogen electrode
ii)Secondary reference electrodes Ex: Calomel and Ag/Agcl
electrodes

6. Reference electrodes
A) Primary reference electrode
e.g. Hydrogen electrode
B) Secondary reference electrode
e.g. Calomel electrode, silver-silver chloride
electrode, mercury-mercury sulphate electrode,
mercuric oxide electrode, glass electrode,
quinhydrone electrode.

7. Secondary Reference Electrode
Calomel Electrode:
Principle: The potential of the calomel electrode depends
upon the concentration of KCl. If KCl solution is saturated,
then its potential is0.2415V and such electrode is called
saturated calomel electrode(SCE).
If KCl solution is 1N,then its potential is 0.280V and such
electrode is called normal calomel electrode(NCE).If KCl
solution is0.1N,then its potential is 0.3338V and such electrode
is called decinormal calomel electrode.(DNCE).
Representation of calomel electrode:
Hg|Hg2Cl2.KCl(xM)

8. Fig: Calomel electrode
Reactions:-
At Anode :- 2Hg → 2Hg+ + 2e-
At cathode: 2Hg+ + 2Cl- + 2e- → Hg2Cl2
------------------------------------------
2Hg + 2Cl- → Hg2Cl2

9. The calomel electrode merits and demerits
Merits
• It is easy to construct and transport and
convenient to handle.
• The potential of the electrode is reproducible and
remains constant
• No separate salt bridge is required for its
combination with other electrode
Demerits
• It should be used above 50 oC Hg2Cl2 starts
decomposing
• It involves handling of poisonous Hg and Hg2Cl2

10. Concentration
of KCl
E0
[v]
saturated KCl 0.241
1M KCl 0.281
0.1M KCl 0.33
4
MEASUREMENT OF ELECTRODE POTENTIAL USING CALOMEL ELECTRODE:
Electrode potential of a given electrode can be measured by using calomel electrode as a
reference electrode.
Example: To measure the electrode potential of zinc, Zinc electrode is coupled with SCE. So zinc
acts as anode and SCE acts as cathode
Ec e l l  Ec a t h o d e  Ea n o d e
Ea n o d e  Ec a t h o d e  Ec e l l
 0 . 2 4 1  Ec e l l
E 2 
Z n / Zn
Applications:
1.It is used as secondary reference electrode in the measurement of single electrode.
2.It is used as reference electrode in all potentiometer determinations and to
measure pH of the given solution

11. T h e n e t cel l r ev er s i b l e el e ct r o d er e a ct i o n i s ,
E  E
0

2 . 3 0 3 R T
logCl  , w h e r e n  1
A g C l ( s )  2 e -  A g  C l -
 n F
 E  E 0
0.0591logCl 
at 2 9 8 K
 Therefore electrode potential of calomel electrode is
depending upon the concentration of KCl. The
electrode is reversible with chloride ions.
 Applications:
Used as secondary reference electrode in ion selective elctrode.

12. Indicator Electrode (Glass Electrode)
The glass electrode consists of a very thin walled glass bulb,
made from a low melting glass having high electrical
conductivity, blown at the end of a glass tube as shown in fig.
The bulb contains 1M HCl solution sealed into glass tube is a
silver wire coated with silver chloride at its lower end. The
lower end of the silver wire dips into the HCl, forming silver-
silver chloride electrode.
When glass electrode is placed in a solution the potential
develops across the glass membrane as a result of a
concentration difference of H+ ions on the two side of the
membrane. The potential of a glass electrode is determined
using standard calomel electrode as shown in fig (a) & (b).

13. Principle
 When two solution of different (H+) are separated by a
thin glass membrane, a potential difference is
developed at the two surface of membrane .The
potential difference developed is proportional to the
difference in H+ of the solution.
 The glass membrane acts as ion exchange i.e. exchange
of Na + of glass with H+

14. Fig (a): Indicator electrode Fig (b):Glass electrode connect calomel electrode

15. Determination of pH using glass electrode:
0.0591
E 0
G  Ec e l l  ES C E
pH 
Ecell  Ecathode  Eanode
Ec e l l  EG  ES C E
Since E SCE is knowing e m f the cell,
E glass can be evaluated.
EG  E 0
G 0.0591 pH
Ecell E 0
G 0.0591 pH  ES C E

16. Ion selective electrodes (ISE)
An ion selective electrode is an indicator electrode which produces
a potential when it is placed in a solution containing a certain ion.
In this electrode, a suitable non porous membrane separate two
solutions, containing similar ions of different concentration and
acts as an electrochemical membrane. The remarkable property of
such electrode is that a potential difference is developed on either
side of the membrane which is proportional to the concentration
difference.
“ Ion selective electrode is one which selectively responds to a
specific ion in a mixture and the potential developed at the
electrode is a function of the concentration of that ion in the
solution”

17. ISE
• Also known as indicator electrodes
• Respond directly to the analyte
• Used for direct potentiometric measurements
• Selectively binds and measures the activity of one
ion (no redox chemistry)

18. ISE have many advantage
 Relatively inexpensive and simple to use
 Robust and durable in field and laboratory
environment .
 Can be used for colored, turbid samples
 Can be used for wide range of temperature.
 Mechanical strong
 Resistant to chemical attack
 Resistant to solvent attack

19. • An ion-selective electrode (ISE), also known as
a specific ion electrode (SIE), is a transducer
(or sensor) that converts the activity of a specific
ion dissolved in a solution into an electrical
potential, which can be measured by a voltmeter
or pH meter.
• An ideal I.S.E. consists of a thin membrane across
which only the intended ion can be transported.
• The transport of ions from a high conc. to a low one
through a selective binding with some sites within
the membrane creates a potential difference.

20. Types of Ion selective electrodes
Glass membrane electrode- H+ , Na+, Ag +, K +
Liquid membrane electrode-Ca
Solid-state electrode-F-
Gas sensing electrode-H2S,NO2,CO2
Enzyme electrodes-NH3,amines

21. GLASS MEMBRANE ELECTRODE
• Glass electrode are responsive to
univalent cations ( H+ , Na+)
• The selectivity for this cation by varying the composition
of a thin ion sensitive glass membrane.
• Example: pH electrodeused for pH measurement
-used as a transducer in various gas and
biocatalytic sensor, involving proton generating or
consuming reaction.

22. • Glass membrane manufactured from SiO2 with
negatively charged oxygen atom.
• Inside the glass bulb, a dilute HCl solution and
silver wire coated with a layer of silver chloride.
• The electrode is immersed in the solution and pH is
measured.

23. e
Construction and working of Ion selective electrod
(ISE) :
Glass electrode: A glass electrode is an ion selective electrode where
potential depends upon the pH of the medium.
1. The glass electrode consists of a glass bulb made up of special type of
glass (sodium silicate type of glass) with high electrical conductance.
2. The glass bulb is filled with a solution of constant pH (0.1MHCl) and
insert with a Ag-AgCl electrode, which is the Internal reference
electrode and also serves for the external electrical contact.
3. The electrode dipped in a solution containing H+ ions as shown in the
figure.
The electrode representation is,
Glass | 0.1M HCl | Ag-AgCl

24. • Ion Selective Electrodes (including the most common
pH electrode) work on the basic principal of the
galvanic cell .By measuring the electric potential
generated across a membrane by "selected" ions, and
comparing it to a reference electrode, a net charge is
determined. The strength of this charge is directly
proportional to the concentration of the selected ion. The
basic formula is given for the galvanic cell:
• Ecell = EISE - ERef

25. WORKING:
• The glass electrode works on the principle that when a thin
glass membrane is placed between two different concentration
of a solution, a boundary potential Eb is developed at layers of
the glass membrane. This potential arises due to difference in
the concentration of H+ ion inside and outside themembrane.
External Solution glass membrane Internal solution
C2=[H+]
E2 Eb
C1=[CONSTANT]=k
E1
• Boundary potential, Eb = E2 – E1

26. Advantages
• This electrode can be used to determine PH in the range 0-
9, with special type of glass even up to 12 can be
calculated.
• It can be used even in the case of strong oxidizing agents.
• The equilibrium is reached quickly.
• It is simple to operate, hence extensively used in various
laboratories.
Limitations
• The glass membrane though it is very thin, it offers high
resistance. Therefore ordinary potentiometers cannot be
used; hence it is necessary to use electronic
potentiometers.
• This electrode cannot be used to determine the PH above 12

27. SOLID STATE ELECTRODE
• Solid state electrode are selective primarily to
anions.
• It may be a homogenous membrane electrode
or heterogeneous membrane electrode.
• Homogenous membrane electrode: ion-selective
electrodes in which the membrane is a crystalline
material (AgI/Ag2S).

28. Homogenous membrane electrode
• homogenous electrode is made up of LaF3 for
determination of F- in water and doped with
Europium flouride
• LaF3 (s) La F2 + F-
• This leads to separation of charge and equilibrium
is established ,leads to potential.

29. Homogenous membrane electrode:
The membrane is made from Lanthanum trifluoride(LaF3)crystal
doped with europium fluoride(EuF2)
The crystal is sealed at the bottom of the polymer containing
internal reference solution (NaF +NaCl or KF or Kcl) consisting
of a reference electrode

30. Determination of Fluoride Ion
• In the lanthanum fluoride electrode, the sensing
element is a crystal of lanthanum fluoride LaF3,
doped with europium fluoride EuF2 to
create lattice vacancies. Such a crystal is an ionic
conductor by virtue of the mobility of fluoride ions
which jump between lattice vacancies.
• An electrochemical cell may be constructed using such
a crystal as amembrane separating two fluoride
solutions.This cell actsas a concentration cell with
transference where the fluoride transport number is 1.
As transference of charge through the crystal is almost
exclusively due to fluoride, the electrode is highly
specific to fluoride

31. Working :
1. EuF2 produces holes in the crystal lattice of Laf3 through
which F- ions can pass.
2. When the electrode is in contact with the sample solution a
potential develops across the membrane which depends on
the difference F- concentration since the concentration on
the F- in the internal solution is fixed the potential developed
across the membrane is related to F- concentration

32. Heterogeneous membrane electrode:
Heterogeneous membrane electrode consisting
of solid crystalline material (AgI/Ag2S)incorporated
with polymer like PVC or silicon
When the electrode membrane is in contact with a
solution containing chloride ions an eletrode
potential develops. This potential is measured
against constant reference potential.

33. LIQUID MEMBRANE
ELECTRODE
• Liquid membrane is a type of ISE based on
water-immiscible liquid substances produced in
a polymeric membrane used for direct
potentiometric measurement.
• Used for direct measurement of several polyvalent
cations (Ca ion) as well as a certain anions.

34. • The polymeric membrane made of PVC to
separate the test solution from its inner
compartment.
• Contains standard solution of the target ion.
• The filling solution contains a chloride salt for
establishing the potential of the internal Ag/AgCl
wire electrode.

35. GAS SENSING ELECTRODE
• Available for the measurement of ammonia,
carbon dioxide and nitrogen oxide. A nitrate ion
responsive electrode is for NO2 while sulphide
ion selective electrode for H2S
• This type of electrode consist of permeable membrane
and an internal buffer solution.
• The pH of the buffer changes as the gas react with it.
• The gas permeable membrane is made of a
hydrophobic porous polymer. The gas in the test
solution diffuses through the membrane and reacts
with the internal filling solution to form the ions.

36. Working:
• The electrode does not detect the presence of
molecular gas but rather an ion into which the gas is
converted after it passes through the membrane.
• The gas in the test solution diffuses through the
membrane and reacts with the internal filling solution
to form the ion. These ions are detected using gas
sensing electrode

37. Enzyme based membrane
These electrode use enzyme to convert
substance in the solution into ionic products
which are measured using ion selective
electrode. The enzyme is immobilized at the
surface of the electrode

38. Enzyme based membrane
Working:
When the electrode is immersed into a solution
containing urea,NH4+ ios are produced which
through the gel.
CO(NH2)2 +H2O+2H+ 2NH4 +CO2
The boundary potential is developed due to
difference in concentration of NH4 on either side
of the membrane
The potential developed is measured using a glass
electrode as reference electrode

39. Advantages of IonSelectiveElectrode
(ISE) Technique
• When compared to many other analytical techniques,
Ion-Selective Electrodes are relatively inexpensive and
simple to use and have an extremely wide range of
applications and wide concentration range.
• Under the most favorable conditions, when measuring
ions in relatively dilute aqueous solutions and where
interfering ions are not a problem, they can be used
very rapidly and easily.
• They are particularly useful in applications where only
an order of magnitude concentration is required, or it
is only necessary to know that a particular ion is below
a certain concentration level.

40. • They are invaluable for the continuous monitoring
of changes in concentration for example in
potentiometric titrations or monitoring the uptake of
nutrients, or the consumption of reagents.
• They are particularly useful in biological/medical
applications because they measure the activity of
the ion directly, rather than the concentration.
• ISEs are one of the few techniques which can
measure both positive and negative ions.
• They are unaffected by sample colour or turbidity.
• ISEs can be used in aqueous solutions over a
wide temperature range. Crystal membranes can
operate in the range 0 C to 80 C and plastic
membranes from 0 C to 50 C.

41. LIMITATION
• Precision is rarely better than 1%.
• Electrodes can be fouled by proteins or other organic
solutes.
• Interference by other ions.
• Electrodes are fragile and have limited shelf life.
• Electrodes respond to the activity of uncomplexed ion.
So ligands must be absent.

42. APPLICATION
 Ion-selective electrodes are used in a wide variety of applications for
determining the concentrations of various ions in aqueous solutions. The
following is a list of some of the main areas in which ISEs have been used.
 Pollution Monitoring: CN, F, S, Cl, NO3 etc., in effluents, and naturalwaters.
 Agriculture: NO3, Cl, NH4, K, Ca, I, CN in soils, plant material, fertilisersand
feedstuffs.
 Food Processing: NO3, NO2 in meat preservatives.
 Salt content of meat, fish, dairy products, fruit juices, brewing solutions.
 F in drinking water and other drinks.
 K in fruit juices and wine making.
 Corrosive effect of NO3 in canned food
 F in skeletal and dental studies.

43. ELECTROCHEMICAL METHODS
Electrochemical methods are analytical techniques that use a
measurement of potential, charge or current to determine an
analyte concentration or to characterize an analytes chemical
techniques.
These methods are divided into 5 major groups
Potentiometry
Voltametry
Coulometry
Conductometry
Dielectrometry

44. CONDUCTOMETRY
Principle
The ability of any ion to transport charge
depends on the mobility of the ion, mobility of ion is
affected by factors like the charge on ion, size and
mass of ion and extent of solvation.

45. Important laws, used in conductometry
 Ohm’s Law : It is written as, I α E
Unit = Ω(ohm).
 Conductance (c): It is written as C = 1/R,
Unit = ohm or mho (ohm-1)
 Specific resistance (ρ): It is written as R α l/a or R=ρ.l/a
Unit = ρ is ohm.cm
 Specific conductance (k): It is written as K =C.l/a
Unit = mho.cm-1
 Cell constant: It is written as K = C.(l/a)
Unit= cm-1
 Equivalent conductance (λv): λv = K x 1000/C
Unit = λv = ohm-1.cm2.gm-equiv-1
 Molar conductance (µ):µ = µ µ = K x 1000/Molarity
Unit = mho.cm2.gmol-1

46. Important laws, Definitions used in conductometery
Based on the conductance of electrical current through
electrolyte solutions similar to metallic conductors.
ohms law : The electric conductance in accordance with ohms
law which states that the strength of current(I) passing through
conductor is directly proportional to potential difference and
inversely to resistance.
I=V/R
where I = strength of current
V= potential difference
R= resistance

47. DEFINITIONS AND RELATIONS
Ohms law: According to this law, the strength of current
(I) flowing through a conductor is directly proportional to
the potential difference (E) applied across the conductor
and inversely proportional to the resistance (R) of the
conductor.
I = E/R
Conductance: It implies the ease with which the current
flows through conductor, thus the conductance is
reciprocal to resistance.
C= I/R

48. Resistance refers to the opposition
to the flow of current.
For a conductor of uniform cross section(a)
and length(l); Resistance R,
a
l
l l
R l and R R
a a
    
Where is called resistivity or
specific resistance.

Resistance

49. Conductance
The reciprocal of the resistance is called conductance. It is denoted by C.
C=1/R
Conductors allows electric current to pass through them. Examples are metals,
aqueous solution of
acids, bases and salts etc.
Insulators do not allow the electric current to pass
through them.
Examples are pure water, urea, sugar etc.
Unit of conductance is ohm-1 or mho or Siemen(S)

50. Specific resistance: (ρ) is the resistance offered by a substance is
directly proportional to 1cm length and inversely proportional
to 1sq.cm cross sectional surface area
It is written as R α l/a or R=ρ.l/a
Unit of measurement is ohm cm.
Specific conductivity: (kv) is the conductivity offered by a
substance of 1cm length and 1sq.cm surface area,Unit of
measurement is mhos cmˉ1
Specific conductance (k): It is written as K =C.l/a
Unit = mho.cm-1

51. DEFINITIONS ANDRELATIONS
Specific resistance: (ρ) is the resistance offered by a substance is
directly proportional to 1cm length and inversely proportional
to 1sq.cm cross sectional surface area
It is written as R α l/a or R=ρ.l/a
Unit of measurement is ohm cm.
Specific conductivity: (kv) is the conductivity offered by a
substance of 1cm length and 1sq.cm surface area,Unit of
measurement is mhos cmˉ1
Specific conductance (k): It is written as K =C.l/a
Unit = mho.cm-1

52. Specific conductance
1
 

Specific Conductivity
 
  
 
K
a
x Conductance
Unit of specific conductance is ohm–1cm–1
SI Unit of specific conductance is Sm–1 where S is Siemen
a
But ρ = R
K
a.R
 
l/a is known as cell constant
Conductance of unit volume of
cell is specific conductance.

53. DEFINITIONS ANDRELATIONS

54. Molar and Equivalent conductance
Molar conductance: This may be defined as “the conductance of
a solution containing 1gm mole of electrolyte
It is denoted by µvand is measured in mhos
λv = K x 1000/M Unit = λv = ohm-1.cm2.per mole
Equivalent conductance (λv): Conductivity of solution by all
the ions produced by one gram equivalent in V volume
(v=1000/c)
λv = K x V
λv = K x 1000/C
Unit = λv = ohm-1.cm2.gm-equiv-1

55. DETERMINATION OF CELLCONSTANT
Cell constant :Cell constant is the ratio of the distance
between two electrodes and area of the electrode.
If two electrode 1 cm apart and having area A then l/a is fixed
R = ρx [where x=l/a= cell constant]
Specific conductance= Cell constant X Conductance
X = cell constant = R
ρ
cell constant=1observed conductivity
1specific conductivity
Specific conductivity= cell constant X observed conductivity

56. Determination of cell constant :-
It is calculated by following formula,

57. Variation of conductance with Dilution.
The conductivity of solution increases on dilution.
The specific conductivity decreases on dilution (as
number of ions decreases w.r.t. to volume).
The equivalent and molar conductivities increase with
dilution.
The equivalent and molar conductivities tend to
acquire maximum value with increasing dilution.
[Maximum at dilution]
Variation of molar conductance with concentration

58. Effect of Dilution on Conductivity
Specific conductivity decreases on dilution.
Equivalent and molar conductance both increase with dilution and reaches
a maximum value.
The conductance of all electrolytes increases with temperature.
concentration, (mole L )
–1 1/2
CH COOH (weak electrolyte)
3
KCl (strong electrolyte)

59. Factors Affecting Conductivity
Conductometric analysis is based on the measurement of the
electrical conductivity of the solution due to the mobility of
cations and anions towards respective electrodes.
The electrical conductivity is entirely due to the movement of
ions.
The various factors are:
1. Number of ions per ml: Greater the number of ions per ml in a
solution, greater is the specific resistance. At higher concentration
of solution, the number of ions per ml is higher.
2.Charge of ions: Higher the charge on ions, greater is the
conducting ability Eg:Mg++ has more conductivity than Na+,So2
has more conductivity than NO3-

60. The mobility of an ion:
Smaller the size of ion, greater is its mobility and
conducting ability.Eg H+ions have highest mobility
due to its smalllest size.NH4+ has lesser conductivitry
than Na+
Effect of temperature on conductivity :-
The conductance of the solution increases with
increase in temperature due to,
 Increase in the velocity of ions.
Decrease in the viscosity of the medium.
Decrease in the interaction between the ions.

61. Concentration:
At higher concentration, the degree of dissociation
is lower. Hence number of ions per ml is slightly
lower.
The specific conductivity decreases on dilution (as
number of ions decreases w.r.t. to volume).
The equivalent and molar conductivities increase
with dilution

62. Kohlarusch’s Law:-
The law states that at infinite dilution, each ion
migrates independently of its co-ions and contributes
definite share to the total equivalent conductance of
the electrolyte
The equivalent conductance of an electrolyte at infinite
dilution is equal to the sum of the equivalent
conductance of the component ions.
Mathematically it is written as, λ∞ =λ anion + λ cation

63. Application of Kohlarusch’s Law :-
Calculation of molar conductivity at infinite dilution
for weak electrolytes:
Calculation of degree of dissociation:
Calculation of dissociation constant for a weak
electrolyte:
Calculation of solubility of sparingly soluble salt:

65. CONDUCTOMETER

66. CONDUCTIVITY
METER
BURETTE
CONDUCTIVTY
CELL
MAGNETIC
STIRRER

67. ACID-BASE OR NEUTRAL TITRATIONS
STRONG ACID- STRONGBASE
• Eg. HCl vs NaOH
WEAK ACID – STRONGBASE
• Eg. CH3COOH vs NaOH
STRONG ACID – WEAKBASE
• Eg. HCl vs NH4OH
WEAK ACID – WEAKBASE
• Eg. CH3COOH vs NH4OH

68. Types of Conductometric Titrations:
 Titration of strong acid (HCl) with strong base (NaOH):
HCl + NaOH → NaCl + H2O

69. STRONG ACID- STRONG BASE
Fall in conductance due to replacement of high conductivity
hydrogen ions by poor conductivity of sodium ions.
Rise in conductance due to increase in hydroxyl ions.

70.  Titration of a weak acid (CH3COOH) and a strong
base (NaOH) :-
CH3COOH + NaOH → CH3COO-Na+ + H2O

71. WEAK ACID – STRONGBASE:
Initial decrease in conductance followed by increase due to
NaOH
Steep rise due to excess of NaOH

72. STRONG ACID – WEAKBASE
Fall in conductance due to replacement of hydrogen by
ammonium ions
Conductance remain constant due to suppression of NH4OH
by NH4Cl.

73. WEAK ACID – WEAKBASE
Increase in conductance due to excess of CH3COOH
Constant conductance due to suppression of NH4
OH by
CH3COOH

74. ADVANTAGES OF CONDUCTOMETRIC
TITRATIONS
Does not require indicators since change in conductance is
measured by conductometer.
Suitable for coloured solutions.
Since end point is determined by graphically means accurate
results are obtained with minimum error.
Used for analysis of turbid suspensions, weak acids, weak bases,
mix of weak and strong acids.
Temperature is maintained constant throughout the titration.
This method cab be used with much diluted solutions.
Be seen by eye.

75. DISADVANTAGES
Increased levels of salt in solutions masks the conductivity
changes in such cases it does not gives accurate results.
Applications of conductometric titrations to redox systems is
limited because, high concentration of hydronium ions in the
solutions tends to mask the change in conductance.

76. pH Metry :-
The concept of pH was first introduced by Danish chemist
Sorensen. The equation that defines pH is given as
follows:
pH= -log[H+]
which is read as: the pH is equal to minus the log of the H+
concentration.
For example :-
H+ concentration is very low, lets say about 0.0000001M,
then the pH is,
pH= -log[.0000001] which is the same as -log[1 X 10-7]
The term log [1 X 10-7] = -7
- (-7) = 7

77. Buffer solutions A buffer solution is one which maintains a fairly
constant PH even when small amount of acid or alkali is added
Types of Buffer solutions :
 Acidic buffer:
Acidic buffer solution contains equimolar quantities of a weak
acid and its salt with strong base. For example: acetic acid,
CH3COOH and sodium acetate I.e. CH3COONa. A solution
containing equimolar quantities of acetic acid and sodium acetate
maintains its pH value around 4.74.
 Basic buffer:
Basic buffer solution contains equimolar quantities of a weak
base and its salt with strong acid. For example: ammonium
hydroxide i.e. NH4OH and ammonium chloride I.e. NH4Cl. A
solution containing equimolar quantities of ammonium hydroxide
and ammonium chloride maintains its pH value around 9.25.

78. pH metric Titrations:-
Titration of StrongAcids (HCl + NaOH ) against
strong base (NaOH) :-
Principle:
That titration where the end point is measure
by change in pH is known as pH metric titration. On
addition of NaOH in mixture of HCl- NaOH
Reactions will occur:
HCl +NaOH→H2O+NaCl

79. Neutral point
The term "neutral point" is best avoided.
• The term "equivalence point" means that
the solutions have been mixed in exactly
the right proportions according to the
equation.
• The term "end point" is where the
indicator changes colour. As you will see
on the page about indicators, that isn't
necessarily exactly the same as the
equivalence point.

80. Titration curves for strong acid v strong base
All the following titration curves are based on both acid and alkali having
a concentration of 1 mol dm3. In each case, you start with 25 cm3 of
one of the solutions in the flask, and the other one in a burette.
Although you normally run the acid from a burette into the alkali in a
flask, you
may need to know about the titration curve for adding it the other way
around
as well. Alternative versions of the curves have been described in
most cases.
We'll take hydrochloric acid and sodium hydroxide as typical of a strong
acid and a strong base.
Running acid into the alkali
•You can see that the pH only falls a very
small amount until quite near the
equivalence point. Then there is a really
steep plunge.

81. Titration curves for strong acid v strong base
Running alkali into the acid
• This is very similar to the
previous curve except, of
course, that the pH starts off
low and increases as you
add more sodium hydroxide
solution.
• Again, the pH doesn't
change very much until you
get close to the equivalence
point. Then it surges
upwards very steeply.

82. Differential curve of titration

83. pH metric Titrations:-

84. Thank you