Spectroscopic pH Measurement Using Phenol Red Dye

1
Spectroscopic pH Measurement
Using Phenol Red Dye
By
Tewodros Adaro
A Thesis Presented to the School of Graduate Studies Addis
Ababa University in Partial Fulfilment of the Requirements
for the Degree Master of Science in Physics.
Addis Ababa, Ethiopia
June 2010
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Addis Ababa University
School of Graduate Studies
Spectroscopic pH Measurement Using Phenol
Red Dye
By
Tewodros Adaro
Department of Physics
Addis Ababa
Approved by the examining Board:
Prof. A.V. Gohlap Advisor____________
Dr.Gizaw Mengistu Examiner__________
Prof. A.K. Chaubey Examiner__________
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ACKNOWLEDGEMENTS
My special thanks go to my advisor prof.A.V Gohlap. Besides bringing the
research area to my attention, his direction, keen insight, guidance, and skilful
pushes to get me explore had a huge impact both in this research and on my
academic development. I also would like to express my heartfelt gratitude to
Dr.Gizaw Mingistu for his all rounded support.
My special thanks are also due to my friends Yirga Taming, Wasihun Temam
and Zemene for their all rounded contribution and support in the course of
conducting this research work.
To all others who helped me in completing this study, I am equally grateful.
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Table of contents
Acknowledgements
Table of contents V
List of tables Vii
List of figures Viii
List of figures iX
Symbols used X
Abstract Xi
Introduction 1
CHAPTER ONE 4
1. Absorption 4
1.1. Theory of Absorption 4
1.2. Absorption of Electromagnetic Radiation 5
1.3. Absorption spectroscopy 6
1.4. Lambert-Bouguer law 9
1.4.1 Transmittance and Absorbance 10
1.5. Beer- Lambert law 11
1.5.1 Limitation to Beer’s law 13
1.6 Practical Applications of Absorption spectroscopy 14
CHAPTER TWO 16
2. pH measurement by absorption spectroscopy 16
2.1 Activity and the Definition of pH 16
2.1.1 Hydrogen ion activity 16
2.2 Buffer solution 16
2.2.1 Buffering agent 17
2.3 pH indicators 18
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2.3.1 Phenol red (phenolsulphonephthalein, ) 19
2.4 Determination of pH of indicator dye 20
2.5 Applications of pH measurement 25
CHAPTER THREE 31
3. Experimental 31
3.1 Basic of UV/VIS Spectrophotometer 31
3.1.1 Instrumental component 31
3.1.2 General arrangement experimental
set up of spectrometer 31
3.2 Reagents 32
3.3 Preparation of Buffers 32
3.3.1 Preparation of Acetate Buffer Solution (pH 3-6) 32
3.3.2 Preparation of phosphate buffer solutions (pH 7- 11) 32
3.4 Preparation of phenol red stock solution 33
3.5 preparation caffeine sample solution 33
3.6 Measurements 33
3.6.1 Solutions prepared for Absorption spectra
measurement of phenol red as a proton donor
and proton acceptor form. 34
3.6.2 Solutions prepared for Absorption spectrum
measurement of phenol red in buffer solution 34
3.6.3 Sample prepared for Absorption spectrum
measurement of phenol red in caffeine 36
3.6.4 Sample prepared for Absorption spectrum
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measurement of phenol red in packed
drinking spring-water. 36
CHAPTER FOUR 37
4. Result and Discussion 37
4.1 Result 37
4.1.1 Color response of phenol red 37
4.1.2 Absorption spectrum of pure acidic
and basic form of phenol red 37
4.1.3 Absorption spectrum of phenol red in
different pH value buffer solutions 38
4.1.4 Absorption spectrum of phenol red in
different pH value spring waters 39
4.2 Discussion 41
4.2.1 Response spectra of pure acid (A) and
pure base (B) form of phenol red 41
4.2.2 Response spectra of phenol red in buffer solutions 42
4.2.3 Phenol red determination 44
4.2.4 Spectroscopic pH calculation of packed
drinking spring-water 46
4.2.5 Comparison between stated values of pH and
spectroscopic measured pH values of spring water 48
CHAPTER FIVE 50
5. CONCLUSION AND RECOMMENDATION 50
5.1 Conclusion 50
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List of tables
Table-2.1 color of different indicators dye 19
Table-3.1 volume of 0.1M acetic acid and 0.1M sodium acetate
mixed to get Acetate buffer 32
Table-3.2 volume of 0.1M disodium hydrogen orthophosphate,
0.1M hydrochloric acid and 0.1M sodium hydroxide
to get Phosphate buffer 33
Table-3.3
Table-3.4
Table 4.1 absorbance of phenol red at = 432.0nm and
= 558.4nm in pH=3.0 and pH=12 buffer solutions 38
= 558.4nm in pH 5,6,7,8 and11 buffer solutions 39
= 558.4nm in five types of packed drinking spring-water. 40
Table 4.4 calculated value of log
[ ]
[ ]
for various pH value of
buffer solution. 46
Table 4.5 calculated pH values for packed drinking spring-water 48
Table 4.6 calculated pH and standard deviation (sd) values
for packed drinking spring-water 48
Table 4.7 comparison in pH of spectroscopically calculated and
the stated values for spring-water. 49
List of figures
Fig-1.1 Transitions to excited state 8
Fig-1.2 Light absorption and transmission by phosphate-molybdenum
blue compound. Schematic diagram showing maximum light
absorption (and minimum light transmition) at =640nm. 11
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Fig-1.3 Interaction of light and matter 12
Fig-2.1 (a) Color response of phenol red at different pH value
(b) Powder form of phenol red
(c) chemical structure of phenol red. 20
Fig-3.1. Basic Experimental set up of spectrometer 31
Fig-4.1 Absorption spectrum of pure acidic and
pure basic form of phenol red 37
Fig-4.2 Absorption spectrum of phenol red
in different pH value buffer solution 38
Fig-4.3 Absorption spectrum of phenol red
in different packed spring water 39
Fig-4.4 Spectra of (a) the undissociated and (b) the dissociated
form of phenol red 42
Fig-4.5 Absorption spectra of phenol red doped sol-gel 43
Fig-4.6 Absorbance vs pH (a) at a wave length of 432nm and
(b) at a wave length of 558.4nm 43
Fig-4.7 A plot of pH vs log
[ ]
[ ]
. 45
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Abstract
The optical spectroscopic technique using phenol red dye was used to
measure pH. Buffer solutions having known pH values were used to
characterize phenol red equilibrium dissociation constant. Optical (Absorption)
spectra of phenol red in these buffer solutions were used to calibrate the dye
equilibrium dissociation constant. of phenol red was found to be 7.89511.
Once the dye equilibrium dissociation constant has been characterized, pH of
unknown solutions can be obtained from their dye spectra. Using this method
pH of caffeine and five different types of packed drinking spring water is
calculated. The error of this method is found to be 0.005 pH unit.
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Introduction
Accurate measurement of pH is important in several diverse field, such as
process control, reaction equilibrium and kinetics, environmental research to
monitor sea water chemistry and natural water quality, biomedical research
and in oilfields [1]. Measurement of pH is one of the most important and
frequently used tests in water chemistry . Practically every phase of water
supply and waste water treatment need pH determination e.g. acid-base
neutralization, water softening, precipitation, coagulation, disinfection, and
corrosion control, is pH dependent. pH is used in alkalinity and carbon dioxide
measurements and many other acid base equilibra.
The International Union of Pure and Applied Chemistry (IUPAC) have issued
guidelines for the standard potentiometric technique for pH measurements.
These methods and calibrating standards are, however, recommended only for
278-3230k, 0.101325Mpa and ionic strengths below 0.1mol/kg water [2]. The
reasons for these measurement constraints are the uncertainty in liquid
junction potential and reference electrode stability at high temperature,
pressure and ionic strength and the lack of calibrating standards.
Spectroscopic measurement of pH with very high accuracy using pH-
sensitive dye is a well established laboratory technique for ambient conditions
since the early 1900s [3, 4]. More recently, this technique has been shown to
improve precision for seawater pH measurements [19]. Using equimolal tries
buffers for total pH scales; dye equilibrium dissociation constant was
characterized at 0.010325Mpa pressure as a function of temperature (2930K to
3030K) and salinity over the narrow salinity range characteristics of sea water
(30 to37 salinity ranges corresponding to ionic strength of~0.53 to0.66 mol/kg
water). Yao and Byrne [5] have also applied this technique for fresh water pH
measurement, where potentiometric methods can prove to be problematic.
Using phosphate buffers, the dye equilibrium dissociation constant was
characterized at 0.101325Mpa pressure as a function of temperature (2830k to
3030k) and ionic strength (0 to .016 mol/ kg water). The Davies equation,
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which is valid at low ionic strengths, was used to calculate the activity
coefficients. Marty et al. [6] describe an autonomous spectroscopic pH sensor
for in situ measurements of natural waters that have low ionic strength (~0.001
mol/kg water) and temperature range of 277-2930k. E.Wang et al. [7] describe
a fast and long term optical sensors for pH based on sol-gels that the sensor
has a response time of less than 20s, the response completely reversible and its
life time is over 12 months.
These references cite the advantage of the spectroscopic technique with
respect to low-drift, reproducibility, and rapidness of the measurement as
compared to the standard glass electrodes. Furthermore, since the
measurement depends only on the molecular properties of the indicator dye,
once the dye equilibrium dissociation constants have characterized, it eliminate
the need for calibration prior to every measurement.
The spectroscopic method depends up on the direct determination of the
ratio of molecular species (neutral molecule) to ionized species. In this respect,
the data are no more different from those obtained by potentiometric titration;
the thermodynamics of ionic and anionic interaction do not depend on the
experimental method used to determine them [8]. For this purpose, the
spectrum of non-ionized species is obtained, using a buffer solution whose pH
is so chosen that the compound to be measured is present wholly as this
species. This spectrum is compared with that of pure ionized species similarly
isolated at another suitable pH. By using this at various pH values,
intermediate between those at which the spectra of the two species were
obtained. This is possible because a series of two component mixture is formed
in which the ratio of the two species depends solely up on pH at which the
solution is optically measured [8].
Dissociation constant of substances can be determined by several
different methods. The potentiometric, chromatographic, electrophoretic
methods also have been used widely. But a method based on spectrometry has
been still used widely by the help of improving computer programs [8]. In most
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of these methods a physical property to the analyte is measured as a function
of the pH of the solution and resulting data are used for the determination of
dissociation constant.
This work shows how the spectroscopic pH measurement technique can
be applied over the IUPAC recommended guideline for the standard
potentiometric techniques. Dye equilibrium dissociation constants have been
calibrated for room temperature, atmospheric pressure and lower ionic
strength because, pH calibrating buffer are available at 0.010325Mpa pressure
in the 273-3230k temperature range and for the 0.05-0.1 mol/kg water ionic
strength range. This work then shows how to use these pH values to calibrate
the dye technique. Once these dyes are characterized, they can be used to
measure pH of unknown solutions without any calibration requirement prior to
each measurement. These advantages make this a potentially attractive pH
measurement technique.
This thesis is structured as follows:
Chapter one deals with details of absorption. In this, theory of
absorption, absorption of electromagnetic wave, absorption spectroscopy
and beer-lambert law are included.
Chapter two deals with definition of pH, pH indicator dyes and
mathematical calculation of determination of pH of indicator dye.
Chapter three is devoted to experimental parts.
Chapter four is totally about result and discussion.
Chapter five is about conclusion and recommendation.
The objectives of this thesis are:
To determine pH spectroscopically using phenol red dye.
To compare stated values of pH of packed drinking spring-water to
spectroscopically measured values.
To compute the pH of caffeine.
To show the technique can be applicable for different type of solutions.
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CHAPTER ONE
1. Absorption
1.1. Theory of Absorption
Most objects around us are not self-luminous but are nevertheless visible
because they scatter the light that falls up on them. Most objects are colored,
however because they absorb light, not simply they scatter it. The colors of an
object typically arise because materials selectively absorb light of certain
frequencies, while freely scattered or transmitting light of other frequencies.
Thus if an object absorbs light of all visible frequencies, it is black. An object is
red if absorbs all (visible) frequencies except those our eye perceive to be “red”
wave length roughly between about 6300 and 6800 Ao), and so on.
The physics of the absorption process is simplest in well-isolated atoms.
These are found most commonly in gases. White light propagating through a
gas is absorbed at the resonance frequencies of the atoms or molecules, so that
one observes gaps in the wave length distribution of the emerging light.
The absorbed energy is particularly converted in to heat (translational
kinetic energy of the atoms) when excited atoms (or molecules) which have
absorbed radiation collide with other particles. The absorbed radiation is also
partially reradiated in all directions at the frequency of the absorbed radiation.
This is called resonance radiation, or resonance fluorescence.
Absorption in liquids and solids is much more complicated than in gases.
In liquids and amorphous solids such as glass, the absorption lines have such
large width that they overlap, water for example, is obviously transparent in
visible, but absorbs in the near infrared, i.e. at infrared wave lengths not far
removed from the visible. Its absorption curve is wide enough, in fact that it
extends in to the red edge of the visible. The weak absorption in the red portion
of the visible spectrum explains why things appear green when one is
sufficiently submerged under water.
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A broad absorption curve covering all visible wavelengths except those in
particular narrow bands is characteristics of the molecules of dye. The
absorption radiation is converted in to heat before it can be reradiated. Such
broad absorption curves and fast quenching rate require the high molecular
number densities of liquids and solids.
1.2. Absorption of Electromagnetic Radiation
The process whereby the intensity of a beam of electromagnetic radiation
is attenuated in passing through a material medium by conversion of the
energy of the radiation to an equivalent amount of energy which appears within
the medium; the radiant energy is converted into heat or some other form of
molecular energy. A perfectly transparent medium permits the passage of a
beam of radiation without any change in intensity other than that caused by
the spread or convergence of the beam, and the total radiant energy emergent
from such a medium equals that which entered it, whereas the emergent
energy from an absorbing medium is less than that which enters, and, in the
case of highly opaque media, is reduced practically to zero.
No known medium is opaque to all wavelengths of the electromagnetic
spectrum, which extends from 1010 m infrasound radio waves, whose
wavelengths are measured in kilometers, through the infrared, visible, and
ultraviolet spectral regions, x-rays and gamma rays, of wavelengths down to
10 14 m(1022Hz) cosmic radiation. Similarly, no material medium is transparent
to the whole electromagnetic spectrum. A medium which absorbs a relatively
wide range of wavelengths is said to exhibit general absorption, while a
medium which absorbs only restricted wavelength regions of no great range
exhibits selective absorption for those particular spectral regions. For example,
the substance pitch shows general absorption for the visible region of the
spectrum, but is relatively transparent to infrared radiation of long wavelength.
Ordinary window glass is transparent to visible light, but shows general
absorption for ultraviolet radiation of wavelengths below about 310
nanometers, while colored glasses show selective absorption for specific regions
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of the visible spectrum. The color of objects which are not self-luminous and
which are seen by light reflected or transmitted by the object is usually the
result of selective absorption of portions of the visible spectrum. Many colorless
substances, such as benzene and similar hydrocarbons, selectively absorb
within the ultraviolet region of the spectrum, as well as in the infrared.
1.3 Absorption spectroscopy
Absorbance spectroscopy, commonly referred to as spectrophotometery,
is the analytical technique based on measuring the amount of light absorbed
by a sample at a given wave length [9]. spectrophotometery ,particularly in the
VIS and UV portions of the electromagnetic spectrum ,is one of the most
versatile and widely used techniques in chemistry and the life sciences.
Molecular absorption spectroscopy in the ultraviolet (UV) and visible (VIS)
is concerned with the measured absorption of radiation in its passage through
a gas, a liquid or a solid, the wave length region generally used is from 190 nm
to about 1000nm ,and the absorbing medium is at room temperature.
A molecule or part of a molecule that can be excited by absorption is called
chromophores. Organic chromophores which absorbs strongly in the UV or
visible portions or the spectrum.
Molecular excitation energy is usually dissipated as heat (kinetic energy)
by the collision of the excited molecules with another molecules (e.g. a solvent
molecule), as the molecule returns to the ground state. In other cases, the
excitation energy is dissipated by the emission of light in a process called
“fluorescence”. In both cases, the intensity of the light transmitted by a
collection of chromospheres is less than the intensity of the incident light.
An excited molecule can possess only one of a set of discrete amounts
(quanta) of energy described by the laws of quantum mechanics. These
amounts are called the “energy levels” of the molecule. In UV/VIS
spectrophotometery, the major energy levels are determined primarily by the
possible spatial distributions of the electrons and are called electronic energy
levels, and to a lesser extent by viberational energy levels, which arise from the
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various modes of vibrations of the molecules (e.g. the stretching and bending of
various covalent bonds)
The energy and wave length of absorption is defined by the difference
between energy levels of and electronic transition. This can be expressed by
=( )
(1.3.1)
E1 is the energy level of the molecule before absorption
E2 is an energy level reached by absorption
If all transitions were between only the lowest vibrational levels of the
ground state and the first excited state, then an absorption spectrum would
consist of narrow, discrete peak. However, the transition from one electronic
level to the next level can occurs between many vibrational levels also
Since energy differences between vibrational levels within an electronic
level are small compared to the minimum energy difference between electronic
levels, the electronic transition consists of a cluster of very closely spaced
spectral peaks. Each peak has significant width, comparable to the spacing
between the peaks. This has the effect that the peaks overlap so much that a
single broad peak, called an electronic absorption band, results.
For most molecules, absorption wave length corresponding to transitions
between the ground state and any vibrational level of the first excited state, fall
in the range of ultraviolet and visible light.
Low –energy transitions are also possible between vibrational levels within
a single electronic level. This transitions produce radiation in the infrared
range. Fig 1.1 illustrates the transitions between the ground state and any
vibrational level of the first excited state.
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Fig-1.1 Transitions to excited state
Molecules which absorb photons of energy corresponding to wave lengths
in the range 190 nm to about nm, exhibit UV/VIS absorption spectra the
quantized internal energy ( )of as molecule in its electronic ground or excided
state can be approximated , with sufficient accuracy for analytical purposes, by
the following equation
(1.3.2)
is the electronic energy
is the vibrational energy
the rotational energy
Absorption of a photon results in a change of electronic energy,
accompanied by changes in the vibrational and rotational energies. Each
vibronic transition, i.e. a particular electronic plus vibrational transition,
corresponds to an absorption band consisting of rotational lines. In liquids and
solids, the rotational lines are broad and overlap so that no rotational structure
is distinguishable [9].
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1.4 Lambert-Bouguer law
A relationship between the absorption of light in a purely absorbing
medium and the thickness of the medium was first determined in 1729 by
Bouguer (1729) some years later Lambert (1760) derived the mathematical
expression for the relationship, known as the lambert-Bonguer law. If a
homogeneous medium is thought of as being constituted of layers of uniform
thickness set normally to the beam, each layer absorbs the same fraction of
radiation incident on it. If I is the intensity to which a monochromatic parallel
beam is attenuated after traversing a thickness l of the medium, and I0 is the
intensity of the beam at the surface of incidence (corrected for loss by reflection
from this surface), the variation of intensity throughout the medium is expressed
by:
(1.4.1)
Which describes how each successive layer of the medium absorbs the same
fraction of the incident intensity for a constant , the latter known as the
absorption coefficient with units of inverse length (usually mm-1). For incident
intensity I0, therefore, the transmitted intensity I through a distance will be:
(1.4.2)
The absorption coefficient can thus be interpreted as probability that a
photon will be absorbed by the medium per unit length. The reciprocal of the
absorption coefficient, known as the absorption length, is the distance required
for the intensity of the beam to fall to e-1 of the initial intensity. When equation
(1.4.2) is expressed in base 10 logarithms as:
10 (1.4.3)
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19
Then the constant k = 2.303 is known as the extinction coefficient at radiation
of the wavelength considered.
Equation (1.4.3) shows that as monochromatic radiation penetrates the medium,
the logarithm of the intensity decreases in direct proportion to the thickness of
the layer traversed. If experimental values for the intensity of the light emerging
from layers of the medium of different thicknesses are available (corrected for
reflection losses at all reflecting surfaces), the value of the extinction coefficient
can be readily computed from the slope of the straight line representing the
logarithms of the emergent intensities as functions of the thickness of the layer.
1.4.1 Transmittance and Absorbance
Suppose the intensity of the beam of light entering the sample( 0) and
excited the sample( ). Taking the ratio get an indication of what fraction
of the light entering the sample was found exiting the sample. This ratio is
called the transmittance.
Transmittance: (1.4.4)
Absorbance (A): is defined as the log10 of the reciprocal of the
transmittance. This is a measure of how much light is absorbed by the sample,
rather than transmittance through it.
Absorbance: = log 1
log (1.4.5)
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1.5 Beer- Lambert law
This law refers to the effect of the concentration of the absorbing medium,
that is, the mass of absorbing material per unit of volume, on the absorption.
This relation is of prime importance in describing the absorption of solutions of
an absorbing solute, since the solute’s concentration may be varied over wide
limits, or the absorption of gases, the concentration of which depends on the
pressure.
According to Beer’s law, each individual molecule of the absorbing material
absorbs the same fraction of the radiation incident upon it, no matter whether
the molecules are closely packed in a concentrated solution or highly dispersed
in a dilute solution. The relation between the intensity of a parallel
monochromatic beam which emerges from a plane parallel layer of absorbing
solution of constant thickness and the concentration of the solution is an
exponential one, of the same form as the relation between intensity and
thickness expressed by Lambert’s law.
In 1852 Beer determined that the absorption coefficient of a compound is
linearly related to its concentration diluted in a non-absorbing medium (Beer,
1852)[10].
(1.5.1)
Where is known as the specific absorption coefficient substituting for µ
in the Lambert-Bouguer law gives what is known as Beer-Lambert law.
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(1.5.2)
Expressing the Beer – Lambert law in log10 gives
= log (1.5.3)
Where /2.303
Then absorbance related to extinction coefficient, concentration and path
length by an equation:
(1.5.4)
Where
is the specific extinction coefficient,
is concentration and
is path length.
Fig 1.3 interaction of light and matter
In a solution containing a mixture of n absorbing compounds, the total
absorbance is the sum of the individual extinction coefficients multiplied by the
distance .
= ( )
= ( ) (1.5.5)
Beer’s law is true for a particular solution, the plot of log (I0 /I) against the
concentrations for solutions of different concentrations, measured in cells of
constant thickness, will yield a straight line, the slope of which is equal to the
molar extinction coefficient. While no true exceptions to Lambert’s law are
known, exceptions to Beer’s law are not uncommon.
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Such exceptions arise whenever the molecular state of the absorbing solute
depends on the concentration. For example, in solutions of weak electrolytes,
whose ions and undissociated molecules absorb radiation differently, the
changing ratio between ions and undissociated molecules brought about by
changes in the total concentration prevents solutions of the electrolyte from
obeying Beer’s law. Aqueous solutions of dyes frequently deviate from the law
because of dimerization and more complicated aggregate formation as the
concentration of dye is increased.
1.5.1 Limitation to Beer’s law
The Beer-Lambert law can be split into two parts with Beer’s law stating
that there is a linear relationship between absorbance and concentration at a
fixed path length, and Lambert’s law stating that there is a linear relationship
between absorbance and path length at a fixed concentration.
As this investigation uses a fixed path length only Beer’s law will be discussed.
In practice there are frequently deviations from the direct proportionality of
Beer’s law but these deviations are well known spectrophotometer phenomena,
and can be categorized as; real limitations to Beer’s law, instrumental
deviations and chemical deviations.
A. Real limitations to Beer’s law
The real limitations to Beer’s law concern its application to concentrated
solutions when it should only be applied to dilute solutions: when solutions
contain high concentrations of ions (>0.01M) these charged species affect the
charge distribution on adjacent analyte molecules, which in turn can change
the high absorbing character of the molecule. In the case of a dilute absorber in
the presence of a non-absorbing electrolyte, the direct proportionality between
absorbance and concentration of the absorber can also be disturbed by the
electrolyte concentration.
B. Chemical deviations
Chemical deviations are the apparent deviations from Beer’s law that
occur when an analyte changes in the presence of the solvent to form
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compounds that have a different light absorbing character from the parent
species. The ionization reactions of acidic or basic indicators are examples of
this behavior. As the indicator concentration increases so does its influence on
the pH, and when the pH changes so do the proportions of the different ionized
species which in turn changes the absorptivity of the solution.
C. Instrumental deviations
There are two main types of instrumental deviations, those due to the
presence of polychromatic radiation, and those due to the presence of stray-
light. Polychromatic deviations occur when more than one wavelength of light
is present and the radiant power of these wavelengths is absorbed in different
proportions by the analyte. The difference between the molar absorptivities of
the compound at the different wavelength determines the size of this Beer’s law
deviation. However, if the analyte has a consistent absorptivity within the
polychromatic range, this instrumental error will not be appreciable.
Light scattering inside the detection area of the spectrophotometer
causes stray-light deviations from Beer’s law. The resulting absorbance
deviation increases with increasing absorbance because the stray-light
represents an increasingly significant part of the signal that reaches the
detector. Such instrumental deviations from Beer’s law always cause
underestimates of the analyte concentration.
1.6 Practical Applications of Absorption spectroscopy
Absorbance measurements allow the following
Determination of the concentration of a substance
Kinetic assay of certain chemical reactions
The identification of materials
The most common use of absorbance measurements is to determine the
concentration of a solute. This can be done if the absorption coefficient ( ) is
known and Beer’s law is obeyed.
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In practice, we do not generally rely on published values of because
this quantity may be very sensitive to idiosyncrasies of reagent preparation and
instrument design[9].
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CHAPTER TWO
1. pH measurement by absorption spectroscopy
2.1 Activity and the Definition of pH
2.1.1 Hydrogen ion activity
pH was originally defined by Sørensen in 1909 in terms of the
concentration of hydrogen ions (in modern nomenclature) as log
where is the hydrogen ion concentration in mol dm–3, and = 1 mol dm–3 is
the standard amount concentration. Subsequently, it has been accepted that it
is more satisfactory to define pH in terms of the relative activity of hydrogen
ions in solution.
log log (2.1.1)
Where is the relative (molality basis) activity and is the molal activity
coefficient of the hydrogen ion H+ at the molality , and is the standard
molality. The quantity pH is intended to be a measure of the activity of
hydrogen ions in solution. However, since it is defined in terms of a quantity
that cannot be measured by a thermodynamically valid method, eq. 2.1.1 can
be only a notional definition of pH[2].
Because of ionic interactions in all but very dilute solutions, it is
necessary to use the “activity” of an ion and not its molar concentration. Use of
the term pH assumes that the activity of the hydrogen ion, , is being
considered. The approximate equivalence to molality [ ] can be presumed only
in very dilute solutions (ionic strength< 0.1).
The pH value of a highly dilute solution is approximately the same as the
negative common logarithm of the hydrogen ion concentration.
2.2 Buffer solution
A buffer solution is an aqueous solution consisting of a mixture of a weak
acid and its conjugate base or a weak base and its conjugate acid. It has the
property that the PH of the solution changes very little when a small amount of
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acid or base is added to it. Buffer solutions are used as a means of keeping pH
at a nearly constant value in wide variety of chemical application [11].
2.2.1 Buffering agent
A buffering agent adjusts the PH of a solution. The function of a buffering
agent is to drive an acidic or basic solution to a certain pH state and prevent a
change in this pH, Buffering agents have variable properties. Some are more
soluble than others; some are acidic while others are basic. As pH managers,
they are important in many chemical applications, including agriculture, food
processing, medicine and photography.
A. What a buffering agent is
Buffering agents can be either the weak acid or weak base that would
comprise a buffer solution. Buffering agents are usually added to water to form
buffer solutions. They are the substances that are responsible for the buffering
seen in these solutions. These agents are added to substances that are to be
placed into acidic or basic conditions in order to stabilize the substance. For
example, buffered aspirin has buffering agent, such as , that will maintain
the pH of the aspirin as it passes through the stomach of the patient. Another
use of a buffering agent is in anti acid tablets, whose primary purpose is to
lower the acidity of the stomach.
B. How a buffering agent works
The way buffering agents work is seen in how buffer solutions work.
Using Henderson–Hasselbakh equation we get an equilibrium expression
between the acid and conjugate base. As a result we see that there is little
change in the concentrations of the acid and base so therefore the solution is
buffered. A buffering agent sets up this concentration ratio by providing the
corresponding conjugate acid or base to stabilize the pH of that which it is
added to.
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2.3 pH indicators
pH indicators are usually weak acids or weak bases that changes their
color depending on their dissociation (protonation) state. Sometimes both forms
are colored, sometimes only one. In most cases you may assume that too
completely change color of bicolored indicator pH must change by 2 units.
However, human eye is more sensitive to some colors than to others, thus
some color changes can be perceived over wider pH range.
A number of naturally occurring or synthetic organic compounds
undergo definite color changes in well-defined pH ranges. A number of
indicators that are useful for various pH ranges are listed in Table 2.1. The
indicators assume different hues within the specified pH range. These are
used as liquid solutions or some as pH papers. The change in color
occurs over a wide range of pH change and therefore pH value cannot be
measured accurately. Further, the turbidity and color of the sample may cause
interference.
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Table 2.1 color of different indicators dye
2.3.1 Phenol red (phenolsulphonephthalein, )
Phenol red exists as a red crystal (fig 2.1. (b)) that is soluble in water. Its
solubility is 0.77 g/L in water and 2.9 g/L in ethanol. It is a weak acid with
pka range from 7.9 to 8.0 at 200C. A solution of phenol red is used as a pH
indicator; it has an orange color when dissolved in water and its color exhibits
a gradual transition from yellow to red over the pH range 6.8 to 8.2. Above pH
8.2 phenol red turns a bright pink (fuchsia) color as shown in figure2.1.(a).This
color change is caused by phenol red losing protons (and changing color) as the
pH increases. In crystalline form, and solution under very acidic condition (low
pH), the compound exists as a Zwitterion as in the structure shown in figure
2.1 (c), with the sulfate group negatively charged, and the ketone group
carrying an additional proton. This form is sometimes symbolically written
as and is orange-red. If the pH is increased, the proton from the ketone
Indicator Acid
color
pH Basic
color
Methyl violet Yellow 0-1.6 Violet
Thymol blue Pink 1.2-2.8 Yellow
Methyl Yellow Red 2.9-4.0 Yellow
Brom-Phenol blue Yellow 3.0-4.7 Violet
Methyl orange Pink 3.1-4.4 Yellow
Brom-cresol green Yellow 3.8-5.4 Blue
Methyl red Red 4.2-6.2 yellow
Litmus Red 4.7-8.2 Blue
Chlorophenlo red Yellow 4.8-6.4 Red
Brom-cresol purple Yellow 5.2-6.8 Purple
Bromothymol blue Yellow 6.0-7.6 Blue
Phenol red Yellow 6.4-8.0 Red
Cresol purpl Yellow 7.4-9.0 Purple
Thymol blue Yellow 8.0-9.6 Blue
Phenolphthalein colorless 8.0-9.8 Pink
Thymolphthalien colorless 9.3-10.5 Blue
Alizarin yellow G colorless 10.1-12.1 Yellow
Indigo carmine Blue 11.4-13.0 Yellow
Trinitrobenzene colorless 12.0-14.3 orange
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group is lost, resulting in the yellow negatively charged ion denoted as . At
still higher pH, the phenol's hydroxide group loses its proton, resulting in the
red ion denoted as .
In several sources, the structure of phenol red is shown with the sulfur
atom being part of a cyclic group, similar to the structure of phenolphthalein.
However, this cyclic structure could not be confirmed by X-ray crystallography.
(a)
(b) (C)
Fig.2.1. (a) color response of phenol red at different pH value (b) powder
form of phenol red (c) chemical structure of phenol red.
2.4 Determination of pH of indicator dye
pH indicator dyes are weak acid, and their dissociation equilibrium can
be represented as shown[1,20,21]
(2.4.1)
Phenol red(pH indicator)
below pH
6.8
above pH
8.2
6.8 8.2
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Fig 2.2. Dissociation of phenol red
The fraction of the dye existing in the acid (A) and base (B) form depends on the
pH of the solution. The individual spectra of these two forms are different and,
hence, the measured dye spectrum (which is a mole-fraction weighted
combination of the acid and base spectra) changes as the fraction of each form
with pH. The equilibrium constant for equation (2.4.1) is specified by
=
[ ]
[ ]
(2.4.2)
Multiplying both side by –log10
log log log
[ ]
[ ]
+ log + log
[ ]
[ ]
(2.4.3)
Were log ; is the thermodynamic equilibrium constant for the dye
dissociation (e.q.2.4.1), and is a function of temperature and pressure;[A],[B]
are concentration of the acid and base form of the dye in the dye-sample
mixture, respectively; and are activity coefficients of the acid and base
forms of the dye, and a function of temperature, pressure and ionic strength of
solution. a is activity.
This equation is more commonly written as
+ log
[ ]
[ ]
(2.4.4)
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log (2.4.5)
Because includes the activity coefficients, it is no longer only
function of pressure and temperature, but also a function of ionic strength.
For any given pH, the dye spectrum is a mole fraction weighted linear
combination of the acid-only and base-only spectra. Using two-wave length
measurements, one can determine the concentration ratio of base-to-acid form
as below.
= [ ] + [ ] (2.4.6)
According to the additivity of Beer-Lambert law
[ ] [ ] (2.4.7)
[ ] [ ] (2.4.8)
And defining
= (2.4.9)
Simultaneously solving equations (2.4.7) and (2.4.8) yields
[ ] = (2.4.10)
[ ] =
1
(2.4.11)
And the concentration ratio becomes
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[ ]
[ ]
=
1
(2.4.12)
Where
are wave length at which maximum absorbance of pure acid and
base form of the dye respectively.
are absorbance measured at a wave lengt respectively.
is optical path length
T is total dye concentration in solution–dye mixture.
are molar absorption coefficents at wave length for A,B
rspectively.
is absorbance ratio defined by equation(2.4.9)
The base-to-acid concentration ratio can also be obtained by reggration
using spectral A values at more than two wave length of the dye spectra when
available [1]
Equation (2.4.4) and (2.4.12) may now be combined to give
+ log (2.4.13)
Where the e-values are the ratio the molar absorption coefficients defined by
equations (2.4.19 to 2.4.21)
The individual spectra of pure acid and base form of the dye are
obtained from dye solution in extreme pH values (above 3 to 4 units away from
the ) were the dye exists only in the acid or base form [1]. These end points
spectra are used to determine the molar absorption coefficients in equation
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(2.4.13). The two spectra represent the absorptions spectra for the proton
donor ( ) and proton acceptor ( ) forms of indicator. For example, for indicator
phenol red the proton donor form has an absorption maximum ( ) between
400nm and 450nm, while the acceptor form has an absorption maximum
( ) between 500nm and 600nm. Note that at the absorption maximum for
the proton acceptor form, the donor form absorbs light very weakly and it is
vice versa [17].
Consider concentration of dye solution is equal for both pure acid and base
form of the dye. Thus
[ ] = [ ] (2.4.14)
Applying Beer-Lambert for pure acid and base forms the dye at two
wave length .
[ ] (2.4.15)
[ ] (2.4.16)
[ ] (2.4.17)
[ ] (2.4.18)
Where
are maximum absorbance of pure acid and base form of the
dye at wave length respectively.
are absorbance of pure acid and base form of the dye at wave
length respectively.
According to equation (2.4.14) pure acid and base form of the dye are
equal. Applying this equation to equation (2.4.15) to (2.4.17) the ratio of molar
absorption coefficients in equation (2.4.13) becomes
= = (2.4.19)
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= = (2.4.20)
= = (2.4.21)
3.5 Applications of pH measurement
Agriculture
The pH of soils is important since plants grow best within a rather narrow pH
range. The optimum pH varies with each type of plant. In hydroponics, pH
control is even more important than in the soil since too high or too low pH can
cause precipitation of some of the chemicals. Helpful soil bacteria grow best in
slightly acid soil. Plant nutrients form insoluble compounds if soil pH is too
high. Toxic amounts of some metals become available if pH is too low.
Brewing
pH is important for proper ageing and for all stages of the brewing process. A
decrease in pH decreases the solubility of the bitter parts of hops and permits
the use of stronger hops without an increase in harshness. The beer should be
at a pH of 3.9 to 4.1 when bottled, to ensure stability while on the shelf.
Corrosion Prevention
The corrosion of iron occurs below pH 4.3 but a semi-protective layer is formed
above 4.3. Acid soils may be below this level. A more resistant coating is formed
above pH 10.5. The thickness of oxide coatings can be estimated by emf
measurements. If the emf is near that of the oxide the coating is complete and
non-porous but if the emf is near that of the base metal then the coating has
little protective value.
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Dairy Industry
Since milk curdles at a pH of 4.7 it must not be allowed to drop to this value.
The ageing of cheese can be followed by both emf and pH measurements. For
example, a pH of 4.9 is about right for cheddar cheese. Ice cream can be
spoiled by the addition of fruits or juices with too low a pH. The pH of cooling
brines must be controlled to prevent corrosion of the pipes.
Example: Cheese Production
pH measurements are required for soft, fresh cheese and hard, mature cheese.
Cheese contains a large amount of protein and fat so the choice of electrode is
important. Often puncture electrodes are used.
Why are pH values measured?
pH measurements enable both production control during the cheese
fermentation, as well as quality control.
The pH development during the first hours and days is characteristic for every
cheese variety
Dyeing
Processes such as bleaching and dyeing with different types of dye must be
made with a definite pH in order to obtain good results and still not damage
the fabrics. Wool, for example, must be neutralized to a definite pH for effective
dyeing, and in addition, the pH will vary depending upon the nature of the dye
itself. The acid content of the dye also has to be determined which is normally
done by means of a titration using a pH meter to determine the end point.
Electrical Equipment
The pH of feed waters should be controlled in order to prevent pipe and boiler
corrosion. In most cases minimum corrosion occurs between a pH of 7.4 and
8.0. Feed water can be monitored in order to detect certain types of
contamination. The soda lime softening process requires a pH of 9.4 for the
removal of calcium and 10.6 for the removal of magnesium.
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Fermentation Reactions
Each fermentation process requires a specific pH for the best results. A change
in pH with some bacteria even changes the product. The pH during a
fermentation process changes by itself, and must be adjusted periodically. This
maintains optimum conditions and prevents manufacture of unwanted or even
harmful by-products.
Fertilizers
The pH of acid type fertilizers is controlled in order to prevent waste of acid,
and to ensure a more uniform product.
Flour Milling
The quality of flour can be determined by making pH measurements, with the
better flours having a lower pH. The overall range is from 5.9 to 6.5. The
buffering qualities of flours are determined by noting the decrease in pH with
the addition of a measured amount of acid.
Gelatin and Glue Manufacturing
The properties of gelatin and glue vary considerably with the pH during
manufacture. pH needs to be controlled accurately to ensure a consistent
product.
Iron and Steel
pH measurements determine effectiveness of pickling baths and neutralization
of waste pickle baths. Sand used in sand casting can be improved by pH
control. Proper pH makes the sand hold its shape better.
Jam and Jelly Manufacturing
Jams and jellies have narrow ranges for proper gelling. A pH of 3.3 is best for
jelly. At 3.1 it becomes stiff and at 3.5 quite tender. No gelling occurs at all
above 3.5
Laundries
The efficiency of soaps and detergents can be improved by proper pH control.
Undyed cottons can stand a maximum pH of 11, wool about a pH of 10, colored
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cloths 9.6 and silks 9.2. In general, the higher the pH the more efficient the
washing process, but the pH should not exceed the maximum value for the
material being cleaned. Proper pH of starch solutions helps prevent sticking
during pressing.
Leather
Close pH control of leather processing is required to obtain maximum efficiency
without damaging the leather. The pH of tanning and dyeing baths actually
determine the texture and color of the finished product. The deharing process
is normally done at a pH of 12.3, but this must be completely neutralized to
assure good keeping qualities of the leather. Minimum swelling of the collagen
is obtained at pH 4.7.
Meat and Fish Processing
The pH measurement of meat and fish gives a good indication of the keeping
qualities and freshness of the product. This is a typical application for a
puncture electrode.
Metal Finishing
The effectiveness of the pickling and cleansing baths are determined by pH
measurements. The pH of plating baths determines the quality and speed of
the plating process. Some alloys can be plated if very strict pH control is
maintained. The plating thickness can be found during destructive testing by
noting the change in millivolt readings when penetration of a coating has been
accomplished. This is an industry where colorimetric methods cannot be used
directly.
Neutralization
Neutralizing acids or bases is best controlled electrometrically with a pH meter
which indicates the neutral point more precisely than any other method. The
pH meter is especially useful in colored solutions where a color indicator is of
no use.
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Printing
The pH of paper and inks must be controlled to assure proper penetration and
drying of the ink. Too high a pH causes gumminess and too low a pH slows up
the drying process.
Pharmaceuticals
Antibiotics produced from molds are grown at a precise pH. Incorrect pH can
possibly produce a poison rather than a medicine. Many pharmaceuticals must
be prepared using very close pH control.
Sewage
The pH of sewage is controlled to assure efficient coagulation of sludge’s. A pH
of either 3.4 or 7.4 may be used for good results. Also the pH of the effluent
water must be controlled to prevent contamination. The best digestion range is
6.8 to 7.6. The pH of filtration depends upon the chemicals used; for example,
pH 3.4 for ferric chloride and 4.4 for alum.
Swimming Pools
pH levels in swimming pools should be maintained near the neutral range or
slightly alkaline to prevent skin irritations. High pH accelerates deposition of
solid salts in heater lines and filters. Low pH causes corrosion of iron pipes etc.
Tropical Fish Breeding
Expensive tropical fish thrive within definite pH ranges. Each species has its
own best pH environment, which is even more critical during breeding. The
Neon Tetra Fish, for example, prefers water as close to pH 7.0 as possible,
while an Angel Fish requires pH 6.8. The general range for freshwater fish is
pH 6.0 to 8.0. Salt water aquariums should be kept at 8.3. If the pH of the salt
water gets as low as 7.0 the fish become sickly.
Water
The pH of water sources such as rivers, lakes and oceans is measured to study
natural conditions of wildlife. These tests are made by oceanographic
institutes, fish and wildlife service’s and water authorities. pH measurements
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also assist in determining the extent of pollution in domestic and industrial
supplies. In measuring the pH in water, there are two extreme situations. One
is the pH measurement in pure water (boiler feed water), and the other is the
pH measurement in waste water (sewage purification plants).
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CHAPTER THREE
3. Experimental
3.1 Basic of UV/VIS Spectrometer
3.1.1 Instrumental component
There are five essential components required for most absorption
spectrometers. [18]
These are;
a. A source or sources of radiation covering the required wave – length
range.
b. A means for selecting a narrow band of wave length- the device used for
this is called the monochrometer.
c. Facilities for holding the cells or cuvettes containing the sample
solution and the blank in the monochromated radiation beam.
d. A device or devices capable of measuring the intensity of the radiation
beam transmitted through the cells- this is the detector and is usually a
photo detector.
e. A display or output device to record the measured value.
3.1.2 General arrangement experimental set up of spectrometer
The general arrangement of these components for a simple single beam
spectrometer is shown in figure 3.1below
Fig3.1. Basic Experimental set up of spectrometer
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3.2 Reagents
Phenol red ( ), acetic acid ( ), Sodium acetate
( ), disodium hydrogen orthophosphate ( .12 ), hydrochloric
acid ( ), sodium hydroxide( ), de-ionized water. All the chemicals were
purchased from chemistry department of Addis Ababa University.
3.3 Preparation of Buffers
Two types of buffer solutions were prepared to have different values of
pH. Acetate buffer solution (pH 3-6) and phosphate buffer solution (pH 7-11)
3.3.1 Preparation of Acetate Buffer Solution (pH 3-6)
Acetate buffer solution was prepared by
1) 0.1M acetic acid ( ) and
2) 0.1M sodium acetate ( )
Mixing the above two in the following proportion (Table-3.1) to get the required
pH and the exact values were achieved by adjusting with the addition of 6M
HCl or 6M NaOH.
pH Volume of
0.1M acetic acid (ml)
Volume of
0.1M sodium acetate (ml)
3 245.6 4.4
4 211.8 38.2
5 89.3 160.7
6 13.1 236.9
Table-3.1 volume of 0.1M acetic acid and 0.1M sodium acetate mixed to get
Acetate buffer
3.3.2 Preparation of phosphate buffer solutions (pH 7- 11)
Phosphate buffer was prepared by
1) 0.1M disodium hydrogen ortophosphate ( . 12 )
2) 0.1M hydrochloric acid ( )
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3) 0.1M sodium hydroxide ( )
Mixing the above three in the following proportion (Table-3.1), we get the
required pH and the exact values were achieved by adjusting with the addition
of 6M HCl or 6M NaOH.
pH Vol. of phosphate (ml) Vol. of 0.1M (ml) Vol. of 0.1M (ml )
7 189 61 -
8 238.8 11.2 -
9 238.7 11.3 -
10 241.6 - 8.4
11 241.3 - 8.7
Table-3.2 volume of 0.1M disodium hydrogen orthophosphate,0.1M
hydrochloric acid and 0.1M sodium hydroxide to get Phosphate buffer
3.4 Preparation of phenol red stock solution
Phenol red stock solution ( ) was prepared by dissolving
10mg of phenol red in 100ml of distilled water [7].
3.5 preparation caffeine sample solution
1.85x10-4M caffeine sample solution was prepared by dissolving 9.0mg of
caffeine in 250ml distilled water.
3.6 Measurements
The pH of the buffer solutions checked using a digital pH meter (portable pH
meter pH-13) calibrated at 20 ± 2 °C with standard buffers solutions of pH 7.0
and 4.0. Absorbance measurements were made by a UV/Vis spectrometer
Lambda-19. Absorption spectra were taken from 320 to 700nm with 2nm
increment for each form of phenol red solutions.
Before taking absorbance measurements for each phenol red solutions
(having different pH values) the blank is used as a control.
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After blanking, pouring out the blank solution and fill the same quartz
cuvette with sample solution prepared and absorbance of the sample solution
was taken.
3.6.1 Solutions prepared for Absorption spectra measurement of phenol
red as a proton donor and proton acceptor form.
Three solutions were prepared as shown in table-3. First the instrument
was blanked by using blank solution. After blanking the cuvette was filled with
solution in tube one and record its spectrum from 320 to 700nm. Repeat the
process with the solution in tube two.
Blank Tube
One
(ml)
Two
(ml)
One
(ml)
Two
(ml)
A Distilled water 2.0 2.0 1.0 1.0
B 1.41 × 10 phenol red ---- ---- 1.0 1.0
C 0.1M acetate buffer ,pH = 3 8.0 ----- 8.0 -------
D 0.1Mphosphate buffer, pH = 12 ---- 8.0 ------ 8.0
Table-3.3 Amounts added for Absorption spectra measurement of phenol red
as a proton donor and proton acceptor form.
3.6.2 Solutions prepared for Absorption spectrum measurement of phenol
red in buffer solution
Solutions were prepared containing the components indicated in table-3.4.
Blank solutions were scanned and used as a background correction. Then scan
spectrometer and recorded the absorbance of each of sample tubes (in table3.4)
by pouring each in to the same cuvette.
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Blank (ml) Tube (ml)
0.1M acetate buffer, pH = 5
De-ionized water
1.41 × 10 phenol red
8.0
2.0
8.0
1.0
1.0
0.1M acetate buffer, pH = 6
De-ionized water
8.0
2.0
8.0
1.0
1.0
0.1M phosphate buffer, pH = 7
De-ionized water
8.0
2.0
8.0
1.0
1.0
De-ionized water
8.0
2.0
8.0
1.0
1.0
De-ionized water
8.0
2.0
8.0
1.0
1.0
Table 3.4 Amounts added for Absorption spectra measurement of phenol red in
different ph value buffers.
3.6.3 Sample prepared for Absorption spectrum measurement of phenol
red in caffeine solution.
For spectroscopic pH measurement of caffeine, the sample was done by
mixing 40ml of caffeine, 5ml of distilled water and 5ml of phenol red solution.
The absorbance of the sample was measured against mixture of 40ml caffeine
solution and 10ml distilled water baseline (background correction) solutions
which contain no phenol red.
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3.6.4 Sample prepared for Absorption spectrum measurement of phenol
red in packed drinking spring-water.
For spectroscopic pH measurement, five types of packed drinking spring-
water were purchased from the market in Addis Ababa. These packed water
samples were selected based on the difference of the stated pH values in their
bottle cover. The analyses were performed as follows.
1. 40ml spring-water and 5ml distilled water were added to five100ml clean
glass flasks.
2. 5ml phenol red solution (1.41x10-4M) was added to each 100ml glass
flasks containing water samples.
3. Each sample solution was loaded in to the same cuvette one after the
other. The absorbance of each solution was measured against mixture of
40ml spring-water and 10ml distilled water baseline (background
correction) solutions which contain no phenol red.
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CHAPTER FOUR
4. Result and Discussion
4.1 Result
4.1.1 Color response of phenol red solution
Phenol red solution made of phenol red dye powder dissolved in distilled
water has an orange color. It changed to pink when in contact with a pH 8, 11,
12 solutions and changed to yellow when in contact with a pH 3, 5, 6 and
caffeine and water like yes for life solutions and red in the rest four packed
drinking spring-water.
4.1.2 Absorption spectrum of pure acidic and basic form of phenol red
Fig.4.1 shows an absorption spectrum of totally protonated (pure acid form)
and totally deprotonated form (pure basic form) of phenol red in two extreme
pH value buffer solutions. Spectral changes are the results of acid-based
equilibria. These changes are completely reversible with the variation in pH.
When we have two absorbing species which are inter convertible, then their
spectra may overlap. The wavelength at which overlap occurs is called the
isosbestic point, and the absorbance at this wave length is independent of the
position of equilibrium, and depends only on the total amount of the substance
present [18].
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From fig.4.1 phenol red solution in a pH=3.0 buffer solution has maximum
absorbance at = 432.0nm and it has a maximum absorbance at = 558.4nm in
pH=12.0 buffer solution. The value of absorbance at these two wave lengths is
shown in table 4.1.
phenol red
solution
pH A ( = 432.0nm) A ( = 558.4nm)
purely acidic 3.0 0.3051 0.001
purely basic 12.0 0.0297 0.8576
Table 4.1 absorbance of phenol red at = 432.0nm and = 558.4nm in
pH=3.0 and pH=12 buffer solutions
4.1.3 Absorption spectrum of phenol red in different pH value buffer solutions.
Fig 4.2 shows an absorption spectrum phenol red in various pH value of buffer
solution.
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From fig 4.2 the value of absorbance of phenol red at = 432.0nm and =
558.4nm are shown in table 4.2.
pH of buffer measured
by pH-meter.
A A
5.0 0.2532 0.0023
6.0 0.2529 0.0128
7.0 0.2360 0.1166
8.0 0.1312 0.4327
11.0 0.0253 0.7199
pH 5,6,7,8 and11 buffer solutions
4.1.4 Absorption spectrum of phenol red in caffeine solution
Fig 4.3 shows absorption spectrum of phenol red in caffeine solution.
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From fig-4.3 absorbance value of phenol red in caffeine solution at = 432.0nm
and = 558.4nm are shown in table-4.3. The data in table-4.3 is for the three
trials, but the absorption spectrum in fig-4.3 is one of the spectrums among
the three.
Sample Trial A A
Caffeine
1 0.2134 0.0066
2 0.2139 0.0061
3 0.2137 0.0064
caffeine.
4.1.5 Absorption spectrum of phenol red in different pH value spring waters.
Fig 4.4 shows absorption spectrum of phenol red in various types of packed
spring drinking waters.
From fig 4.4 absorbance value of phenol red in different spring water sample
at = 432.0nm and = 558.4nm are shown in table 4.3. The data in table-4.4 is
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for every three trial but the absorbance spectrum in fig-4.4 is only for the first
trial of every sample solution.
Water sample Trial A A
Abyssinia springs 1 0.1194 0.405
2 0.1189 0.4084
3 0.1166 0.2932
Aqua safe 1 0.1367 0.3217
2 0.1357 0.3226
3 0.1353 0.325
Yes for a better life 1 0.2479 0.0107
2 0.2474 0.0109
3 0.2476 0.0114
Origin 1 0.1559 0.2402
2 0.1556 0.2411
3 0.1549 0.2427
Real spring 1 0.1749 0.2042
2 0.1744 0.2053
3 0.1738 0.2063
Table 4.4 absorbance of phenol red at = 432.0nm and = 558.4nm in five
types of packed drinking spring-water.
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4.2 Discussion
4.2.1 Response spectra of pure acid (A) and pure base (B) form of phenol red
Fig 4.1 shows the spectra of phenol red in purely acidic (pH=3) and purely
basic (pH=12) solutions. The absorption spectra of phenol red in purely acidic
and basic solutions resemble to the spectra of phenol red in fig 4.5 [12,13].
Absorbance value of phenol red in buffer solutions measured at two wave
length. Purely acidic phenol red solution has maximum absorbance at a wave
length of 432.0nm while, purely basic phenol red solution has maximum
absorbance at a wave length 558.4nm and this absorption peaks are in
agreement with a statement described in chapter two which is, pure acidic
form of phenol red has an absorption maximum ( ) between 400nm to
450nm, While the pure basic form has an absorption maximum ) between
500nm and 600nm.
The reason for the spectral shift is seen by looking at the chemical
species present at two extremes. Fig-4.5 (left) is the species at low pH (acidic)
conditions. There is some resonant structure in the molecule, but noticed that
the three rings are not in conjugation. For the molecule on fig-4.5 (right) at
high pH (basic) conditions, there is more resonance in the structure since the
three rings are now in conjugation with one another (i.e. the electrons are
delocalized across all three rings).This is the reason why the basic structure
absorb light at a longer wavelength than the acid form –it has a greater degree
of conjugation [14].
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4.2.2 Response spectra of phenol red in buffer solutions
Fig 4.2 shows the spectra of phenol red solution in 3, 5, 6, 7, 8, 11, and pH
value buffer solution. The spectral properties of the dye are strongly
dependent on the pH value. The absorption spectra of phenol red in buffer
solution reassemble to the spectra of phenol red doped sol-gel (fig 4.6)[7]. At
low pH, the absorbance maximum is at 432.0nm, and this peak decreases as
pH of solution increases (fig4.7 (a)), where as the absorbance maximum of
phenol red doped sol-gels is 400nm at low pH. There is only slight red shift of
absorbency maximum in aqueous solution compared to the sol-gels, e.g. the
absorbance maximum are 558.4 and 560 in aqueous solution and in sol-gel,
respectively. In the range of pH 6.00-11.00, the peak at 558.4nm continually
increases with the increased pH (fig4.7 (b)). There is an isosbestic point at
480nm.The absorbance at this point (isosbestic point) is dependent only on the
total indicator concentration because the extinction coefficients of the
protonated and the unprotonated form are equal at this wavelength [15].
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4.2.3 Phenol red determination
The determination of acidity constants by UV spectroscopy is an ideal
method when the compound is too insoluble for potentiometry or when its
value is particularly low or high. Under suitable conditions, it is the
most accurate method, as all measurements being taken in very dilute
solutions. The spectroscopic technique is based on the fact that, for solutions
containing only the fully protonated or the totally nonprotonated species, there
will be an absorption due to both the free base (neutral molecule) and
conjugate acid. The procedures depends upon the direct determination of the
ratio of neutral molecule to ionized species in series of non-absorbing buffer
solutions of known pH [16].
Absorbance of phenol red in purely acidic and basic form at a wave
length of 432.0nm and 558.4nm are shown in table4.1. Thus, from table 4.1 it
is shown that = 0.8576, = 0.001, = 0.0297 = 0.8576. substituting
these values in equations (2.4.18) to (2.4.20). The ratio of molar absorption
coefficients was calculated. The values are
= = = 0.00328 , = = = 2.81 and = = = 0.0346
These e- values are comparable to the values published by Robert-Baldo. The
e-values are extinction coefficient ratios and are either constants or
functions of temperature, which are published together with the pKa
values for indicator like phenol red (Robert-Baldo et al. 1985). According
to Robert-Baldo e = 0.0038 ,e = 2.6155 and e = 0.04718 [15].
Having obtained the spectra of phenol red in various pH buffer solution
(fig.4.2), the value of log
[ ]
[ ]
was determined (table (4.5)) by using the second
term of equation (2.4.13).
Fig.4.8 shows a plot of pH vs log
[ ]
[ ]
. By least square fitting of (e.q.2.4.13)
to the plot, the value of is determined to be 7.8951±0.0270 (intercept
at log
[ ]
[ ]
= 0 ).
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pH(pH meter) A A A [B]
[A]
log
[B]
[A]
5.0 0.2532 0.0023 0.0091 0.0021 -2.6850
6.0 0.2529 0.0128 0.0506 0.0168 -1.7730
7.0 0.2360 0.1166 0.4940 0.1776 -0.7505
8.0 0.1312 0.4327 3.2980 1.3235 0.1217
11.0 0.0253 0.7199 28.4550 655.05 2.8160
[ ]
[ ]
for various pH value of buffer solution.
4.2.4 Spectroscopic pH calculation of caffeine.
sample Trial A A A [B]
[A]
log
[B]
[A]
pH
Caffeine
1 0.2134 0.0066 0.0309 0.0777 -1.1096 6.7855
2 0.2139 0.0061 0.0285 0.0709 -1.1488 6.7463
3 0.2137 0.0064 0.0299 0.0768 -1.1146 6.7805
[ ]
[ ]
for various pH value of caffeine
The mean pH of caffeine measured spectroscopically is 6.7768 ±0.0005 the
precession of the measurement is 0.0005 pH unit. pH of caffeine is 6.9 and the
difference between the measured and the theoretical value is 0.1232 pH unit.
4.2. Spectroscopic pH calculation of packed drinking spring-water.
The pH level is a measure of the hydrogen ion content in water. The
higher the concentration, the more acidic the water is. The lower the
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57
concentration, the more basic the water is. Pure water has a pH of 7. Water
with dissolved gases, minerals and chemicals has a pH range of 5.5 to 8.5.
Natural waters usually have pH values in the range of 4 to 9, and most are
slightly basic because of the presence of bicarbonates and carbonates of alkali
and alkaline earth metals.
Alkalinity in water comes from a high concentration of carbon-based mineral
molecules suspended in the solution. Water with high alkalinity is said to be
"hard." The most prevalent mineral compound causing alkalinity is calcium
carbonate, which can come from rocks such as limestone or can be leached
from dolomite and calcite in the soil. Fresh drinking water should have an
alkalinity level of 20 to 200 milligrams of calcium carbonate per liter of water.
Water can become alkaline by using a water ionizer which separates ordinary
tap water into two forms, acidic and alkaline. The alkaline water is then used
for drinking and the acidic water can be used for washing needs. Alkalinity can
be created by electrically combining certain minerals such as calcium, with
water molecules, which makes the water colloidal. Minerals are derived from
rocks but not all rocks are alkaline in nature. Lakes, rivers and streams can
naturally be alkaline if many minerals are present; creating what is typically
called "hard" water.
Spectrophotometric pH determination is based on the absorbance spectra of
phenol red dye in water samples, it is calculated that phenol red has a
value of 7.8951 and colored protonated and deprotonated forms. From the
absorbance ratio of the unprotonated and protonated forms, the pH can be
calculated using Equation.2.4.13. The calculated pH values of different water
samples for each trial are shown in the table-4.7 below.
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Water sample Trial A A A [B]
[A]
log
[B]
[A]
pH
Abyssinia
springs
1 0.1194 0.4050 3.3920 1.3663 0.1355 8.0307
2 0.1189 0.4084 3.4348 1.3859 0.1417 8.0368
3 0.1166 0.2932 3.3191 1.3331 0.1249 8.0199
Aqua safe 1 0.1367 0.3217 2.3533 0.9104 -0.0407 7.8544
2 0.1357 0.3226 2.3773 0.9206 -0.0359 7.8592
3 0.1353 0.3250 2.4021 0.9311 -0.0310 7.8641
Yes for a better
life
1 0.2479 0.0107 0.0432 0.0142 -1.8469 6.0483
2 0.2474 0.0109 0.0441 0.0145 -1.8372 6.0579
3 0.2476 0.0114 0.0461 0.0153 -1.8164 6.0787
Origin 1 0.1559 0.2402 1.5407 0.5779 -0.2381 7.6569
2 0.1556 0.2411 1.5495 0.5814 -0.2355 7.6596
3 0.1549 0.2427 1.5668 0.5883 -0.2304 7.6647
Real spring 1 0.1749 0.2042 1.1675 0.4318 -0.3648 7.5303
2 0.1744 0.2053 1.1772 0.4355 -0.3610 7.5341
3 0.1738 0.2063 1.1870 0.4393 -0.3572 7.5379
Table 4.7 calculated pH values of five different packed drinking spring-water
The mean pH for each type of packed spring water and the precision in
measurement are shown in table-4.8 below. The average precision of the
measurement is 0.0073 pH unit.
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Water sample mean pH ± SD
Abyssinia springs 8.0292±0.0085
Aqua safe 7.8592±0.0048
Yes for a better life 6.0616±0.0155
Origin 7.6604±0.0039
Real spring 7.5341±0.0038
Table 4.8 calculated pH ± SD values for packed drinking spring-water
4.2.5 Comparison between stated values of pH and spectroscopic
measured pH values of spring water.
Taking together, the 5 comparisons in table -4.9 provide the relationship
( ., 20 ) ( , 25 ) + (0.1869 ±0.5697),
Where our spectrometric measurements are reported on the free hydrogen ion
concentration scale. Since the standard deviation in this relationship (±0.5697)
is attributed to a summation of errors from both stated value in the bottle cover
and spectrophotometric factors, these analyses indicate that the standard error
of each method is smaller than ±0.5697.
Were = ( . , 20 ) ( . , 25 )
Table 4.9 comparison in pH of spectroscopically calculated and the stated
values for spring-water.
Types of spring
waters (Stated value in the
bottle at 25oC)
( by spectroscopy
method at 20oC )
Abyssinia
springs
7.4 8.0292 0.6292
Aqua safe 7.5 7.8592 0.3592
Yes for a better
life
7.0 6.0616 -0.9384
Origin 7.2 7.6604 0.4604
Real spring 7.11 7.5341 0.4241
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The large portion of the apparent noise (or imprecision) observed in this
comparison may in fact be attributable to real differences between the water
sample used in each analysis. Since the comparison were not, initially,
included in the ship’s sampling protocol, the water sample obtained for
spectophotometric analysis were the sample drawn from plastic bottle. The
water sample were purchased from the market, it may be stored for several
months, so the stated value on the cover of the bottle may not be the same.
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61
CHAPTER FIVE
5. Conclusion and Recommendation
5.1 Conclusion
Spectroscopic pH measurement was done using phenol red dye. The use
of pH sensitive dye (phenol red dye) offers an alternative method of measuring
pH. Absorption spectrum of phenol red in different pH value buffer was taken
and of phenol red was characterized. The advantage of using dye is that,
unlike potentiometric methods, the dye measurement depends only on the
molecular properties of the pH sensitive dye and once the dye equilibrium
dissociation constant has been characterized, it eliminates the need for
calibration prior to every measurement. This technique successfully applied for
pH measurement of caffeine and five types of packed drinking spring-water.
5.2 Recommendation
In this thesis pH was measured spectroscopically using single
indicator dye i.e. phenol red dye. For further study it is better to use a mixture
of indicator dyes instead of using single indicator dye to get more precise
result, since mixture of dyes is more sensitive for every pH than a single
indicator dye.
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Reference
1. B. Raghuraman, G.Gustavson,O.C.Mullins and P.Rabbito” spectroscopic
pH measurement for high temperatures, pressure and ionic
strength” AICHE journal.2006; 52(9):3257-3265
2. Buck RP, Rondini S, Convigton AK, Baucke FGK, Brett CMA, Camoes MF,
Miltom MJT, Mussini T, Naumann R. Pratt KW, SpitzerP,WilsonGS,
” Measurement of pH, Definition, standards and procedure. “Pure
and app.chm.2002; 74(11): 2169-2200.
3. Vogel AI.”Text book of Quantitative Inorganic Analysis.” 3rd ed. John
Wiley and sons.1961.
4. Bates R.”Determination of pH: theory and practice.” John Wiley and
sons: New York: 1964
5. Yao W, Byner RH.” Spectrophotometric Determination of Freshwater
pH Using Bromocresol Purple and Phenol red”,Environ.
Sci.Technol.2001 ;35:1197-1201
6. Martz TD, Carr JJ,French CR, DeGrandper MD. “A submersible
Autonomous Sensor for Spectroscopic pH Measurements of Natural
Waters”,Anal. Chem.2003; 75: 1844-1850
7. Enju Wang, Kwok-Fan Chow, Vivian Kwan , Tammy Chin , Crystal Wong
andAndrew Bocarsly “Fast and long term optical sensor for pH based
on sol-gel” Acta 495(2003) 45-50
8. A. Haken Atkins, Nurullah, Guzide Pekcan:”spectroscopic
determination of pka values for some phenolic compounds in
acetonitriale-water mixture.” Acta chim.solv.2006, 53, 214-218.
9. Andor Technolgy, “Absorption/Transmission/ Reflection (ATR)
spectroscopy”. www.andor.com
10. D. Breuer “Molecular spectroscopy in the ultraviolet and visible
range” 2007, vol.10.
11. Wikipedia, the free encyclopedia
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12. Bernard Valeur, “Molecular Fluoresence: Principles and
Applications.” Wiley-VCH,Weinhein Germany(2002).
13. Joseph R. LaKowicz ,”Principles of Flurescence spectroscopy.”,Kluwer
Academic/Plenum publishers, New York(1999)
14. Cllegg, chao, liu, “Physics 5980s optical spectroscopy (Fall 06).”
15. Karsten Friis, Arne korortzinger, and Douglas W.R Wallace
“ spectrosphotometreic pH measurement in the ocean:
Requirements, design, and tasting of an autonomous charge-
coupled device detector system”, American society of limnology and
oceamgraphy.Inc.2004, 126-136
16. Dilek Elamali “calculation of acidity constants of some substituted
thiazole derivetives using DTF and UV spectroscopic methods.”
8/Aralik 2007.
17. Rendian, “Expermental Methods in Modern Biochemistry.”
18. Ronald C Denney and Roy Sinclair “visible and ultraviolet
spectroscopy.” London, 1987
19. Clayton TD, Byrner RH. “Spectrophotometric seawater pH
measurements: Total Hydrogen Ion concentration scale calibration
of m-Cresol Purple and At-sea Result “1993; 40(10):2115-2129.
20. Serghey A. Shapovalov, Valentina L. Koval , Tatyana A. Chernaya, Andrey
Yu. Pereverzev,Nadezhda A. Derevyanko, Aleksandr A. Ishchenko and
Nikolay O. Mchedlov-Petrossyan “Association of Indopolymethine
Cyanine Cations with Anions of Sulfonephthalein and Xanthene
Dyes in Water “ 2005 ;16(2):232-240
21. Stefan Hulth, Robert C. Aller, Pia Engstro¨m and Erik Selander “A pH
plate fluorosensor (optode) for early diagenetic studies of marine
sediments” 2002;47(1): 212–220
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64
Appendix I
Molarity of a solution is defined as:
Molarity =
moles of solute
liters of solution
Molarity =
grames of solution
molar mass solute X liter of solution
Unit of molarity is M
1M=1mole/L
Dilution: diluting the concentration of known concentration
solution by
M V = M V
0.1M solutions were prepared, for buffer preparation, by dissolving solute in
distilled water and by dilution of concentrated solution as follows:
0.1M Sodium acetate (8.2 g/L)
0.1M Di-sodium hydrogen orto-phosphate (35.8 g/L)
0.1M Sodium hydroxide (4 g/L)
0.1M Acetic acid (1.4ml acetic acid (17.86M) diluted to 250ml)
0.1M Hydrochloric acid (12.5 ml HCl (12.5M) diluted to 250ml)
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Appendix II
Standard deviation : A valuable parameter derived from the normal error
curve and expressed by:
=
(x )
N
Where x is a measured result, is the true mean and N is the number of
results in the set. Unfortunately, is never known and the mean derived
from the set of results has to be used. In these circumstances the degrees of
freedom are reduced by one and an estimate of the true standard deviation is
calculated from:
s =
(x )
1
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Spectroscopic pH Measurement Using Phenol Red Dye

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