1) pH is a measure of how acidic or basic a solution is, defined as the negative logarithm of the hydrogen ion concentration. 2) A pH meter measures the pH of a solution by using a glass electrode and reference electrode to detect the voltage difference created by the solution's hydrogen ion concentration. 3) The pH scale ranges from 0-14, with values below 7 being acidic, above 7 being basic, and 7 being neutral. A change of one pH unit corresponds to a tenfold change in hydrogen ion concentration.

1. Presented By:
Vivek Chourasiya

2.  The hydrogen ion concentration or pH is a
measure of the acidity or alkalinity of a
solution.
 PH = 1
log 10 (H+)
 (H+) is the hydrogen ion concentration of the
solution in moles per liter.
 The pH of solution is defined as the negative
logarithm of the hydrogen ion concentration,
in an aqueous solution.
What is pH

3.  Scale of pH meter: 0 to 14 pH
 An acid solution. : Less than 7.0
 A basic solution. : greater than 7.0
 A neutral solution. : 7
 A change of one pH unit corresponds
to a 10
 old change of hydrogen-ion conc. of
the solution.
Scale of pH meter

4. pH scale

5.  When the pair of electrodes or a
combined electrode (glass electrode
and calomel electrode) is dipped in an
aqueous solution , a potential is
developed across the thin glass of the
bulb (of glass electrode).
 The e. m. f. of complete cell (E) formed
by the linking of these two electrodes at
a given solution temp. is therefore
Principle

6. E = Eref - Eglass
 Eref is the potential of the stable calomel
electrode which at normal room temp. is
+0.250V.
 Eglass is the potential of the glass electrode
which depends on the pH of the soln. under test.
 The resultant e.m.f. can be recorded
potentiometrically by using vacuum tube
amplifier.
 Variations of pH with E may be recorded directly
on the potentiometer scale graduated to read pH

7. Determination of pH

8.  Litmus paper is an indicator, which can be
used to determine acidic and basic
solutions.
 Normally this comes as a paper stripe.
Mixture of water-soluble dyes extracted
from lichens like Roccella tinctoria are
absorbed into filter paper stripes to make
litmus paper.
 In this mixture, there are about 10- 15 types
of dyes. There are two types of litmus
papers as blue and red.
pH strips

9.  Red litmus paper is used to test basic
solutions. Red litmus papers turn blue when
encountered with a basic solution.
 Where as blue litmus papers turn red when
encounter with an acidic solution.
 The neutral litmus papers are purple in color.
The color change of litmus papers is taking
place over pH range 4.5-8.3 at 25 °C.

10. Advantage
 Readings are instantaneous, and easy to take.
 Litmus papers can be used by anyone without
any expertise knowledge.
Disadvantage
 that they can not be used to determine the pH
value below or above the above mentioned
value.
 The strength of the acidity or the basicity also
cannot be determined using litmus paper.

11.  pH papers are indicators which are easy to
use. They are common in every chemical
laboratory.
 They come in various types as rolls,
stripes, etc.
 Though pH meter is the best equipment to
measure the pH value accurately, pH
papers are the best alternative when quick
and approximate measurements are
needed.
pH paper

12.  With the pH paper, a color chart is
provided. When the pH paper in the
solution, where pH should be
determined, the paper will show a
certain color.
 This color change is relevant to the
pH of the solution.

13.  Colorimetric Determination of pH is the variation
in the intensity of the colour of a solution with
changes in concentration (or pH).
 The colour may be due to an inherent property
of the constituent itself (e.g. MnO4 − is purple)
or it may be due to the formation of a coloured
comp suitable reagent (e.g. indicator).
 By comparing the intensity of the colour of a
solution of unknown concentration (or pH) with
the intensities of solutions of known
concentrations (or pH), the concentration of an
unknown solution may be determined.
Colorimetric Determination of the pH

14. 1. Estimate the pH of the unknown solution
using an universal indicator paper.
2. Select an indicator.
3. Select a buffer system.
4. Prepare a series of solutions with various pH
values (10 ml from each, pH steps in the
series are 0.2 - 0.2)
5. Add 1 or 2 drops (strictly the same amount)
of indicator to each of the solutions and
finally to the unknown solution. Compare the
colour of the unknown solution to the
Follow the following procedure

15.  A pH meter is an electronic instrument
used for measuring the pH (acidity or
alkalinity) of a liquid
 Special probes are sometimes used to
measure the pH of semi-solid substances.
 A typical pH meter consists of special
measuring probes (a glass electrode and
a reference electrode) connected to an
electronic meter that measures and
displays the pH reading.
pH meter

16.  The first commercial pH meter
was built around 1936 by
Radiometer in Denmark and by
Arnold Orville Beckman in the
United States.
History

17.  pH is defined as the negative logarithm of
hydrogen ion concentration.
pH= -log [ H+ ]
p = power
H = hydrogen
[H+ ] = hydrogen ion concentration
 The pH of a solution can be measured by
the pH meter.
 The glass electrode is an half cell and the
calomel electrode is another half cell.
Principle

18. The glass electrode contains Ag, AgCl, and HCL.
All these
remain in the ionized state.
Ag <—————> Ag+ + e–
AgCl <—————> Ag+ + Cl-
HCL <—————> H+ + Cl–
All the above three equilibrium reactions are
balanced. In
the glass electrode, H+ is generated.
Following reactions take place in
the electrodes

19. The calomel electrode contains Hg, Hg2Cl2 , and
KCL.
Here also the following series of ionizations take
place.
Hg <————> Hg+ + e-
Hg2Cl2 <————> 2Hg+ + 2Cl-
KCL <————> K+ + Cl-
The above equilibrium reactions are balanced. In
the
calomel electrode, H+ is not generated.

21. 1. Reference Electrode
 Standard Hydrogen
electrode
 Calomel electrode
 Silver-Silver Chloride
electrode
Important Components of pH
Meter

22.  Indicator electrode indicates
the potential or pH of a
solution in comparison to a
reference electrode of a
known potential.
INDICATOR ELECTRODES

23.  It consists of a very thin bulb about 0.1 mm thick
blown on to a hard glass tube of high resistance.
 The bulb contains 0.1 mol/litre HCL connected to
a platinum wire via a silver-silver chloride
combination.
Glass Electrode

24. a) General
I. electrodes based on determination of cations
or anions by the selective adsorption of these
ions to a membrane surface.
II. Often called Ion Selective Electrodes (ISE)
III. Desired properties of ISE’s
 minimal solubility – membrane will not dissolve
in solution during
 Measurement – silica, polymers, low solubility
inorganic compounds
‚Need some electrical conductivity
ION SELECTIVE ELECTRODE

26.  Response is very rapid
 Chemically resistant to oxidizing &
reducing agents, dissolved gases,
salts etc.
 When Lithia -silica glasses are used,
it can be used over the entire pH
range.
ADVANTAGES

27.  It is extremely fragile
Minute abrasions on the surface
of the tip, damages the electrode
 It cannot be used with simple
potentiometers, because of the
high resistance.
DISADVANTAGES

28.  it consists of a antimony rod dipped
into a solution, whose potential or pH
to be determined.
 Antimony oxide is formed on
exposure to air
Antimony – Antimony Oxide Electrode

29. ADVANTAGES
 It can be used from pH 3 to pH 8. it can be
used even up to pH 12
 It is not easily poisoned or damaged
 Can be used even with viscous fluid
DISADVANTAGES
 This electrode cannot be used in presence
of dissolved oxygen, oxidizing agents,
complexing agents etc..

30.  Which has a standard potential on
its own
 And its potential does not change
to whichever solution it is dipped.
 E.g. hydrogen electrode, saturated
calomel electrode & silver-silver
chloride electrode
 Most commonly used is saturated
calomel electrode
REFERENCE ELECTRODE

31.  A potential difference will be generated between the two
platinum electrodes by the different active hydrogen-ion
concentrations in the solutions. The relationship is
expressed by the NERNST equation:
where:
E = potential difference (mV)
R = gas constant (8,31439 J x mol-1 x K-1)
F = Faraday constant (96495,7 C x mol-1)
T = absolute temperature in Kelvin (K)
n = charge number of the measured ion (in this case nH = 1)
C1 = active H-ion concentration in solution C1
C2 = active H-ion concentration in solution C2
Nernst equation

32.  It consists of a platinum
wire in a inverted glass
tube. Hydrogen gas is
passed through the tube
at 1 atm.
 A platinum foil is attached
at the end of the wire. The
electrode is immersed in
1M H+ ion solution at
25°C.
 The electrode potential of
SHE is zero at all
temperatures.
Standard hydrogen electrode(SHE)

33.  It is rather difficult to regulate the pressure
of the H2 gas to be at exactly 1atm
throughout the experiment.
 If the solution contains any oxidizing agent,
the H2 electrode cannot be used.
 Excess of H2 bubbling out carries little HCl
with it and hence the H+ concentration
decreases. In such a system, it is difficult to
maintain the concentration of HCl at 1M.
Limitations

34.  It consists of a glass tube containing saturated
KCl connected to a platinum wires through
mercury-mercurous chloride paste.
Calomel electrode

36.  The silver/silver chloride reference
electrode is a widely used reference
electrode because it is simple,
inexpensive, very stable and non-toxic.
 it is mainly used with saturated potassium
chloride (KCl) electrolyte, but can be used
 with lower concentrations such as 1 M KCl
 and even directly in seawater.
Silver- Silver chloride electrode

38. Combination electrode

39.  An acidic solution has far more positively
charged hydrogen ions than an alkaline
one, so it has greater potential to produce
an electric current in a certain situation.
 In other words, it is a bit like a battery that
can produce a greater voltage.
 A pH meter takes advantage of this and
works like a voltmeter: it measures the
voltage (electrical potential) produced by
the solution
Working Mechanism

40.  When two electrodes (or one probe containing
the two electrodes) are dipped into solution,
 some of the hydrogen ions in the solution
move toward the glass electrode and replace
some of the metal ions in its special glass
coating.
 This creates a tiny voltage across the glass
the silver electrode picks up and passes to the
voltmeter.
 Reference electrode acts as a baseline or
reference for the measurement.

41.  A voltmeter measures the voltage generated
by the solution and displays it as a pH-
measurement.
 An increase in voltage means more hydrogen
ions and an increase in acidity, so the meter
shows it as a decrease in pH; in the same
way, a decrease in voltage means fewer
hydrogen ions, more hydroxide ions, a
decrease in acidity, an increase in
alkalinity, and an increase in pH.
 ↑ voltage = more H+/less OH- = ↑ acidity =
↓pH
 ↓ voltage = less H+/more OH- = ↓ acidity =
↑pH

42.  Calibration should be performed with at least
two standard buffer solutions that span the
range of pH values to be measured.
 For general purposes buffers at pH 4 and pH
10 are acceptable.
 The pH meter has one control (calibrate) to set
the meter reading equal to the value of the first
standard buffer and a second control (slope)
which is used to adjust the meter reading to
the value of the second buffer.
 A third control allows the temperature to be
Calibration

44.  Standards
- pH measurements cannot be more accurate than standards
(±0.01).
 Acid Error
- At high [H+], the measured pH is higher than actual pH, glass
is saturated.
 Hydration of glass
- A dry electrode will not respond to H+ correctly.
 Temperature
- Calibration needs to be done at same temperature of
measurement
 Cleaning
- Contaminates on probe will cause reading to drift until
properly cleaned or equilibrated with analyte solution
Errors in pH Measurements

45. Applications of pH meter
1.
• For the diagnosis of various disordersof humanin body
2.
• Agriculture
3.
• Brewing
4.
• Corrosion Prevention
5.
• Deying
6.
• Jam and jelly manufacturing