Study about pH meter, How it Works?,pH meter principle, its application, types of indicators used in pH determination, types of Buffers etc,..

1. pH – Principle, Measurement, Buffers and Buffer
index
(Subject: Applied Environmental Chemistry and
Microbiology)
Presented by: Under the Guidance of:
MELROY CASTALINO Dr. Veena Kumar Adi
Mtech (Environmental engineering)
B.I.E.T Davanagere.

2. Contents
 INTRODUCTION
 DEFINITION OF pH
 MEASUREMENT OF pH
 GLASS ELECTRODE
 WORKING OF GLASS ELECTRODE
 APPLICATION OF PH DATA IN ENVIRONMENTAL
ENGINEERING
 BUFFERS
 BUFFER INDEX
 APPLICATION OF BUFFERS
 REFERENCE
1

3. INTRODUCTION
pH is universally express the intensity of the acid or alkaline
condition.
Expressing the hydrogen-ion concentration or the hydrogen-ion
activity.
In olden days acids and bases were originally distinguished by
their taste, Acids litmus paper red. Bases taste turn litmus paper
blue.
Concept of pH, initially proposed by the Danish chemist
Sørensen in 1909.
2

4. DEFINITION OF pH
 pH is the negative logarithm of hydrogen ion concentration of an aqueous
solution.
 It is an important parameter to determine the quality of water.
 The pH value is expressed as:
pH = −log10[H+] = log(
1
H+)
measured in units of mol/lt of hydrogen ions.
 pH scale is usually represented as ranging from 0 to 14.
 Solutions with pH less than 7 are acidic and solution with pH greater than
7 are basic.
 pH is temperature dependent.
3

5. pH Scale:
4

6. pH Contd.,
The pH of natural water usually lies in the range of 6.0 to 8.5.
Scientists often use logarithmic scales for quantities that vary
widely in magnitude. So, pH is measured in log scale.
Ex: pH 4 , [H+]= 0.0001, [H+]=10−4
.
pOH is a measure of hydroxide ion OH−
ion concentration.
pOH= 14- pH
pH has no direct adverse effect on health.
Higher value of pH results in scale formation in heating
apparatus.
5

7. MEASUREMENT OF pH
The pH of any aqueous solution can be measured in different ways
such as.
 Indicators
 Glass electrode
 Hydrogen electrode
6

8. Indicators
Indicators are substances whose solutions change colour due to
changes in pH.
Colour changes in well-defined pH ranges.
In general, these compounds are weak acids or weak bases.
Figure shows pH scale with indicator.
7

9. Properties of
various acid-
base
indicators.
Indicators Acid colour pH range Base colour
Methyl violet Yellow 0-2 Violet
Malachite green (acidic) Yellow 0-1.8 Blue-green
Thymol blue (acidic) Red 1.2-2.8 Yellow
Bromphenol blue Yellow 3-4.6 Blue
Methyl orange Red 3.1-4.6 yellow-orange
Bromcresol green Yellow 3.8-5.4 Blue
Methyl red Red 4.4-6.2 Yellow
Litmus Red 4.5-8.3 Blue
Bromthymol blue Yellow 6-7.6 Blue
Phenol red Yellow 6.8-8.4 Red
Metacresol purple Yellow 7.6-9.2 Purple
Thymol blue (alkaline) Yellow 8-9.6 Blue
Phenolphthalein Colourless 8.2-9.8 pink
Thymolphthalein Colourless 9.3-10.5 Blue
Alizarin yellow Yellow 10.1-11.1 Lilac
Malachite green Green 11.4-13 colourless 8

10. Indicator Contd.,
Phenolphthalein changes from colourless to pink in the pH range 8.2
to 8.4 and has been the colour indicator commonly used in
environmental analysis.
pH levels of 3.7 to 4.5. the indicator commonly used in the past for
such purpose was methyl orange.
The colour of hydrangea a natural type of indicator.
9

11. Picture shows various colours of hydrangea flowers.
10

12. GLASS ELECTRODE
Would develop a potential related to the hydrogen-ion activity with-
out interference from most other ions.
 Its use has become the standard method of measuring pH.
Glass electrode is used most routine laboratory purpose, being
capable of measuring pH within a ±0.1 pH unit.
Glass electrode which generates a potential varying linearly with the
pH of the solution in which it is immersed.
11

13. Construction of Glass Electrode
Typical three-component composition of 72% SiO2, 22%
Na2O, 6% CaO.
Contains an internal standard acidity solution typically 0.1
M HCl, sometimes.
with the addition of AgCl wire.
12

14. Reference electrode
 Purpose of a reference electrode is to provide continuity to the
electrical circuit.
 since the potential across a single half-cell cannot be measured.
 A reference electrode should satisfy the following basic
requirements:
1) The potential developed should be independent of temperature.
2) The potential developed should be independent of H+ ion
concentration.
3) The potential developed should not change with time.
13

15. Glass electrode connected with reference electrode
14

16. Combination Electrode
Sometimes the reference and measuring electrodes are
housed together, This type of electrode is known as
Combination Electrode.
The combination is dipped in the solution whose pH is to
be measured.
The output voltage is the difference between the e.m.f.s
generated by the measuring glass electrode and the
reference electrode.
15

17. Figure for Combination Electrode
16

18. WORKING OF GLASS ELECTRODE
 A potential is developed across the two surfaces of this glass bulb, when dipped in
aqueous solution.
 This potential is sensitive to the H+ ion concentration, having a sensitivity of 59.2
mv/pH at 250C.
 The buffer solution inside the glass bulb has a constant H+ ion concentration and
provides electrical connection to the lead wire.
 When you place the bulb into a solution, the hydrogen ions in the solution will
move toward the glass electrode.
 This will cause a tiny voltage across the glass blub, voltage is picked up by the
silver wire and passed to the voltmeter
 The voltmeter will amplify and change the voltage value into a value we can use,
pH units
 The H+ ion activity difference between the solution inside the electrode and the
external one is measured as electric potential difference. Then, converted in pH
units (0-14) by the instrument.
17

19. Nernst equation
 The relation between the potential of a cell to the concentrations of participating ions is
Nernst equation named after Walter Nernst.
 It is a Nernst concentration cell with potential controlled by the activities of H+ on either
side of a very thin glass membrane. Latter is the bottom part of a bulb at the end of a glass
tube containing a reference solution of fixed value.
Em = 2.303
RT
F
log(
H+
sample
H+
standard
)
 Taking into account that [H+]sample is a constant and known quantity and that for the
sample solution it holds pHs = -log[H+]standard, eqn becomes
E = Econstant − Em
E = Econstant − 2.303
RT
F
pHs
 Where, R=8.314 (universal gas constant)
T= 298 K or 25 °C (absolute temperature)
F= 6×1023
× 1.6× 10−19
(Faraday constant)
18

20.  By substituting above values we get,
E = Econstant − 0.058 pH
Application of Nernst Equation:
 Determination of ion concentration.
 It is also used to calculate the potential of an ion charge across a
membrane.
 It is used in solubility product and potentio-metric titration.
 It is also used in pH measurements.
19

21. Calibration of pH meter
 For pH meters to be accurate, they have to be properly calibrated.
 So they usually need testing and adjusting before you start to use
them.
 You calibrate a pH meter by dipping it into buffers and adjust the
meter accordingly.
 Standard buffer solution of pH 4 and pH 9.2, dip the pH meter.
 Another important consideration is that pH measurements made this
way depend on temperature.
 Some meters have built-in thermometers
 Automatically correct their own pH measurements as the
temperature changes.
20

22. Hydrogen electrode
 A hydrogen electrode is made by platinum wire or a platinum plate.
 solution is saturated with hydrogen gas.
 The electrode potential is inversely proportional to pH of the
solution.
 The hydrogen electrode was found to be suitable device for
measuring hydrogen-ion activity.
 With its use, it was found that pure water dissociates to yield a
concentration of hydrogen ions equal to about 10−7
mol/L.
H2O ⇌ H+ + OH−
21

23. Hydrogen electrode Contd.,
 Since water dissociates to produce one hydroxide ion for each hydrogen
ion, it is obvious that about 10−7 mol/L of hydroxide ion is produced
simultaneously. By substitution into the equilibrium equation, we obtain
H+ OH−
H2O
= K
 But, since the concentration of water is so extremely large and is
diminished so very little by the slight degree of ionization, it may be
considered as constant (its activity equal 1) and Eq. can be written as
H+ OH− = Kw
 And for pure water at about 25°C,
H+ OH− = 10−7 × 10−7 = 10−14
This is known as the ion product or ionization constant for water
22

24. Hydrogen electrode Contd.,
 However, this method is not appropriate for daily use because of the
effort and expense involved
 With the inconvenience of handling.
23

25. Application of pH data in Environmental Engineering
It is important in almost every phase of environmental
engineering and science.
In the field of water supplies, it is a factor that must be
considered in chemical coagulation, disinfection, water
softening and corrosion control.
Determination of pH is one of the important objective in
biological treatment of waste water.
pH value or range is of immense value for any chemical
reaction.
It is used in the calculation of carbonate, bicarbonate, CO2
corrosion, stability index and acid base equilibria.
24

26. BUFFERS
 Buffers may be defined as substances in solution that offer resistance to
change in pH.
 Mixtures of weak acids and their conjugate bases or weak bases and their
conjugate acids.
 The fundamental basis for an understanding of buffer action is through
titration curve.
25

27.  The pH at the midpoint in the titration is given by pKA for weak acids and
by
 pKw − pKB for weak base.
 Fig. titration curve for weak acids
26

28.  They may also be used effectively within ±1 pH unit of the pK value
by changing the relative concentration of acid to salt or base to salt.
 Consider the equilibrium relationship for weak acid.
KA =
H− A−
HA
 By rearranging, we obtain
1
H+
=
1
KA
A−
HA
 Thus 𝑝𝐻 = 𝑝𝐾𝐴 + log
𝑠𝑎𝑙𝑡
𝑎𝑐𝑖𝑑
pKA = -log1.8 x 10-5 (at 25°C)
27

29. Example of How a Buffer Works
Consider a solution containing both acetic acid, CH3COOH (weak acid) and
CH3COONa (strong base).
CH3COOH ionizes partially as
CH3COOH ⇌ CH3C00−
+ H+
When strong base is ionizes
CH3COONa ⟶ CH3COO−
+ Na+
If a little strong acid like HCl is added 𝐻+
ions will combine with CH3COO−
ions of the solution
to form a weak acid CH3COOH molecule
H+
+ CH3COO−
⇌ CH3COOH
Hence the H+ ion concentration of the solution doesn’t increase i.e, pH doesn’t decrease.The
amount of strong acid or base that a buffer can neutralize is called the buffer capacity. 28

30. Numerical example for calculating the pH of a buffer solution.
Q. Calculate the pH of a buffer solution containing 0.01 M acetic acid and 0.01 M sodium acetate. Then calculate the
pH after enough HCl is added to given concentration of 0.001 M.
Sol.
We know the Eq.
pH = pKA + log
salt
acid
pH = 4.74 + log(
0.01
0.01
)
𝑝𝐻 = 4.74 + 0
pH = 4.74
when the HCl is added, and we assume that all the HCl dissociate then above Eq. becomes
pH = 4.74 + log(
0.01 − 0.001
0.01 + 0.001
)
pH = 4.74 – 0.09
pH = 4.65 29

31. BUFFER INDEX
 The buffering capacity of a solution can be indicated quantitatively by the
buffering index (sometimes called buffer intensity) β,
 which is defined as the slope of a titration curve of pH versus moles of
strong base added (𝐶 𝐵) or moles of strong acid added (𝐶𝐴).
β =
dCB
dpH
= −
dCA
dpH
Buffer index , β = 2.303[
Kw
H+ + H+ +
CTKA H+
KA+ H+ 2
Where,
𝐾𝐴 = 1.8 × 10−5
𝐾 𝑤 = 10−14
30

32. Numerical example for Buffer index
Q. A buffer solution has been prepared by adding 0.2 mol/L of acetic acid and 0.1 mol/L of
acetate. The pH of the solution has been adjusted to 5.0 by addition of NaOH. How much
NaOH (mol/L) is required to increase the pH to 5.1?
Sol.
Total concentration CT= 0.2+0.1=0.3 mol/L
[H+] =10−5
We know the Eq.
β = 2.303[
Kw
H+
+ H+ +
CTKA H+
KA + H+ 2
β = 2.303[
10−14
10−5
+ 10−5 +
0.3 × 1.8 × 10−5 × 10−5
1.8 × 10−5 + 10−5 2
β = 0.16 mol/L of base per pH unit
NaOH required= (5.1-5) × 0.16
NaOH required= 0.016 mol/L
31

33. Application of buffers
 Buffer solutions are very useful for chemical manufacturing.
 Essential for many biochemical processes.
 Buffer solutions are necessary to keep the correct pH for enzymes in many
organisms to work.
 Industrially, buffer solutions are used in fermentation processes
 Setting the appropriate conditions for dyeing fabrics.
 They are also used in chemical analyses and syntheses, and for the
calibration of pH meters.
32

34. Reference
 Clair N. Sawyer and Perry L. McCarty “Chemistry for Environmental
Engineering and Science” Fifth Edition.
 A. E. Mirsky and M. L. Anson. “ A description of the glass electrode and its use
in measuring hydrogen ion concentration.”
 Bates, Roger G. “Determination of pH: theory and practice”.
 Butler J. N. “Ionic Equilibrium: A Mathematical Approach”.
 Scorpio, R. “Fundamentals of Acids, Bases, Buffers & Their Application to
Biochemical Systems”
33

35. IMPORTANT QUESTIONS FOR EXAMS
 What is Glass electrode? With the principle involved, explain the
measurement of pH using glass electrode? (May/June 2010)
 What is Buffer? Explain buffer action with a suitable example? (May/June
2010)
 Explain the mechanism of pH stability of a buffer solution with suitable
example? (Dec 2010)
 Explain pH and p(x) concept? (Dec 2010)
 Using approximate equations, calculate the pH of i) 0.01M(600 mg/l)
acetic acid solution ii) 0.01M phosphoric acid solution. Given pKA of
acetic acid= 4.74 ; pKA of phosphoric acid= 2.12. (Dec 2010)
 What is meant by pH of a solution? How is the pH of a solution
determined experimentally, using glass electrode? (Dec 2011)
34

36. Contd..,
 Explain with neat diagram, the functioning of a pH meter? (Dec 2015/Jan
2016)
 What are Buffer? How do they function? (Dec 2015/ Jan 2016)
 Calculate the pH of a buffer solution containing 0.4M calcium acteate
and 0.2M acetic acid. Given that Ka= 1.8*10^-5. (Dec 2015/Jan 2016)
 With a labeled diagram, explain the functioning of pH electrode? (Dec
2016/Jan 2017)
35

37. THANK YOU