CHEM-352

A term paper on

INSTRUMENTAL METHODS OF ANALYSIS

i

ChEm-352

A term paper on

Instrumental methods of analysis

Submitted By

Name: Md. Hasib Al Mahbub

Student Id.: 0902045

Group No: 05 Section: A2

Level-3, Term-1

Department Of Chemical Engineering

Submitted To

Dr. Md. Rafique Ullah, Professor, Department Of Chemistry, BUET

Dr. Md. Nazrul Islam, Associate Professor, Department Of Chemistry, BUET

Date of Submission: 11th June, 2013 Partners Id: 0902038

0902040

0902044

Bangladesh University of Engineering and Technology

ii

PREFACE

Analytical chemistry is the art and science of chemical characterization. Instrumental method

of analysis is the part of analytical chemistry. These methods are based on the measurement of

optical, electrical, thermal and other properties. The instrumental methods of analytical

chemistry used in laboratory and research work can be mastered if chemical instrumentation is

studied in its own way. It is always desirable to get almost perfect result with high reliability

within a very short time and this is the reason behind the introduction of digital instruments in

the field of analytical chemistry that has created a revolution in this discipline. In the way of

gradual development and faster adoption of change in the field of science & technology we,

the chemical engineers, may contribute relentless effort and initiative to spread laboratory

researches and innovations to the large scale operation with the nourishment of knowledge and

skill of instrumentation in analysis. For this purpose, this course (Chem-352) has been

introduced to familiarize the students with modern analytical methods associated with

instrumentation which will lead us to future professional skills. The results obtained, by the use

of these instruments are accurate other than the classical methods: At the same time, the

consumption of time is much smaller than the classical methods. Some instruments, which have

the recorder and computer facilities, have created a new era in analytical chemistry. So,

knowledge for operation and handling of those instruments are badly needed and this course is

very much important from this point of view. To understand in depth of the instrumental

analysis, five experiments were performed in the laboratory. This term paper is just an

introduction of instrumental methods of analytical chemistry and covers the most common

instrumental methods. It is hoped that this term paper will provide the reader very basic idea

about the instrumental analysis and modern equipment used in the field of analytical chemistry.

iii

Acknowledgement

At first, I want to express utmost gratitude to the almighty Allah to let me complete this term

paper on the course “Chem-352 Instrumental Methods of Analysis”. I am thankful to my

honorable teacher Dr. MD. Nazrul Islam and Dr. Md. Rafique Ullah who guided and helped

me at every stage to perform a successful substantial work. His constructive comments made

me able to complete this work successfully. I would also like to thank the lab assistants who

helped our group to establish the experimental setup. I thank my group partners for their

continuous co-operation. I am also grateful to my classmates who helped me by sharing their

knowledge and ideas about this experiment. And at last, I state my gratitude to the Department

of Chemistry, BUET.

Md. Hasib Al Mahbub

10th June, 2013

iv

Table of contents

Topics Page no

Preface ii

Acknowledgement iii

Table of contents iv

List of Illustrations v

Introduction 7

Historical background 8

Analytical Chemistry 10

Approaches of analytical chemistry 12

Instrumental Methods of Analysis 14

Applications of Analytical Chemistry 17

Process of Instrumental Analysis 18

Limitations 20

Potentiometric Titration 22

Conductometric Titration 28

pH-metric Titration 38

Spectrophotometric analysis 46

Electrogravimetric Analysis 56

Conclusion 63

Bibliography 64

Appendices 65

Appendix-A: Experimental results for potentiometric titration 66

Appendix-B: Experimental results for Conductometric titration 71

Appendix-C: Experimental results for pH metric titration 75

Appendix-D: Experimental results for Spectrophotometric analysis 81

Appendix-E: Experimental results for Electrogravimetric analysis 83

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List of ILLUSTRATIONs

Topics Page No Figure-1: Block diagram of an analytical instrument showing the stimulus

and measurement of response

15

Figure-2: End point determination of potentiometric titration 24

Figure-3: Potentiometer in Laboratory 24

Figure-4: Conductometric titration curve for strong acid and strong base. 29

Figure-5: Conductometric titration curve for weak acid and strong base 29

Figure-6: Conductometric titration curve for acetic acid titrated using

solution of sodium

30

Figure-7: Conductometric titration curve for the hydrochloric acid – acetic

acid mixture titrated using solution of sodium hydroxide.

31

Figure-8:Graph of Conductance vs. volume of NaOH 32

Figure-9: Schematic diagram of wheatstone bridge 33

Figure-10: Flask type conductance cell 34

Figure-11: Conductometer used in Laboratory; Model (CMD=750) WPA,

Linton Cambridge company

34

Figure-12: pH vs volume of HCl graph for determination of end point 40

Figure-13: Operating principle of glass electrode method 41

Figure-14: Scheme of typical pH glass electrode 41

Figure-15: Combined electrode(glass and calomel) in pH metric Titration 42

Figure-16: The electromagnetic spectrum 47

Figure-17: A typical absorbance spectrum 48

Figure-18: Components of a Spectrophotometer 49

Figure-19: Absorbance vs Wavelength graph 50

Figure-20:Absorbance vs. concentration graph 50

Figure-21: Diagram of a single-beam UV/vis spectrophotometer. 51

Figure-22:Working principle of a spectrophotometer 52

Figure-23: Samples of different concentration 53

Figure-24:Spectrophotometer in the lab 53

Figure-25:Basic construction of an electrolytic cell 56

Figure-26: Schematic Diagram of Electrogravimetric analysis 58

Figure-27: Instrument used in laboratory for electrolysis 59

Figure-28:Copper deposition on the platinum electrode 59

Graph-A.1: EMF of cell vs. Volume of K2Cr2O7 68

vi

Graph-A.2: dV

dE vs. Volume of K2Cr2O7

69

Graph-A.3: 2

2

dV

Ed vs. Volume of K2Cr2O7(ml)

70

Graph-B.1: Specific conductance vs. Volume of HCl graph 74

Graph-C.1: pH vs. Volume of HCl graph 77

Graph-C.2: V

pH

vs. Volume of HCl graph

79

Graph-C.3: 2

2

V

pH

vs. Volume of HCl graph

80

Graph-D.1: Graph of absorbance(A) vs concentration of Iron 82

Table-A: Stages in Instrumental analysis 18

Table 1:Observed and calculated data for volume,cell EMF, first and

second derivative of cell EMF

66

Table 2: Table for volume of NaOH (ml) & conductance (ms) 71

Table 3: Data for Determination of the Strength of Oxalic Acid 71

Table 4: Observed and Calculated data for pH-matric titration 75

Table 5: Table for observed data of absorbance of different concentration 81

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Introduction

Instrumental methods of chemical analysis have become the principal means of obtaining

information in diverse areas of science and technology. The speed, high sensitivity, low limits

of detection, simultaneous detection capabilities, and automated operation of modern

instruments, when compared to classical methods of analysis, have created this predominance.

Professionals in all sciences base important decisions, solve problems, and advance their fields

using instrumental measurements. As a consequence, all scientists are obligated to have a

fundamental understanding of instruments and their applications in order to confidently and

accurately address their needs. A modern, well-educated scientist is one who is capable of

solving problems with an analytical approach and who can apply modern instrumentation to

problems. With this knowledge, the scientist can develop analytical methods to solve problems

and obtain appropriately precise, accurate and valid information. The goal of a chemical

analysis is to provide information about the composition of a sample of matter. In some

instances qualitative information concerning the presence or absence of one or more

components of the sample suffices, in other instances quantitative data are sought. Regardless

of the need, however, the required information is finally obtained by measuring some physical

property that is characteristically related to the component or components of interest.

Instrumental method of chemistry is the part of analytical chemistry. It is defined as the

measurement of a physical property of a sample. Instrument have bought for mankind such

comfort as man can never dreamed of even in a fairly land. Knowledge of analytical chemistry

is highly desirable in all people who use analytical data to arrive at a decision in their work.

Accurate result can be obtained by the use of these instruments other than the classical methods.

At the same time, the consumption of time is much smaller than the classical method. Some

instruments that have the recorder and computer facilities have created a new area in analytical

chemistry, so knowledge for operations and handling of those instruments a badly needed and

this course is very much important from this point of view.

The objective of this course (Chem-352) is to introduce the students with the instruments,

which are now currently used in analytical laboratories. To understand the depth of the

instrumental analysis, five experiments are performed in the laboratory. Those are as follows:

Potentiometric titration

Conductometric titration

8

pH metric titration

Spectrophotometric analysis

Electrogravimetric analysis

It is remarkable that accuracy, precision, reproducibility are the main theme in instrumental

methods of analytical chemistry. So instrumentation has become part and parcel to modern

chemistry.

HISTORICAL BACKGROUND

Much of early chemistry (1661-~1900AD) was analytical chemistry since the questions of what

elements and chemicals were present in the world around us and what are their fundamental

natures is very much in the realm of analytical chemistry. There was also significant early

progress in synthesis and theory which of course are not analytical chemistry. During this

period significant analytical contributions to chemistry include the development of systematic

elemental analysis by Justus von Liebig and systematized organic analysis based on the specific

reactions of functional groups. The first instrumental analysis was flame emissive spectrometry

developed by Robert Bunsen and Gustav Kirchhoff who discovered rubidium (Rb) and caesium

(Cs) in 1860.

Most of the major developments in analytical chemistry take place after 1900. During this

period instrumental analysis becomes progressively dominant in the field. In particular many

of the basic spectroscopic and spectrometric techniques were discovered in the early 20th

century and refined in the late 20th century. The separation sciences follow a similar time line

of development and also become increasingly transformed into high performance instruments.

In the 1970s many of these techniques began to be used together to achieve a complete

characterization of samples. Starting in approximately the 1970s into the present day analytical

chemistry has progressively become more inclusive of biological questions (bioanalytical

chemistry), whereas it had previously been largely focused on inorganic or small organic

molecules. Lasers have been increasingly used in chemistry as probes and even to start and

influence a wide variety of reactions. The late 20th century also saw an expansion of the

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application of analytical chemistry from somewhat academic chemical questions to forensic,

environmental, industrial and medical questions, such as in histology.

1950's - This was a pretty dull field. pH meters and single wavelength spectrophotometers, and

electrochemical techniques were used. Lots of titrations, gravimetric analysis. Some important

work was done to lay the theoretical groundwork. Data was primarily one dimensional.

Experiment Number

1960's - Invention of Gas Chromatography and Atomic Absorption spectrophotometry make

trace analysis possible and reasonably easy. Analysis of ppm and ppb levels of metals and

organics in the environment begins. Text book triples in size. Scanning spectrophotometers

become common. Thus data representations were now two dimensional.

Experiment Graph

1970's - Invention of liquid chromatography and the common use of mass spectrometry for

analytical chemistry begins. GC and AA reach new limits of sensitivity allowing part per

trillion trace analysis. We begin to find virtually everything almost everywhere. Surface

analysis of thin layers becomes common. Analytical chemistry is brought to bear on problems

of the environment, energy, and biological and physiological analysis.

1980's - Continued strides in trace analysis and in identification of trace components through

interfaced (hyphenated) methods. Computers appear to control instruments, manipulate data,

and run experiments Robots appear to conduct complete analytical schemes Multidimensional

data representations add new dimensions to data interpretation It becomes possible to detect

single atoms of many substances. Three dimensional data presentations become important

(MS-MS and 2D NMR).

Experiment 3D Graphic

1990’S - Detection limits continue to drop to the point that one can begin to discuss detection

of single atoms and molecules. Sample preparation is also a major focus as old, time consuming

methods of solvent extraction are replaced with fast, automated procedures. Multi-channel

analysis becomes the major thrust.

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Analytical chemistry Analytical chemistry is the study of the separation, identification, and quantification of the

chemical components of natural and artificial materials.

The goal of analytical chemistry is to provide information is finally obtained by measuring

some physical properties that is characteristically related to the component or components of

interest.

The recent rapid advantages in our knowledge of the physical world parallel a similar rapid

advance in the science of analytical methods. As a chemical engineer, the beginning student

soon becomes aware of the importance of modern analytical methods in all branches of science.

Biologists, geographycists, engineers-all practice chemistry to some extent, and nearly all

chemists practice analytical chemistry. Chemistry is the study of the composition and behavior

of the natural world. Anyone wishing to know more about the composition of substances must

employ analytical methods to determine the kinds and amounts of compounds, elements,

atoms, and sub-atomic particles present in a given sample, as well as to examine the detail

arrangements of the various species. And to study the behavior of materials, analytical methods

must be used before, during, and after certain reactions or “changes”. Analytical chemistry

based on the theoretical principles and instrumentation of all methods. Analytical methods are

usually classified according to the property that is observed in the final measurement process.

There are two large divisions based on the kinds of material with-

Organic

Inorganic

Both of these can be subdivided into two groups:

Qualitative analysis.

Quantitative analysis

Analytical Chemistry

Quantitative Analysis

Qualitative Analysis

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Qualitative analysis: Qualitative analysis is the study of material and the separation and

identification of the component that make up the component that make up the compound and

the mixture.

Quantitative analysis: Quantitative analysis is the determination of the amount of the

constituent of a substance. Quantitative analysis can be further subdivided in terms of the types

of method employed.

There are in general three classes of methods.

Volumetric analysis

Gravimetric analysis

Physio-chemical(instrumental) analysis

Until about 1920 nearly all the analysis were founded on two main properties - volume and

mass. As a consequence gravimetric and volumetric procedures have come to be known as

classical method of analysis.

Volumetric Method

In this method the final determination of the quantity of a substance desired is made by a direct

or indirect measurements of a volume of standard reagent reacting with the substance in a

chemical reaction.

Gravimetric Method

This method requires the separation in a weighable form of the constituent being determined

from a measured amount of the substance containing the desired constituent. The final

measurement in this case is one of the weights.

Physico-Chemical Method

This method involves the final measurements of some physical properties of the composition

of the system as a whole.

Qualitative Analysis

Volumetric Analysis

Gravimetric Analysis

Physico Chemical Analysis

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Approaches of Analytical Chemistry

Most modern analytical chemistry is categorized by two different approaches such as analytical

targets or analytical methods.

By Analytical Targets

Bioanalytical chemistry

Material analysis

Chemical analysis

Environmental analysis

Forensics

By Analytical Methods

Spectroscopy

Mass Spectrometry

Spectrophotometry and Colorimetry

Chromatography and Electrophoresis

Crystallography

Microscopy

Electrochemistry

Traditional analytical techniques

Although modern analytical chemistry is dominated by sophisticated instrumentation, the roots

of analytical chemistry and some of the principles used in modern instruments are from

traditional techniques many of which are still used today. These techniques also tend to form

the backbone of most undergraduate analytical chemistry educational labs. Examples include:

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Titration

Titration involves the addition of a reactant to a solution being analyzed until some equivalence

point is reached. Often the amount of material in the solution being analyzed may be

determined. Most familiar to those who have taken college chemistry is the acid-base titration

involving a color changing indicator. There are many other types of titrations, for example

potentiometric titrations. These titrations may use different types of indicators to reach some

equivalence point.

Gravimetry

Gravimetric analysis involves determining the amount of material present by weighing the

sample before and/or after some transformation. A common example used in undergraduate

education is the determination of the amount of water in a hydrate by heating the sample to

remove the water such that the difference in weight is due to the water lost.

Inorganic qualitative analysis

Inorganic qualitative analysis generally refers to a systematic scheme to confirm the presence

of certain, usually aqueous, ions or elements by performing a series of reactions that eliminate

ranges of possibilities and then confirms suspected ions with a confirming test. Sometimes

small carbon containing ions are included in such schemes. With modern instrumentation these

tests are rarely used but can be useful for educational purposes and in field work or other

situations where access to state-of-the-art instruments is not available or expedient.

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Instrumental methods of Analysis

Instrumental methods of analysis assume the supporting role of an indispensable tool in

advancing the state of knowledge in the fields of inorganic, organic and physical chemistry.

“Instrumental methods” are based on the measurement of optical, thermal and other properties.

Modern instrumental methods generally involve sophisticated instrumentation and they often

possess a high degree of selectivity and hence allow more rapid analysis

Few features clearly distinguish instrumental methods from the classical ones. Some

instrumental methods are more sensitive than classical techniques, but others are not. It is

necessarily true that instrumental procedures employ more sophisticated and more costly

apparatus than classical procedures. An instrument for chemical analysis does not generate

quantitative data instead simply chemical information to a form that is more readily observable.

Thus the instrument can be viewed as a communication device. It accomplishes this purpose in

several steps that includes:

Generation of a signal.

Transmission of the signal to one of different nature (called transduction)

Amplification of the transformed signal and

Presentation of the signal as a displacement on a scale or on the chart of

recorder.

An instrument for chemical analysis does not generate quantitative data but converts chemical

information to a more readily observable form. Sensitivity of the instrument is increased by

amplification of the original signal or its transuded form. Since amplification is commonly

accomplished electronically, instrumental analysis has developed tremendously now a day with

the advance of electronics. These methods are specially applied in industrial laboratories where

economy in time and manpower is gained by their use in routine work.

Summaries of major analytical fields where instrumental methods are extensively applied are

given below-

Molecular analysis

I) Nuclear magnetic resonance ii) Infrared spectroscopy

iii) UV absorption and UV fluorescence

Spectrophotometric Analysis

I) X- ray diffraction and absorption

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ii) Radio tracer techniques iii) Mass spectrometry

Thermal analysis

I) Gas chromatography ii) Electromicroscopy

Elemental Analysis

I) Emission spectrograph ii) Flame photometry

iii) Atomic absorption spectroscopy.

The Nature of Measurement

The steps involved in a chemical determination are the following:

Generation of signal

Detecting and transducing the signal after its interaction with (or generation by) a sample of

interest.

Amplification

Processing or modification (conversion of the signal into a form suitable to operate a read out).

Output or read out (as a line on a chart, pointer deflection on a meter, digital reading etc.). Not

all steps are required in every measurement, nor is every step necessary associated with a

separate entity; instrument modules or subassemblies that perform to or more functions are

common.

Choice Of Method Of Measurement

A measurement begins with an examination of the chemical system of interest. Kinds of

information that must be learned about available samples are:

Physical state

Expected constituents and percentage ranges.

Amounts

Whether a measurement of a property of one constituent will be affected by the presence of

other constituents.

Figure-1: Block diagram of an analytical instrument showing the stimulus and

measurement of response

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Different instrumental methods are-

Spectroscopy

Spectroscopy measures the interaction of the molecules with electromagnetic radiation.

Spectroscopy consists of many different applications such as atomic absorption spectroscopy,

atomic emission spectroscopy, ultraviolet-visible spectroscopy, x-ray fluorescence

spectroscopy, infrared spectroscopy, Raman spectroscopy, nuclear magnetic resonance

spectroscopy, photoemission spectroscopy, Mössbauer spectroscopy and so on.

Mass Spectrometry

Mass spectrometry measures mass-to-charge ratio of molecules using electric and magnetic

fields. There are several ionization methods: electron impact, chemical ionization, electrospray,

fast atom bombardment, matrix assisted laser desorption ionization, and others. Also, mass

spectrometry is categorized by approaches of mass analyzers: magnetic-sector,quadrupole

mass analyzer, quadrupole ion trap, Time-of-flight, Fourier transform ion cyclotron resonance,

and so on.

Crystallography

Crystallography is a technique that characterizes the chemical structure of materials at the

atomic level by analyzing the diffraction patterns of usually x-rays that have been deflected by

atoms in the material. From the raw data the relative placement of atoms in space may be

determined.

Electrochemical Analysis

Electroanalytical methods measure the potential (volts) and/or current (amps) in an

electrochemical cell containing the analyte. These methods can be categorized according to

which aspects of the cell are controlled and which are measured. The three main categories are

potentiometry (the difference in electrode potentials is measured), coulometry (the cell's

current is measured over time), and voltammetry (the cell's current is measured while actively

altering the cell's potential).

Standard Curve

A standard method for analysis of concentration involves the creation of a calibration curve.

This allows for determination of the amount of a chemical in a material by comparing the

results of unknown sample to those of a series known standards. If the concentration of element

or compound in a sample is too high for the detection range of the technique, it can simply be

17

diluted in a pure solvent. If the amount in the sample is below an instrument's range of

measurement, the method of addition can be used. In this method a known quantity of the

element or compound under study is added, and the difference between the concentration

added, and the concentration observed is the amount actually in the sample.

Applications of Analytical Chemistry

Analytical chemistry has played critical roles in the understanding of basic science to a variety

of practical applications, such as:

Biomedical applications

Environmental monitoring

Quality control of industrial manufacturing

Forensic science and so on.

The recent developments of computer automation and information technologies have

innervated analytical chemistry to initiate a number of new biological fields. For example,

automated DNA sequencing machines were the basis to complete human genome projects

leading to the birth of genomics. Protein identification and peptide sequencing by mass

spectrometry opened a new field of proteomics.

Also, analytical chemistry has been an indispensable area in the development of

nanotechnology. Surface characterization instruments, electron microscopes and scanning

probe microscopes enables scientists to visualize atomic structures with chemical

characterizations.

Analytical chemistry is pursuing the development of practical applications and commercial

instruments rather than elucidating scientific fundamentals. This may be an arguable difference

from overlapping science areas such as physical chemistry and biophysics, although there

aren’t any distinct boundaries among disciplines in contemporary science and technology.

However, this aspect may attract many engineers' interest; thus, it is not difficult to see papers

from engineering departments in analytical chemistry journals.

Among active contemporary analytical chemistry research fields, micro total analysis system

is considered as a great promise of revolutionary technology. In this approach, integrated and

miniaturized analytical systems are being developed to control and analyze single cells and

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single molecules. This cutting-edge technology has a promising potential of leading a new

revolution in science as integrated circuits did in computer developments.

Process of Instrumental Analysis

Instrumental analysis may be stated as the application of a process or series of processes to

identify and/or quantify a substance, the components of a solution or mixture, or the

determination of the structures of chemical compounds.

When a completely unknown sample is presented to an analyst, the first requirement is to

ascertain what substances are present in it. Having ascertained to the nature of the constituents

of a given sample, the analyst is then frequently called upon to determine how much of each

component, or of specified components, is present.

Stages of Analysis

The objective and purpose of the analysis has to be sensibly assessed before selecting an

appropriate procedure. A complete instrumental analysis, even for a single substance, involves

a series of steps and procedures. Each one of them has to be carefully considered and assessed

in order to minimize errors and to maintain accuracy and reproducibility. The steps are listed

in the table below:

Table-A: Stages in Instrumental analysis

Stages Examples of procedure

1. Sampling Depends on the size and physical nature of

the sample

2. Preparation of analytical sample Reduction of particle size, mixing for

homogeneity, drying, determination of

sample weight or volume.

3. Dissolution of sample Heating, ignition, fusion, use of solvent(s)

dilution

4. Removal of interferences Filtration, solvent extraction, ion exchange,

chromatographic separation

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5. Sample measurement and control of

instrumental factors

Standardization, calibration, optimization,

measurement of response; absorbance,

emission signal, potential, current

6. Calculation Calculation of analytical result(s) and for

the sample, statistical evaluation of data

7. Presentation of data Printout, data plotting, storage(archiving)

Selecting the Method

Important factors which must be taken into account when selecting an appropriate method of

analysis include

The nature of the information which is sought

The size of sample available and the proportion of the constituent to be determined,

and

The purpose for which the analytical data is required

According to the supplied information and required data generation instrumental analyses may

be classified into four types as follows-

Proximate analysis determines the amount of each element in a sample but not the

compounds that are present.

Partial analysis determines selected constituents in the sample.

Trace constituent analysis is a type of partial analysis that determines specified

components present in very minute quantity.

Complete analysis determines the proportion of each component in the sample.

On the basis of sample size, analytical methods are often classified as follows:

Macro : > 0.1g

Meso (semimicro) : 10-2 g to 10-1 g

Micro : 10-3 g to 10-2 g

Submicro : 10-4 g to 10-3 g

Ultra micro : < 10-4 g

Trace : 102 to 104 µg g-1 (100 – 10000 ppm)

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Limitations

The instruments used in instrumental analysis are expensive enough that most ordinary

laboratories of third world countries cannot afford them. There are some instruments in

our laboratory which are unique of their kind in our country. So, unavailability of

instruments is a major problem for institutions and organizations of least developed

countries.

The special technique of instrumental analysis requires special treatment of samples to

be prepared for the instruments. To prepare these samples a large number of steps have

to be followed.

An instrumental method is a cost effective technique when it is used in large number of

operations. But for an occasional, non-routine analysis it is often better to use classical

method.

Some instruments are so sophisticated and sensitive that any unmindful operation can

cause a notable damage.

Most of the instruments are manufactured in developed countries and the parts, as well

as sales services, are not available everywhere so, a little damage in any parts may

results the invalidation of the instrument as a whole.

21

Experiments performed

EX

PE

RIM

EN

TS

POTENTIOMETRIC TITRATION

pH METRIC

TITRATION

CONDUCTOMETRIC TITRATION

ELECTROGRAVIMETRIC METHOD

SPECTROPHOTOMETRIC METHOD

22

Potentiometric titration

SAMPLE:Fe2+ SOLUTION TITRATED WITH K2Cr2O7 Based on the electrical properties of metals and their ionic solution, several analytical

techniques have been developed. The branch of quantitative analysis involving the estimation

of substances in solution by the measurement of potentials is known as potentiometric analysis

or potentiometric titration. The potential of an electrode relative to a reference electrode varies

in a predictable way with the activities of species, which undergo oxidation or reduction

reactions at the electrode surface. In potentiometric titration, the measurement of potential

serves to locate the equivalents points for titration. Potentiometric titration yields primarily

stoichiometric information. The essential condition to conduct this titration is that the reaction

must be involved the addition or removal of some ion for which an electrode is available.

Theory

Potentiometry deals with the measurement of potential. The cell potential depends on the

potential of the electrodes which in turn depends on the activity of the solution, temperature,

electrodes material, etc. Ordinarily the potentiometric titration involves measuring and

recording the cell potentials after each addition of reagent. This method can be applied to locate

the end point of a titration reaction provided the activity of at least one of the substances

involved can be followed by means of a suitable indicator electrode whose emf is given by

“Nernst equation”. “Nernst equation” can be explained as follows:

When a metal M is emerged in a solution containing its own ion M+ then an electrode potential

is established the value of which is given by the Nernst equation.

RT

E = E0 + --------- ln aM n+

nF

Where,

E= emf of the electrode,

Eo= emf of the electrode at standard condition

n = number of electron transferred during reaction

F= Faraday’s number

Potentiometric titration based upon the quantitative relationship between the emf of a cell is

given by the distribution of potential

Ece l l = E re f + E ind i ca to r+E junc t i o n

23

The reference electrode potential and the junction potential is assumed to remain more or less

constant. So the change of potential of a suitable indicator electrode is observed as a function

of volume added of a titrant of precisely known concentration.

Chemistry Involved

To estimate the amount of Fe++ in a solution a platinum electrode (reference electrode) was

connected to a potentiometer. The solution was titrated against standard K2Cr2O7 solution.

Reaction

Calomel electrode (left): Hg2Cl2 + 2e- = 2Hg + 2Cl -

Platinum (right): Fe++ - e - = Fe+++ , E0 = 0.76v

Cr2O7 - - + 14H+ + 6e - = 2Cr+ + + 7H2O

Emf of the cell is given by,

Ec e l l = E l e f t- E r ig h t (oxidation potential)

This reaction can be divided in to two half-cell reactions. Here ferrous ion is oxidized to ferric

by leaving electrons and these are picked up by dichromate ion. The conversion to ferric ion

from ferrous causes an electromotive force, which can be recorded by the potentiometer. The

endpoint occurs when the entire ferrous ion is converted to ferric ion and this is denoted by the

abrupt change in the magnitude of potential. Here potassium dichromate was used to help the

conversion of ferrous ion to ferric ion. The electrode emf is given by

E = E0 + 0.0591 log[ Fe+ ++]/[ Fe++]

A platinum wire dipped in to the solution picked up the potential and served as an electrical

contact between the electrodes. Plot of potential against volume of titrant added, was used to

find the endpoint of the titration. For more satisfactory deduction of the endpoint slope of the

above curve against the volume of titrant added was constructed. By such a plot it is possible

to obtain a sharp maximum value and hence within narrow limits, the volume of titrant added.

End Point Determination

The end point of a titration may be established any of several ways. The most straightforward

method involves a visual estimation of the midpoint in the steeply rising portion of the curve.

24

Various mechanical methods have been suggested for evaluation of this point; it seems

doubtful, however, these improve the accuracy of estimation greatly.

End point may be determined graphically by plotting the change in potential per unit in volume

of reagent (dE/dV) as a function of the average volume of reagent added. The end point is taken

as maximum in the curve and is obtained by extrapolation of experimental points.

Figure 2: End point determination of potentiometric titration

Instrumentation

The essential instrument used for this experiment was a potentiometer with milivolt scale

connected with saturated calomel electrode and platinum electrode. The potentiometer operated

according to the Poggendroff compensation principle. Ordinarily the titration involves

measuring and recording the cell potential after each addition of reagent. The solution is added

in large increments at the beginning, as the end point is approached the increments are made

smaller.

Figure 3: Potentiometer in Laboratory

Procedure

Procedure mentioned below is followed during performing this experiment.

A stock solution of iron (ferrous) having the strength 0.1(N) is prepared by dissolving

Mohr’s salt.

25

From the stock solution 10ml sample is taken in a cleaned beaker and 20 ml of 6(N) N2SO4

acid is added in this beaker.

It is then titrated with standard 0.1(N) potassium dichromate solution.

The potential (emf) is recorded after addition of each 2ml of titrant, up to the addition of

8ml of K2Cr2O7 solution. Then the potential is recorded after addition of each 0.1 ml of

the titrant up to 13 ml of K2Cr2O7 solution.

The solution must be stirred during titration.

Reference Electrode

In this experiment a standard calomel half-cell electrode is a secondary standard electrode. In

this electrode mercury and calomel are covered with potassium chloride solution of definite

concentration which may be 0.1M, 1M or standard. These electrodes are referred to as the

deciolar, the molar and the standard calomel electrode (SCE) respectively. These electrodes

have the potentials relative to the standard hydrogen electrode (SHE) at 250c of 0.3358 volt,

0.2824 volt and 0.2444 volt respectively. Calomel electrode represented as Hg|Hg2Cl2(S)/C1-

(C). The electrode reaction is,

Hg2Cl2(S) +2e- = 2Hg (liquid) + 2Cl-

Indicator Electrode

The indicator electrode of a cell is one whose potential is dependent upon the activity and

therefore the concentration of a particular ionic species whose concentration is to be

determined. In direct potentiometry or the potentiometric titration of a metal ion, a simple

indicator electrode will usually consist of a carefully cleaned rod or wire of the appropriate

metal (eg. pt). It is important that the surface of the metal to be dipped in to the solution is free

from oxide films or any corrosive product.

Advantages

The main advantage of the method is that it can be applied for turbid, fluorescent,

opaque, or colored solutions or when suitable visual indicators are unavailable or inapplicable.

It provides inherently more accurate data than the indicator method.

It is useful for detecting the presence of unsuspected species in a solution.

26

In acid base titration, the potentiometric method can be employed with advantage where

the solution has a natural color of their own e.g. ink, vinegar and where titration with indicators

is not possible.

Solutions containing more than one halide can be analyzed in a single titration against

silver nitrate using a silver electrode in potentiometric titration.

The end point can be located precisely, even with dilute solutions and also in the

titration of a mixture.

Disadvantages

Its main disadvantages are the more time and attention requirement and the difficulties

associated with the preparation, standardization and storage of standard titrant solution

Result

Required volume of K2Cr2O7 for titration = 4.45 ml

Strength of Fe++ ion in the solution = 0.02225 (N)

Discussion

The aim of this experiment was to determine the end point of the titration. For this purpose

three curves have been plotted from the experimental data. The experimental plots are very

close to the theoretical plot. So the experimental performance was good.

The potentiometric titration of ferrous ferric ion system is based on the fact that the transfer of

electron takes place rapidly and reversibly when the reduced and oxidized forms have the

essentially identical structure. EMF of the cell constituted by calomel and platinum electrode

depends upon the potential of platinum electrode since potential of calomel electrode is

constant. Potential of platinum electrode in turn depends on the ratio [Fe+3] /[Fe +2] as seen

from Nernst equation. After each addition of K2Cr2O7 this ratio changes and EMF changes

gradually. The critical problem in a titration is to recognize the point at which the quantities of

reacting species are present in equivalent amounts-the equivalence points. The typical titration

curve can be followed point by point, plotting as ordinate successive values of the cell EMF

Vs the corresponding volume of titrant added.

After adding K2Cr2O7 in accurately measured increments a sigmoid curve is obtained which

is similar to the theoretical titration curve for oxidation - reduction reaction. It is seen that EMF

27

changes slowly during the initial portion. In the vicinity of equivalence point a sharp rise is

observed as shown by the vertical portion of the graph. The midpoint in the steeply rising

portion of the curve is then estimated visually and taken as the end point.

A second approach was to calculate the change in potential per unit change in volume of reagent

[that is, ΔE / ΔV] as has been done in Table 1. A plot of this parameter as a function of average

volume leads to a sharp maximum at the end point. The third approach shows that the volume

can be fixed more exactly by estimating the point where the second derivative of the voltage

with respect to volume (that is Δ2E / ΔV2) becomes zero.

All of the methods considered are based on the assumption that the titration curve is symmetric

about the equivalence point and that the inflection in the curve thus corresponds to that point.

This assumption is perfectly valid, provided the participants in the chemical process react with

one another in an equimolar ratio, and also provided the electrode process is perfectly

reversible. Ordinarily, the change in potential within the equivalence point region of these

curves is large enough so that a negligible titration error is introduced if the midpoint of the

steeply rising portion is chosen as the end point. Only when unusual accuracy is desired or

where very dilute solutions are employed must account be taken of this source of uncertainty.

28

Conductometric titration

SAMPLE: HCl ACID & CH3COOH ACID TITRATED WITH NaOH SOLUTION

Titration in which conductivity measurements are made use of, in determining the end point of

acid base reaction, some displacement reactions or precipitation reactions are called

Conductometric titration. The conductance of a solution is a measure of how well a solution

carries a current. It is reciprocal of the electrical resistance and has the unit of ohm-1 or mho. It

is represented by LS.

Theory

The positive ions migrate through the solution towards the cathode, where they pick up

electrons. Anions migrate towards the anode, where they deliver electrons. The net result is a

flow of electrons through the solution. Thus the electrical conductance of a solution at a

constant temperature depends upon the number of displaceable electrons in the solution and

their mobility under the influence of an applied electromotive force. The addition of an

electrolyte to a solution of another electrolyte under conditions producing no appreciable

change in volume will affect the conductance of a solution according to whether or not ionic

reactions occur. If no ionic reaction take place such as in the addition of one simple salt to

another (e.g. KCl to NaNO3), the conductance will simply rise. If ionic reaction occurs, the

conductance may either increase or decrease, thus in the addition of base to a strong acid, the

conductance decreases owing to the replacement of the hydrogen ion of high conductivity by

another ion of lower conductivity. This is the under lying principle of Conductometric titration.

Let us consider the conductance of a solution of a strong electrolyte A+B- will change upon the

addition of a reagent C+D-, assuming that the cation A+ reacts with the anion D- of the reagent.

If the product of the reaction, AD is relatively insoluble or only slightly ionized the reaction

may be written as:

A+B- + C+D- = A+D- + C+B-

Thus in the reaction between A+ ions and D- ions, the A+ ions are replaced by C+ ions during

the titration, as the titration proceeds the conductance increases or decreases, depending on

whether the conductivity of the C+ ions is greater or less than that of the A+ ions.

29

Chemistry of strong acids & weak acids with strong bases

End point

Conductance

Volume of base

Figure 4: Conductometric titration curve for strong acid and strong base.

Reaction: (H+ + Cl-) + (Na+ + OH-) = (Na+ + Cl-) + H2O

Again when a weak acid like acetic acid is titrated against a strong alkali like Sodium

Hydroxide (NaOH) a curve of the following type is obtained. The initial conductivity of the

solution is low because of the poor dissociation of the weak acid. On addition of alkali, highly

ionized sodium acetate (CH3COONa) is formed. The acetate ions at first tend to suppress the

ionization of acetic acid still further due to common ion effect. But after a while the

conductivity begins to increase because the conducting power of highly ionized salt exceeds

that of the weak acid.

Reaction: CH3COOH + Na+ + OH- = CH3COO- + Na+ + H2O

End point

Conductance

Volume of base

Figure 5: Conductometric titration curve for weak acid and strong base.

Immediately after the end point, further addition of NaOH introduces the fast moving hydroxyl

ions and the conductivity value shows a sharp increase. The point of intersection of the two

curves therefore gives the end point.

30

Consider the titration of solution of weak acid, such as acetic acid CH3COOH, using a solution

of strong base, NOH. As we know, the weak acids, as well as other weak electrolytes, are

dissociated into very small extent and they exist in solution essentially in form of the neutral

acid molecules. When a solution of NaOH is added the reaction

occurs and, as is seen, the undissociated molecules of acetic acid are transformed into

dissociated molecules of potassium acetate. The changes are accompanied by increase in

conductivity of the hydroxide solution.

Figure 6: Conductometric titration curve for acetic acid titrated using solution of

sodium.

It should be noted, however that an initial decrease in a conductivity of the solution may be

observed after addition of the first drops of titrant. This minor important effect is related to

neutralization reaction of the protons resulting from dissociation and existing even in a solution

of the weak acid. Thus, an mild increase in conductivity of a titrated solution is observed until

the equivalence point is reached, at which we have a solution of sodium acetate, CH3COONa.

If an excess of titrant, that is the potassium hydroxide solution, is added a sharp increase in

conductivity is observed. This distinct difference in a rate of increase is related to the fact that

the excess OH- anions, as well as the protons, exhibit particular mechanisms of charge

migration. More detailed inspection of the Conductometric titration curve presented in figure

3.4. Indicates that the equivalence point is less sharp than that observed for the strong acid.

Thus, it should be localized as the intersection point of two lines determined by two section of

31

the Conductometric curve. The slope of the first part of the Conductometric curve is dependent

on strength of the acid. It means that it is positive for very weak acid only.

The method of Conductometric titration is thus well adapted to the estimation of mixtures of

acids of differing strengths. When a mixture of strong and weak acid is titrated a plot of

conductance against alkali added takes form of Fig.11.

Figure 7: Conductometric titration curve for the hydrochloric acid – acetic acid mixture

titrated using solution of sodium hydroxide.

As is seen, the Conductometric titration curve is a combination of the diagrams obtained during

the titration of strong and weak acid respectively, where the first endpoint corresponds to a

neutralization of the strong acid present in the sample and the second one is associated with a

neutralization of the weak acid in the solution under investigation. The volume of the alkali

consumed by the latter is given by a difference of the respective volumes.

32

CHEMICAL REACTIONS

Sodium Hydroxide, Acetic acid and Sulfuric acid ionize in the solution in the following forms:

NaOH = Na+ + OH-

HCl = H+ + Cl-

CH3COOH = CH3COO- + H+

Then the following reactions occur:

Na+ + Cl-= NaCl

H+ + OH- = H2O

NaOH + HCl = NaCl + H2O

NaOH + CH3COOH = CH3COONa + H2O

Determination of End Point

Conductometric measurements provide a convenient means for the location of end points in

titrations. To establish a Conductometric end point, sufficient data are needed to define the

titration curve. After correction for volume change, the conductance data are plotted as a

function of titration volume. The point of intersection of two lines is taken as the equivalence

point.

Figure 8: Graph of Conductance vs. volume of NaOH

33

Instrumentation

1. Instrument used in the lab: CMD 750 WPA conductivity meter

2. An electro chemical cell with a pair Pt electrodes

Conductometer:

The conductivity of an electrolyte solution is a function of the concentration of ions in the solution.

Because of this, each solution has a resistance that can be measured with conductivity cells.

L=1/ρ

where, L= conductance and ρ= resistivity of a substance. These are characteristic quality of a

material. ρ of a material depends upon its concentration and other properties.

Conducting sensors are constructed of an insulating material imbedded with metallic pieces. The

metal pieces, serving as sensing elements, are placed at a fixed distance apart and make contact

with the solution. The cell constant, K, for a particular conductivity cell is determined by the

geometry of the cell. The value of K can be estimated by the following equation:

where d is the distance between the sensing elements and A is the surface area of one sensing

element exposed to the solution.

When measuring the resistance of a solution it is impractical to use an ohmmeter, since passing a

current through the solution will lead to errors, such as heating and polarization. Using AC

potential eliminates the latter problem, but an AC Wheatstone bridge minimizes both of the above

problems. In Figure 13, Rf1 and Rf2 are fixed resistors of known resistance; Rb is a resistance

box in which the resistance can be changed very precisely over a wide range; and Rm is the

resistance to be measured, in this case the conductivity cell with the solution of interest.

Figure 9: Schematic diagram of Wheatstone bridge

A

dK

34

The bridge works correctly when V2 = V4 and therefore Rm = Rb(Rf1/Rf2). If the two fixed

resistors are chosen to be identical, then Rm is equal to the reading of the variable resistance

box Rb. (2)

Conductance cell

Different types of cells are used for conductance measurement. A cell which is commonly used

is shown in figure-14. The container is a glass vessel of Pyrex or other resistance glass (A),

which carries two thick platinised platinum foils (PP), securely fixed so that their distances are

not altered. Two metallic wires (BB) sealed to the platinum foils and protected by glass tubes

serve as the leads for connecting to the metallic bridge.

Figure 10: Flask type conductance cell

Figure 11: Conductometer used in Laboratory; Model (CMD=750) WPA, Linton

Cambridge company

35

Method of Operation

20 ml mixture of HCl and acetic acid was taken in a conductivily cell.

Secondary standard NaOH was added in an amount of 0.5 ml each time from a burette.

For every addition the corresponding reading of specific conductances were taken.

A graph of specific conductance vs Volume added of NaOH was plotted

From the graph the end point strength of HCland acetic acid was determined

Advantages

Sometimes the end point in an acid-base titration reaction cannot be determined using a

chemical indicator. This may occur for the following reasons:

a) No suitable indicator

b) No stable end point

c) The solution is too dilute

d) The solution is turbid (suspension of small particles)

e) The solution is brightly or strongly colored.

The end point can still be determined by following some suitable characteristic of the solution

that is conductance. Conductometric titration is suitable for-

Very weak acid, such as boric acid and phenol, which cannot be titrated potentiometrically

in aqueous solution, can be titrated conductometrically.

Mixture of certain acids can be titrated more accurately by Conductometric titration than

by Potentiometric pH methods. Thus mixtures of HCl or any other strong acid and acetic

acid or any other weak acid of comparable strength can be titrated with a weak base (e.g.

aqueous NH3) or with a strong base (e.g. NaOH) reasonably satisfactory end points are

obtained.

The apparatus required for making conductance measurements and performing titration is

generally inexpensive and basically simple in detail. For these reason the measurement of

conductance finds wide acceptance in industry as an analytical tool, both in the laboratory

and in process control.

Conductometric methods may be applied where visual or potentiometric methods fails to

give result owing to considerable solubility or hydrolysis at the equivalent point very week

36

point , such as boric acid and phenol which cannot titrated potentiometrically aqueous

solution, can be titrated conductometrically.

Another important advantage is that this method is as accurate in dilutes in more

concentrated solution.

This method can be employed with colored solution.

Disadvantages

Conductometric titration may be used only to detect the end point of a comparatively simple

system in which there is no excessive amount of reagent present. Thus many oxidation titration

which require the presence of relatively amounts of acid are not suitable to conductometric

titration.

The plot obtained in Conductometric titration cannot use for solutions of different

concentrations of the same species.

Application

Conductometric Titration is applied for determining the level of contamination of water or any

other chemical species, as the conductance of pure species of different substances are constant.

Result

Strength of `HCl= 0.19 N

Strength of CH3COOH = 0.212 N

Percentage of HCl= 35.28%

Percentage of CH3COOH = 64.72%

Discussion

In this experiment strong acid HCl and weak acid CH3COOH is titrated at the same time with

strong base NaOH solution so there should be two end points. According to Graph-C.1

(Appendix-C) two definite end points are observed. The first of these indicate neutralization of

HCl acid (strong acid) and second of these indicate neutralization of CH3COOH acid (weak

acid). So the result is satisfactory.

In this experiment, the Conductometric titration curve for a weak acid (CH3COOH) and strong

acid (HCl) with strong base (NaOH) is depicted.

This curve shows a sharp decrease in conductance till the end point of the reaction between

HCl and NaOH is reached. After this first equivalent point the conductance rise gradually up

37

to the second end point, which indicate the neutralization of CH3COOH acid with NaOH

solution and then the conductance rise sharply again as NaOH introduces the fast moving

hydroxyl ion. This characteristic of the curve is described in brief below –

In the titration of a strong acid (HCl) with base (NaOH), the conductivity of the solution

containing acid decreases and continues to fall with every subsequent addition of alkali till the

end point is reached. After the equivalent point, the further additions of NaOH solution results

in an increase of conductance since the hydroxyl ions are no longer removed in the chemical

reaction in the form of freely ionized water. The point of minimum conductance therefore

coincides with the end point of the titration.

In the titration of a weak acid (CH3COOH) with strong base (NaOH solution), on addition of

alkali highly ionized CH3COONa is formed. The acetate ions at first tend to suppress the

ionization of an acetic acid still further. But after a while the conductivity begins to increase

because the conducting power of highly ionized salt exceeds that of the weak acid. Immediately

after the end point, further addition of sodium hydroxide introduces the fast moving hydroxyl

ions and the conductivity value shows a sharp increase.

In performing the Conductometric titration the following things should be taken under

consideration –

The relative change of conductance of the solution during the reaction and upon the addition

of an excess of reagent largely determined the accuracy of the titration. (Under optimum

conditions this is about 0.5%). Large amount of foreign electrolytes which do not take part

in the reaction, must be absent since those have a considerable effect upon the accuracy.

The accuracy of the method is greater when the acute angle of interception is more and the

points on the graph lie on a straight line.

It is necessary to keep the temperature constant throughout the experiment.

In acid alkali titration, the titrant should be about 10 times stronger than the solution to be

titrated so that the volume change is as little as possible.

Platinised platinum electrode should be used to minimize polarization effects.

The size and separation of the electrodes will be governed by the change of conductance

during the titration. For low conductance solutions (e.g. when solution is extremely dilute)

the electrodes should be large and should be placed closely.

It is not necessary to know the cell constant since relative values are sufficient to permit

locating the equivalence point, &

During titration, if the cell is connected to a pen recorder in addition to the conductivity

meter a plot of change in conductance with addition of reagent can be drawn automatically.

38

P H metric titration

Sample: HCl Acid Titrated With NaOH

Introduction

The purpose of the experiment was to perform the pH metric titration of sodium carbonate by

using HCl as titrant. In an acid base neutralization reaction the neutralization point exhibits a

sharp change in pH with small change in the volume. The process of determining the end point

by monitoring the sharp change in pH is known as pH titration.

Theory

Acidity and basicity of a solution may be described in terms of its hydrogen ion concentration.

Sorenson in 1909 devised a measure of hydrogen ion concentration which is convenient and at

the same time compact. He introduced the term pH and designed it as the negative value of

logarithm of hydrogen ion concentration.

pH=-log10 [H+]

With the introduction of the concept of activity the correct definition is

pH= -log aH

Where, aH = hydrogen ion activity

In acid base titration pH of the test solution changes with every addition of titrant. This change

is very sharp in the vicinity of end point. So, end point can be located from the change of pH.

This is the underlying principle of pH titration.

A variety of electrodes available for pH measurement the hydrogen gas electrode, the glass

electrode, the antimony electrode, the quinhydrone electrode etc. The arrangement of the cells

which are pH responsive is known as pH cell. In this course glass electrode was chosen as to

construct a pH responsive cell.

The glass electrode, actually a membrane electrode comprises a thin walled bulb of pH -

responsive glass sealed to a stem of non-pH-responsive, high resistance glass.

The potential of the glass electrode is given by,

E = Ee + .000198T (pH -pHr)

Where,

pHr= pH of the reference solution contained within the glass bulb

Ee =combination of junction potential

T =Temperature in degree Kelvin

39

From the above equation it can be seen that there exists a relation between the pH of a solution

and the electrode potential of an electrode that is dipped in the solution.

The conventional pH value may be defined in an operational manner, as follows,

pH = pHs + (E-Es).000198T.

Where,

E= emf of pH cell which cantains an unknown solution.

E s= emf of pH cell which contains standard referance of known pH.

So, change in electrode potential following a change in hydrogen ion concentration shows the

change in pH. A pH meter is added to the electrodes to convert the change in electrode-potential

in terms of pH and display it in the monitor)

CHEMISTRY INVOLVED

The combination electrode was dipped in the Na2CO3 solution taken in a beaker. Since the

solution inside the glass tube was maintained at a constant pH potential of the calomel electrode

inserted into it was constant and so was the potential between the HCl solution and inner surface

of the glass tubes and test solution varied with the change of hydrogen ion concentration of the

solution. With every addition of HCl from the burette into the solution of Na2CO3, pH varied

and consequently emf of the combination electrode changed. This emf was recorded by an in

built potentiometer of the pH meter. This potentiometer reading was automatically converted

electrically to a direct reading of pH of the solution.

Reactions involve in the process are-

H+ (solution) + Na+ (glass) = Na+ (solution) + H+ (glass)

HA = H+ + A-

HA

AH

K

HA

AlogH log K log

When 50% acid is neutralized then the concentration of acid and salt will be equal.

In that case, pK = pH

The reaction involved here is completed by two steps. They are as follows:

Na2CO3 + HCl NaHCO3 + NaCl

NaHCO3 + HCl NaCl + H2CO3

CO2 + H2O

40

Determination of end point

pH metric process to determine the end point of titration is an important method. By this

method, for an acid-base titration pH changes gradually as titrant is added. Around the

equivalence point, however the pH changed abruptly. The rate of change is greatest at the

equivalence point. This can be seen with the plot pH vs volume of HCl plot. A plot dpH/dV vs.

volume also shows a sharp peak at the equivalence point.

Figure 12:pH vs volume of HCl graph for determination of end point

To determine the more accurate end point d2pH/dV2 graph is also plotted.

Instrumentation

The essential instruments used for pH measurements are.

(i) pH meter ( pH/ion meter. 150 CORNING) an electronic digital volt meter scaled to read

pH directly

(ii) A combination electrode connected to the pH meter.

The combination electrode contains an indicator electrode (a thin glass electrode) and a

reference electrode (calomel electrode) combined in a single unit. The thin glass bulb is filled

with HCl of const. pH and carry the calomel electrode.

Glass electrode

Glass electrode is potentiometric sensor made from glass of a specific composition. It is an Ion-

selective electrode (ISE) with a transducer (sensor) which converts the activity of a hydrogen

ion (H+) dissolved in a solution into an electrical potential which can be measured by a

voltmeter or pH meter. The voltage is theoretically dependent on the logarithm of the ionic

activity, according to the Nernst equation. The sensing part of the electrode is usually made as

an ion-specific membrane, along with a reference electrode. Ion-selective electrodes are used

in biochemical and biophysical research, where measurements of ionic concentration in an

aqueous solution are required, usually on a real time basis.

41

Figure 13: Operating principle of glass electrode method

Construction

A glass electrode consists of an electrode membrane that responds to pH, a highly isolating

base material to support the unit, solution inside the glass electrode, an internal electrode, a

lead wire, and a glass electrode terminal. The most critical item in this system are-

Figure 14: Scheme of typical pH glass electrode

1. a sensing part of electrode, a bulb made from a specific glass

2. sometimes the electrode contains a small amount of AgCl precipitate inside the glass

electrode

3. internal solution, usually 0.1M HCl for pH electrodes or 0.1M MeCl for pMe electrodes

4. internal electrode, usually silver chloride electrode or calomel electrode

42

5. Body of electrode, made from non-conductive glass or plastics.

6. reference electrode, usually the same type as 4

7. Junction with studied solution, usually made from ceramics or capillary with asbestos or

quartz fiber.

Combined Electrode

Typical modern pH probe is a combination electrode, which combines both the glass and

reference electrodes into one body. The bottom of a pH electrode balloons out into a round thin

glass bulb. The pH electrode is best thought of as a tube within a tube. The inside most tube

(the inner tube) contains an unchanging saturated KCl and a 05M HCl solution. Also inside the

inner tube is the cathode terminus of the reference probe. The anodic terminus wraps itself

around the outside of the inner tube and ends with the same sort of reference probe as was on

the inside of the inner tube. Both the inner tube and the outer tube contain a reference solution

but only the outer tube has contact with the solution on the outside of the pH probe by way of

a porous plug that serves as a salt bridge.

The measuring part of the electrode, the glass bulb on the bottom, is coated both inside and out

with a ~10nm layer of a hydrated gel. These two layers are separated by a layer of dry glass.

The silica glass structure (that is, the conformation of its atomic structure) is shaped in such a

way that it allows Na+ ions some mobility.

Figure 15: Combined electrode (glass and calomel) in pH metric Titration

The metal cations (Na+) in the hydrated gel diffuse out of the glass and into solution while H+

from solution can diffuse into the hydrated gel. It is the hydrated gel, which makes the pH

electrode an ion selective electrode. H+ does not cross through the glass membrane of the pH

43

electrode, it is the Na+ which crosses and allows for a change in free energy. When an ion

diffuses from a region of activity to another region of activity, there is a free energy change

and this is what the pH meter actually measures. The hydrated gel membrane is connected by

Na+ transport and thus the concentration of H+ on the outside of the membrane is 'relayed' to

the inside of the membrane by Na+.

Procedure

Procedure mention mentioned below is followed during performing this experiment:

(1) 20 ml of NaCO3 solution is taken into a beaker

(2) A combined set-up of glass electrode and calomel electrode is dipped into the above

said beaker. This glass set-up of electrodes is connected to pH meter.

(3) 0.2 (N) HCl solution is added in the beaker containing Na2CO3 solution by using a

burette and corresponding pH reading is taken from pH meter.

(4) Each time .2 to .3 ml of 0.2 N HCl acid is added until end point is reached.

Advantages

pH metric titration has some unique advantages. Some of them are given below:

In colored solutions, Indicators cannot detect endpoints where as pH metric titration is

not subjected to this limitation.

Indicators for any acid-base titration must be chosen so that the pH at which the

indicator changes color corresponds more or less to the pH of the solution at the equivalence

point. Hence some information is required concerning the relative strength of the acid base

involved. pH metric titration always yield the equivalence point whether this point comes

exactly at the neutral point or on the acid or basic acid.

With indicators it is some-times impossible or difficult to titrate a polybasic acid or a

mixture of a strong acid and weak acid with a base. In such cases pH metric titration gives

accurate result.

Oxidation titration can be carried out by this method

In precipitation reaction, pH metric titration can be employed without any difficulty.

Titration of weak acid-weak base, strong acid-weak base, weak acid strong base, when

indicator choosing is difficult or indicator does not show distinct color change, then by pH

metric titration end point of the titration can be detected easily.

44

Disadvantage

The main limitations of pH-metric titration arise when strong, alkaline solution is used. In this

case the glass electrode is affected by the alkaline solution. The glass electrode shows very

little response to divalent actions. Special care should be taken in handling this glass electrode,

as it is very expensive otherwise it can be damaged or broken.

Results

Strength of Na2CO3 solution = 0.05N

First dissociation constant, pK1= 7.7

Second dissociation constant, pK2= 5.25

Discussion

The purpose of this experiment was to study the nature of Na2CO3-HCl titration with the

measurement of pH and to plot the titration curve of pH vs. milliliters of titrant added. The pH

values were directly obtained from a pH meter.

The pH curve shows pH changes sharply twice during the process. This indicates that the

overall reaction takes place in two steps. pH level of Na2CO3 solution is maximum before

adding HCl. As HCl is added, pH of the resulting solution decreases gradually. At the first

vertical portion of the curve, pH decreases sharply. This point indicates the completion of 1st

step of the overall reaction:

Na2 CO3 + HCl = NaHCO3 + NaCl.

Complete conversion of CO3-2 to HCO3

- takes place at this point. pH value corresponding to

this point represents the 1st ionization constant of Na2CO3. As the process continues pH

decreases slowly again until the 2nd steeper portion is reached. At this 2nd end point there is

another sharp fall of pH. This indicates the completion of 2nd part of the overall reaction.

NaHCO3 + HCl = NaCl + H2O + CO2

This point indicates the conversion of HCO3- to H2CO3 followed by dissociation of carbonic

acid to H2O and CO2.

To locate the end points accurately, a first derivative plot of Δ pH/ΔV vs. volume is constructed

from the pH curve. This plot shows two sharp peaks, which indicates the locations of two

endpoints. For more accuracy a second derivative Δ2pH/ΔV2 Vs V is plotted.

45

This point indicates the conversion of HCO3- to H2CO3 followed by dissociation of carbonic

acid to H2O and CO2.

From the second derivative vs volume curve the pK values for HCl can be determined,

According to the Henderson’s equation:

pH=pK+ log acid

salt

So when 50% of the acid is neutralized and forms salt then pH= pK, so from the curve the pK1

and pK2 values could be determined.

Temperature changes in the measured liquid affect both the response of the measurement

electrode to a given pH level , and the actual pH of the liquid.

Application

Titration of colored solutions can be done in this method as indicators are not suitable

for those solutions.

Electrolytes of blood and their concentration can be determined in this method

pH of amino acid and its properties in the body can be measured through it.

46

SPECTROPHOTOMETRIC ANALYSIS

Sample:Fe3+ solution

Introduction

Spectrophotometric analysis is the most sophisticated analytical method & usually used to

analyze in ppm or ppb range both qualitatively & quantitatively. In this method the ratio of the

intensities of the incident and the transmitted beams of light is measured at a specific

wavelength by means of a detector (a photocell). The absorption spectrum also provides as a

fingerprint for qualitatively identifying the absorbing substance. The amount of visible light or

other radiant energy absorbed by a solution is measured since it depends on the concentration

of the absorbing substance; it is possible to determine quantitatively the amount present.

Basic principles of spectrophotometry

An absorbance spectrophotometer is an instrument that measures the fraction of the incident

light transmitted through a solution. In other words, it is used to measure the amount of light

that passes through a sample material and, by comparison to the initial intensity of light

reaching the sample, they indirectly measure the amount of light absorbed by that sample.

Spectrophotometers are designed to transmit light of narrow wavelength ranges (see Figure 1

the electromagnetic spectrum). A given compound will not absorb all wavelengths equally–

that’s why things are different colors (some compounds absorb only wavelengths outside of

the visible light spectrum, and that’s why there are colorless solutions like water). Because

different compounds absorb light at different wavelengths, a spectrophotometer can be used to

distinguish compounds by analyzing the pattern of wavelengths absorbed by a given sample.

Additionally, the amount of light absorbed is directly proportional to the concentration of

absorbing compounds in that sample, so a spectrophotometer can also be used to determine

concentrations of compounds in solution. Finally, because particles in suspension will scatter

light (thus preventing it from reaching the light detector), spectrophotometers may also be used

to estimate the number of cells in suspension.

When studying a compound in solution by spectrophotometry, it is put in a sample holder called

a cuvette and placed in the spectrophotometer. Light of a particular wavelength passes through

the solution inside the cuvette and the amount of light transmitted (passed through the

solution—Transmittance) or absorbed (Absorbance) by the solution is measured by a light

meter. While a spectrophotometer can display measurements as either transmittance or

absorbance, in chemistry researchers are usually interested in the absorbance of a given sample.

47

Figure 16: The electromagnetic spectrum. Visible light (400-700 nm) constitutes only a

small portion of the spectrum that ranges from gamma rays (less than 1 pm long) to

radio waves that are thousands of meters long

Because other compounds in a solution (or the solvent itself) may absorb the same wavelengths

as the compound being analyzed, the absorbance of test solution is compared to a reference

blank. Ideally, the reference blank should contain everything found in the sample solution

except the substance you are trying to analyze or measure. The reference blank in this case

would be water alone. The amount of light transmitted through a solution is referred to as

transmittance (T). The transmittance is defined as the ratio of the light energy transmitted

through the sample (I) to the energy transmitted through the reference blank (I0). Since the

compound being tested is not present in the reference blank, the transmittance of the reference

blank is defined as 100%T.

T = I/I0

This number is multiplied by 100 to determine the percent transmittance (%T), the percentage

of light transmitted by the substance relative to the reference blank.

%T = (I/I0 )× 100

A certain portion of the light will be absorbed by the compound in the test cuvette; therefore

its %T will be lower than that of the blank (by definition, 100%).

We measure absorbance (Al, also referred to as Optical Density or ODl, where l is the

wavelength used for the measurements), the amount of light absorbed by a solution.

Absorbance is related logarithmically to transmission thusly.

A = -log T

Again, a reference blank is used. In this case, to ‘zero out’ any light absorbed by anything in

the solution other than the compound of interest. By definition, the absorbance of the reference

blank is set at zero (Al = 0). Visible light (see Figure-20) is composed of wavelengths from

400 to 700 nm (nanometers). When visible light passes through a colored solution, some

48

wavelengths are transmitted and others are absorbed. We see the color of the transmitted

wavelengths. For instance, a red color results when a solution absorbs short wavelengths (green

and blue) and transmits longer wavelengths (red). An absorbance spectrum (a plot of

absorbance as a function of wavelength) is determined to select the optimal wavelength for

analyzing a given compound. The optimal wavelength (Amax) for measuring absorbance is

that wavelength that is most absorbed by the compound in question. This provides maximum

sensitivity for your measurements. A hypothetical absorbance spectrum is shown in Figure -21

Figure 17:A typical absorbance spectrum

The light from the spectrophotometer’s light source (in the case of measurements in the visible

range, a simple incandescent bulb) does not consist of a single wavelength, but a continuous

portion of the electromagnetic spectrum. This light is separated into specific portions of the

spectrum through the use of prisms or a diffraction grating. A small portion of the separated

spectrum then passes through a narrow slit. When you adjust the wavelength on a

spectrophotometer, we are changing the position of the prism or diffraction grating so that

different wavelengths of light are directed at the slit. The smaller the slit width, the better the

ability of the instrument to resolve various compounds. This small band of light then passes

through the cuvette containing the sample. Light that passes through the sample is detected by

a photocell and measured to yield the transmittance or absorbance value (optical density) for

the sample. See Figure-22 for a schematic of a spectrophotometer.

49

Figure 18: Components of a Spectrophotometer

Theory

When light (monochromatic or heterogeneous) falls upon a homogenous medium, a portion of

incident light is reflected, a portion is absorbed with in the medium. And the remainder is

transmitted. If the intensity of the incident light is expressed by Io, that of the absorbed light

by Ia, that of the transmitted light by It, and that of the reflected light by Ir, than,

Io = Ia + It + Ir

For air-glass interfaces, arising from the use of glass cell, it may be stated that about 4 percent

of the incident light reflected. It is usually eliminated by the use of a control, such as a

comparison cell, hence

Io = Ia + It

The two separate laws governing absorption are usually known as Lambert’s Law and Beer’s

Law. In the combined form they are referred to as the Beer-Lambert Law, which is the

fundamental equation of spectrophotometry.

Lambert’s Law states that when monochromatic light passes through a transparent medium,

the rate of decrease in intensity with the thickness of the medium is proportional to the intensity

of the light. The law may be expressed by the following equation:

dI/dt = kl

Where I is the intensity of the incident light of wavelength, t is the thickness of the medium,

and k is a proportionality factor. Integrating equation and putting I = Io,

ln (Io /I t) = kt

I t = Ioe -k t

Where Io is the intensity of the incident light falling upon an absorbing medium of thickness

l, it is the intensity of the transmitted light and k is a constant for the wavelength and the

absorbing medium used.

Beer’s Law states that the intensity of a beam of monochromatic light decreases exponentially

as the concentration of the absorbing substance increases arithmetically. This may be written

in the form:

I t = Io e -k c

I t = Io .10 -k c

Where c is the concentration and k and k are constants. Combination of these two laws gives

Beer-Lambert Law:

50

I t = Io 10 -c l

Since each isolated species of ion atom or molecule will exhibit an individual set of definite

energy levels, it will absorb only the frequencies that corresponds to excitation from one level

to another. In this experiment, the maximum wavelength (the wavelength at which maximum

absorbance occurs at first) was measured. Then with the help of maximum wavelength a curve

of absorbance vs. concentration was prepared by taking different concentration of the key

solution. The ion concentration at its definite absorbance at maximum wavelength is obtained

from this graph.

Chemistry Involved

When a transparent solution of Fe+3 ions is subjected to radiating light it remains inert with

respect to the system. So it is required to convert the Fe+3 ions into equivalent amount of

complexes, which are sensitive to spectrophotometry. This may be done by the addition of

ligands like SCN-. In an acidic Fe+3 ion solution, the excess amount of ligand forms ox-red

colored complex equivalent to Fe+3 ions.

Fe+3+ SCN-=[Fe(SCN)]-3

Instrumentation

Spectrophotometer

UV/Vis spectroscopy is routinely used in the quantitative determination of solutions of

transition metal ions and highly conjugated organic compounds. The instrument used in

ultraviolet-visible spectroscopy is called a UV/vis spectrophotometer. It measures the intensity

of light passing through a sample (I), and compares it to the intensity of light before it passes

through the sample (Io). The ratio I / Io is called the transmittance, and is usually expressed as

a percentage (%T). The absorbance, A, is based on the transmittance:

Figure 19: Absorbance vs

Wavelength graph

Figure 20: Absorbance vs

Concentration graph

51

A = − log(%T)

The basic parts of a spectrophotometer are a light source (often an incandescent bulb for the

visible wavelengths, or a deuterium arc lamp in the ultraviolet), a holder for the sample, a

diffraction grating or monochromator to separate the different wavelengths of light, and a

detector. The detector is typically a photodiode or a CCD. Photodiodes are used with

monochromators, which filter the light so that only light of a single wavelength reaches the

detector. Diffraction gratings are used with CCDs, which collects light of different wavelengths

on different pixels.

Figure 21: Diagram of a single-beam UV/vis spectrophotometer.

A spectrophotometer can be either single beam or double beam. In a single beam instrument

(such as the Spectronic 20), all of the light passes through the sample cell. Io must be measured

by removing the sample. This was the earliest design, but is still in common use in both

teaching and industrial labs.

Instrument used in the lab – Bausch and Laumb spectronic-21

A spectrophotometer must consist of some essential components, such as-

Source

Prism

Screen(passing monochromatic light)

Spectrometer

52

Photo meter

Amplifier

Recorder

Figure 22: Working principle of a spectrophotometer

Each system will have proper slits, lenses, controls etc. curettes are required to hold the sample

which must be transparent to the of our interest and must be of fixed size. For a general

purpose instrument we need a way to produce broad range of with reasonable uniform

intensity. The goal of a wavelength selector is to only allow a specific to reach the detector

at any given time .

Source Prism Screen Spectrometer Photomete

r

Amplifier Recorder

53

Method of Operation

Procedure mentioned below is maintained during conducting this experiment-

1. Ferric solution of concentration 1 to 5 ppm are prepared and then 20ml 10% KSCN and

3ml 6(N) HNO3 acid is added in each of the prepared ferric solution.

2. Solution of concentration of 3 ppm is taken as standard solution.

3.

that each wavelength is adjusted to 100% transmittance for solvent i.e. distilled water.

4. Setting the spectrophotometer at the maximum wavelength, absorbance against each

concentration is evaluated.

5. Then absorbance VS concentration curve is constructed.

6. From the plot, which is a straight line, knowing absorbance of the unknown solution,

concentration of that solution is evaluated.

7. In order to evaluated maximum wavelength, absorbance VS wavelength spectrum is

plotted.

Advantages

Important characteristics of photometric method include –

Spectrophotometer can provide narrow bandwidth of radiation and also can handle

absorption spectra in ultra violet region, which turns spectrophotometric method

dominating over spectrometer or photometer.

Spectrophotometric estimation can provide a simple mean of determining minute quantities

of substance. As for example, it can detect 10-4 to 10-6 mol/liter in any solution.

Figure 23: Samples of different concentration Figure 24: Spectrophotometer in the lab

54

So as a whole it has- (a). wide applicability; (b). high sensitivity; (c). moderate to high

selectivity; (d). accuracy; & (e). convenience.

Disadvantages

For solutions that do not follow Beer-Lambert’s law it cannot be used.

For very solutions (transmittance lie outside the range 0.20 to 0.65)Beer–Lambert’s law

shows deviation. So problem arises in this range.

Results

From Calibration Curve:

If absorbance of a solution is = 0.505

Then Concentration of the solution = 3.52 ppm

Discussion

The prediction of Beer Lambert’s law is verified in this experiment. By plotting absorbance Vs

concentration a straight line should be obtained which provides a linear relation between

absorbance and concentration.

The calibration curve of absorbance versus concentration (at constant wavelength) should pass

through the origin since at the zero concentration (Infinite dilution) absorbance must be zero.

But in this experiment, least square fitting of the curve yields a straight line while passed

straightly above the origin. The following reasons may be responsible for this –

In complete conversion of ferrous ion (Fe 2+) to ferric ion (Fe 3+) may occurred;

Impurities present in the sample solution (solution of 3ppm concentration) which may

absorb light at the absorption wavelength; &

Impurities present in the distilled water used in preparing sample solution.

In the spectrum of Fe [(SCN)6]-3 the highest peak indicates strong (maximum absorption) while

the small peak indicates weak absorption. Maximum absorption occurs at a particular

wavelength but not at any wavelength. The sample compound (Fe3+ solution) doesn’t absorb

radiation appreciably in the provided wavelength regions. It is necessary to form an absorbing

substance by reacting the sample with other reagents. That is why KSCN is added to Fe3+

solution. This forms a complex with Fe3+ solution. A ligand (SCN-) had been added to form

color of the solution, as for colorless solutions no absorption will take place in the

55

spectrophotometer. again the Fe3+ ion concentration is so minute that it couldn’t form any color,

that’s why the ligand (NH4SCN) had to be added.

The plot drown in Spectrophotometric estimation can be used for determining concentrations

of other solutions of the same species whereas in the other methods a plot can be used only for

that respective solution.

56

ELECTROGRAVIMETRIC ANALYSIS

Sample: Brass Sample Containing Certain Amount Of Copper And Zinc

Introduction

Electrogravimetry is a special case of gravimetry where electrolysis is carried out and the

material deposited on one of the electrodes is weighed. In most applications the metal is

deposited on a weighed platinum cathode. This sort of electrolytic precipitation has been used

for over a century for the gravimetric determination of metals.

This topic will describe in brief the theory, chemistry of the process, experimental procedure,

instrumentation, advantage and limitation, result & discussion of electro gravimetric method

related to the experiment performed in the lab.

THEORY

When a direct electric current is passed through an ELECTROLYTE (such as a molten salt or

an aqueous solution of a salt, acid or base), chemical reactions take place at the contacts

between the circuit and the solution. This process is called ELECTROLYSIS.

Figure 25: Basic construction of an electrolytic cell

In Electrogravimetric analysis, the element is to be determined is deposited electrolytically

upon a suitable electrode. Electro deposition is governed by Ohm’s law and by Faraday’s law

of electrolysis.

Ohm’s Law: It expresses the relation between the three fundamental quantities: current,

electromotive force and resistance.

57

The current I is directly proportional to the resistance R, i.e.

I = E/R

Faraday’s Law:

1. The amounts of substances liberated or dissolved at the electrodes of a cell are directly

proportional to the quantity of electricity which passes through the solution.

2. The amounts of different substances liberated or dissolved by the same quantity of

electricity are proportional to their relative atomic (or molar) masses divided the number

of electrons involved in the respective electrode processes.

In the common method of Electrogravimetric analysis, a potential slightly, in excess of the

decomposition potential of the electrolyte under investigation is applied, and the electrolysis is

allowed to proceed without further attention, except perhaps occasionally to increase the

applied potential to keep the current at approximately the same value. For this purpose an

electronic device called potentiostat is used in the circuit. This device controls the potential

difference between the cathode and the reference electrode.

Chemistry Involved

The chemistry involved in this experiment was purely the chemistry of a general electrolysis.

The positive ions of the electrolyte, in this case Cu+ and H+ collected electrons from the

cathode whereas negative ions of the electrolyte, OH- in this experiment, delivered electrons to

anode.

Copper may be deposited from either sulphuric acid or nitric acid solution but here a mixture

of the two acids is employed. When this solution is electrolyzed the following reactions

occurred –

Reaction at Cathode: CueCu 22

Reaction at Anode: eOHOOH 424 22

When the sample was dissolved in conc.HNO3 solution the following reactions took place-

Cu + 4HNO3 = Cu++ + 2NO3- + 2NO2 + 2H2O

Zn + 4HNO3 = Zn++ + 2NO3- + 2NO2 + 2H2O

The beneficial effect of nitrate ion is due to its depolarizing action at the cathode. The reaction

is

58

OHNHeHNO 243 3810

The reduction potential of NO3- is lower than the discharge potential of hydrogen. Here the

formation of ammonium ion occurs to the exclusion of gas evolution.

The beneficial effect of sulphate ion (SO42-) is to the easy liberation of NO3

- for depolarization

and also for liberation of nitrate ion (NO3-) as NO2 when it is warmed. The reaction is –

eONONO 221

23

Instrumentation

Instrument used in the lab- an Electrochemical cell containing platinum electrodes

An electricity source (0.4-0.5 A, 5V)

magnetic stirrer

The essential requirements for a constant current electrolytic determination – a source of

(direct) current, a variable resistance, an ammeter, a voltmeter and a pair of platinum electrodes

were assembled in the laboratory. A platinum metal electrode that is covered with a rough,

large surface area platinum coating. The purpose is to produce an electrode with a large true

area that will be relatively non polarizable

Figure 26: Schematic Diagram of Electrogravimetric analysis

59

In order to carry out the electrolysis properly magnetic stirrer was used for constant stirring of

the solution. Two platinum electrodes were used for electrolysis. A constant electric supply

was maintained by changing the potential at necessary moments.

Figure 27: Instrument used in laboratory for electrolysis

Figure 28: Copper deposition on the platinum electrode

Procedure

Procedure mentioned below was followed during conducting this experiment –

By using an electronic balance weight of brass and platinum gauge cathode were taken;

Brass was then taken into a beaker and concentrated HNO3 acid is added drop wise to

dissolve the brass. The color of the solution at that instant was blue;

1 ml of concentrated H2SO4 acid was then added into the beaker;

A small amount of distilled water was poured into the above said beaker and it is heated

until the solution becomes concentrated and white fumes are observed;

The stirrer was switched on and the electrolysis is continued until the blue color the solution

has entirely disappeared (about 1 hr);

The cathode was sunk into C2H5OH and then kept in open air for quick drying;

After drying, the weight of (cathode + Cu) was taken; &

From the increase in weight of the cathode, the copper content in brass was calculated.

60

Advantages

Applicability of this method in electroplating i.e. coating of metal substances by depositing

on it Ni, Cu, Au, Ag, Pt etc. These non-corrosive metals make metallic substances long

lasting and lustrous.

An interesting application of this method is the direct quantitative separation of constituents

from alloy. The constituent with low deposition potential will be deposited first.

In this method, with rising and lowering the temperature of the solution, physical properties

of the deposition can be controlled.

By applying this process, constituents and also percentage composition of the constituents

in any alloy can easily be evaluated. Thus purity of a substance can be measured.

Limitations

Some limitation of this electro gravimetric method is also present although this process is very

much advantageous. They are listed below –

For gravimetric purpose, an electrolytic deposit should be strongly adherent, dense and

smooth so that the process of washing, drying and weighing can be performed without

mechanical loss or without reaction with the atmosphere.

Fluctuating current density may result in flaky, spongy, power granular, uneven deposits,

which only loosely bound to the electrode.

At sufficiently high values of the current density evolution of hydrogen may occur owing

to the depletion of metal ions near the cathode. If appreciable evolution of hydrogen gas

occurs, the deposit will usually become broken up and irregular which may lead to spongy

and poorly adherent deposits.

If cost of electricity is high then use of this process for coating makes substances costly.

Many metals form smoother and more adherent films when deposited from solutions in

which their ions exist primarily as complexes. Deposition of these metals requires a higher

applied potential than in the absence of the legend. These effects can be large and must be

taken into account when considering the feasibility of an electrolytic determination or

separation.

61

Result

The amount of Cu determined in the brass sample was 76.13%

Discussion

From the result it is clear that complete deposition of Cu was not achieved. This may happened

due to the following reasons:

The brass used here was impure

In preparing Cu solution heating is required to remove excess NO3- as NO2 gas, but in this

case excess heat was given which oxidized Cu into CuO,

Cu + ½ O2 = CuO

Since Nitrite ion resists complete deposition it should be made the HNO3 acid free from

HNO2 either –

- by addition of urea to the solution,

2H+ + 2NO2- + NH2CONH2 = 2N2 + CO2 + 3H2O

- by boiling the HNO3 before adding it; or

- by adding sulphamic acid to the solution,

H+ + NO2- + SO3

-NH2 = N2 + HSO4- + H2O

If any of the above said measures are not performed, complete deposition is quite impossible.

In this experiment the dilute solution is prepared otherwise incomplete deposition of Cu

might take place or the deposition will not adhere satisfactorily to the cathode.

Polarization occurs due to evolution of H2 gas to cathode and it leads to deposition of

spongy and poor quality copper. That’s why the deposition of Cu is carried out in HNO3

media because nitrate ion reduced to ammonium ion at a copper surface,

NO3- + 10H- + 8e- = NH4

+ + 3H2O

During electroplating, the concentration of Cu ion at the electrode is depleted relative

to the rest of the solution during the passage of any appreciable current. This is termed as

concentration over potential. A magnetic stirrer is used to keep this term as small as possible.

In this experiment, low constant current density is used for better deposition.

62

To find out all the Cu deposited or not some water had been added, the cathode is more

dipped and it was tested that whether it becomes reddish or not, when no more Cu deposited

on the cathode then the power was turned off.

The voltage and current supply had been maintained constant throughout the experiment

Caution

Cathode should be removed from the solution at the time when it starts changing its

color from salmon pink to black due to oxide formation (CuO).

Cathode & anode must be removed from the solution before the switch is turned off;

otherwise there is possibility of back reaction.

63

Conclusion There was a time when people used to rub stones to light fire. But those days are history now.

No one will even think about lighting fire from stones now a days. With the similar manner the

classical methods of analytical chemistry will also become obsolete in near future. But the

importance of instrumental analysis is increasing day by day. It should be mentioned here that

instrumental methods of analysis is a must for an engineer who has to deal with instruments all

the times. Whereas a Chemical Engineer has to frequently deal with the problems like

separation of different components of a sample determination of the amount of an element, find

out end point of solution etc. and so on.

To deal with these types of problems, a chemical Engineer should have a good knowledge about

the instrumental methods of analysis, for example, spectrophotometer can be used for the

presence and amount of an element determination in a solution.

Since the modern world has been heading towards instrumentation, where every unit of

parameter comes out directly in digits, where control system is remote control, so a Chemical

Engineer should know the use of different types of instruments, their working methods,

principles and even their limitations.

The wide utility of analytical chemistry has sufficient reasons alone for treating it is an

individual course. The instrumental methods of analytical chemistry of this course (Chem 352)

help the students to learn careful and quantitative laboratory skills and techniques. For the

improvement of this course (Chem352) and to offer the students translucent conception about

instrumental methods of analysis the following measures should be taken under consideration

–

Number of apparatus and instrument in the laboratory should be increased so that every

student can perform each of the experiment individually.

A pen recorder should be provided as when it is connected to a cell during performing

conductometric titration, draws a plot of change in conductance with addition of reagent

used automatically. Drawing plots manually is time consuming and sometime gives

erroneous result.

64

Bibliography

Bahl, B. S., Tuli, G.D., Bahl, A., “Essentials of Physical Chemistry”, Millennium edition,

S.Chand Publications, (2000).

Besset & associates, “Vogel’s Textbook of Quantitative Inorganic Analysis”, 5th

ed.,Addison Wesely Longman Inc,(1999).

Glasstone, S., “An Introduction to Electrochemistry”, International edition, East-West

Press Private Ltd, (1942).

Maron, S. H., Prutton, C. F., “Principles of Physical Chemistry”, 4th edition, McMillan

Publishing Co.,(1965).

Williard, H. H., Merritt, L.L., & Dean, J.A., “Instrumental method of analysis”, 4th ed.,

Nostrand company, London.

65

Appendices

Containing

Observed data

Calculated data

Sample calculation

graphs

66

Appendix-a

POTENTIOMETRIC TITRATION of iron (Fe2+)

Concentration of K2Cr2O7 , S1= 0.1(N)

Concentration of H2SO4=6.0(N)

Volume of ferrous (Fe++) solution, V2=10 ml

Required volume of Kr2Cr2O7 , V1=4.45ml

Table 1: Observed and calculated data for volume, cell EMF, first and second derivative

of cell EMF

Volume of

K2Cr2O7

(ml)

E (Volt)

dV

dE

dE/dV

d2E/dV2

0 0.35125 0.2 0.00625 0.03125 -0.0625

0.2 0.3575 0.2 0.00875 0.04375 0.03125

0.4 0.36625 0.2 0.0075 0.0375 -0.03125

0.6 0.37375 0.2 0.00875 0.04375 0.0625

0.8 0.3825 0.2 0.00625 0.03125 0.03125

1 0.38875 0.2 0.005 0.025 0.03125

1.2 0.39375 0.2 0.00375 0.01875 -1.3E-15

1.4 0.3975 0.2 0.00375 0.01875 -0.0625

1.6 0.40125 0.2 0.00625 0.03125 0.09375

1.8 0.4075 0.2 0.0025 0.0125 -0.03125

2 0.41 0.2 0.00375 0.01875 -0.03125

2.2 0.41375 0.2 0.005 0.025 0.03125

2.4 0.41875 0.2 0.00375 0.01875 -0.0625

2.6 0.4225 0.2 0.00625 0.03125 0.0625

2.8 0.42875 0.2 0.00375 0.01875 -0.09375

3 0.4325 0.2 0.0075 0.0375 0.0625

3.2 0.44 0.2 0.005 0.025 -0.0625

3.4 0.445 0.2 0.0075 0.0375 2.36E-15

3.6 0.4525 0.2 0.0075 0.0375 -0.125

3.8 0.46 0.1 0.005 0.05 -0.25

3.9 0.465 0.1 0.0075 0.075 -3.3E-15

4 0.4725 0.1 0.0075 0.075 0.125

4.1 0.48 0.1 0.00625 0.0625 -0.375

4.2 0.48625 0.1 0.01 0.1 8.88E-15

4.3 0.49625 0.1 0.01 0.1 -13.5

4.4 0.50625 0.1 0.145 1.45 12.125

4.5 0.65125 0.1 0.02375 0.2375 0.9375

4.6 0.675 0.2 0.01 0.05 0.0625

4.8 0.685 0.2 0.0075 0.0375 0.0625

5 0.6925 0.2 0.005 0.025 0.0625

5.2 0.6975 0.2 0.0025 0.0125 -0.03125

5.4 0.7 0.2 0.00375 0.01875 4.16E-16

5.6 0.70375 0.2 0.00375 0.01875 2.78E-15

5.8 0.7075 0.2 0.00375 0.01875 0.016632

6 0.71125 0.118542 -0.0625

67

Sample Calculation

Strength of Fe2+ at the end point

Concentration of K2Cr2O7 ,, S1 = 0.1N

From graph required volume of K2Cr2O7 for titration,V1 = 4.45 ml

Volume of ferrous (Fe++) solution, V2 =20 ml

Concentration of Fe2+ solution, S2 = ?

Now,

V1S1 = V2S2

S2 = (4.450.1) / 20

= 0.02225N

Determination of Concentration of Fe3+ ion For 4 different points of the Graph

For V=3ml K2Cr2O7 ,Ecell=0.4325 V

Now,Ecell=EFe2+/Fe

3+ +Ec

=> EFe2+/Fe

3+ = Ecell - Ec……………….(Eq-1)

Here,Ec =0.244 V for saturated KCl solution

So,from Eq-1, EFe2+/Fe

3+ =0.1885 V

Again, EFe2+/Fe

3+ = E0Fe

2+/Fe3+ -

𝑅𝑇

𝐹 ln

𝑎𝐹𝑒3+

𝑎𝐹𝑒2+ ……..(Eq-2)

Here,

E0Fe

2+/Fe3+ = - 0.771 V

R=8.314 J/(mol.K)

T=320C =305 K (Room temperature)

F=96500 Faraday

From Eq-2, 𝑎

𝐹𝑒3+

𝑎𝐹𝑒2+ = 1.387×10-16

Strength of Fe2+ at the point where 3ml K2Cr2O7 is used can be calculated from equivalence

relation.

4.45 ml 0.1N K2Cr2O7 = 0.02225 N 20 ml Fe2+

1ml 0.1 N K2Cr2O7 = 0.005N 20 ml Fe2+

3 ml 0.1 N K2Cr2O7 = 0.015 N 20 ml Fe2

So, aFe2+ =0.015 N

Now, aFe3+ =1.387×10-16 × aFe2+

=2.08×10-18 N

71

APPENDIX-B

CONDUCTOMETRIC TITRATION of an acid mixture(Hcl+CH3COOH)

Observed and calculated data:

Volume of the NaOH = 10ml

Strength of Oxalic Acid = 0.1N

Table 2: Table for volume of NaOH (ml) & conductance (ms)

Volume of NaOH

(ml)

Conductance

(mS)

Volume of NaOH

(ml)

Conductance

(mS)

0.0 2.30 5.4 1.22

0.3 2.24 5.7 1.25

0.6 2.07 6.0 1.28

0.9 1.89 6.3 1.32

1.2 1.76 6.6 1.35

1.5 1.59 6.9 1.39

1.8 1.48 7.2 1.42

2.1 1.35 7.5 1.46

2.4 1.22 7.8 1.50

2.7 1.11 8.1 1.57

3.0 1.04 8.4 1.65

3.3 1.02 8.7 1.74

3.6 1.03 9.0 1.83

3.9 1.05 9.3 1.94

4.2 1.08 9.6 2.06

4.5 1.11 9.9 2.15

4.8 1.14 10.2 2.24

5.1 1.18 10.5 2.33

Table 3: Data for Determination of the Strength of Oxalic Acid

Observation

No.

Volume of

NaOH(ml)

Burette Reading

Volume of

Oxalic Acid(ml)

IBR(ml)

FBR(ml)

01 10 0 20 20

02 10 20.2 40.2 20

Sample Calculation

Standardization of NaOH by H2C2O4

Here, volume of NaOH, V1 = 10 ml

Volume of H2C2O4 , V2 = 20 ml

72

Strength of H2C2O4 , S2 = 0.1N

Strength of NaOH = S1

Now, V1 × S1 = V2 × S2

S1 = (20 × 0.1)/ 10

= 0.2N

So, the strength of NaOH = 0.2N

Calculation for HCl

Here, the volume of NaOH for neutralizing HCl, V1 = 9.5ml (from graph)

Strength of NaOH, S1 = 0.2N

Volume of HCl solution, V2 = 10ml

Strength of HCl = S2

Again, V1 × S1 = V2 × S2

S2 = (9.5 × 0.2)/ 10

= 0.19N

Now, Equivalent weight of HCl, e = 36.5

Weight of HCl = W1

Again, W1 = S2 e V2

= (0.19 × 36.5 × 10)/ 1000

= 0.06935g

Calculation of CH3COOH

Here, volume of NaOH for neutralizing CH3COOH, V1 = (20.1-9.5) = 10.6ml

Strength of NaOH, S1 = 0.2N

Volume of CH3COOH solution, V2 = 10ml

Strength of CH3COOH = S2

Again, V1 × S1 = V2 × S2

S2 = (10.6 × 0.2)/ 10

= 0.212N

Now, Equivalent weight of CH3COOH, e = 60

Weight of CH3COOH = W2

Again, W2 = S2 e V2

= (0.212 × 60 × 10)/ 1000

= 0.1272g

Percentage (%) of HCl = [W1/ (W1+ W2)] × 100%

73

= [0.06935/ (0.1272 + 0.06935)] × 100%

= 35.28%

Percentage (%) of CH3COOH = [W2/ (W1+ W2)] × 100%

= [0.1272/ (0.1272 + 0.06935)] × 100%

= 64.72%

75

APPENDIX-C

pH METRIC TITRATION

Observed and calculated data

Table 4: Observed and Calculated data for pH-matric titration

Sample Calculation

Concentration of Na2CO3, S1 = 0.05 N

From graph,

Required volume of HCl acid for complete titration, V2= 2.78 ml (from graph)

Volume of ferrous Na2CO3 solution, V1= 10 ml

Now,

V1S1 = V2S2

S2 = 𝟏𝟎×.𝟎𝟓

𝟐.𝟖= 0.1798 N

Calculations for dissociation constants

Volume of HCl for converting Na2CO3 to NaHCO3, V1 = 1.17 ml

Dissociation constant of Na2CO3, pk1 = the corresponding pH for V1/2 ml of HCl

= 7.7 (From Graph)

Volume(V) pH dV d(pH) d(pH)/dV d2(pH)/dV2

0 9.93

0.8 9.13 0.8 0.8 1 6

0.9 8.97 0.1 0.16 1.6 -3

1 8.84 0.1 0.13 1.3 41

1.1 8.3 0.1 0.54 5.4 13

1.2 7.63 0.1 0.67 6.7 -20

1.3 7.16 0.1 0.47 4.7 -8

1.4 6.77 0.1 0.39 3.9 -5.08333

2 6.12 0.6 0.65 1.083333 1

2.2 6.05 0.2 0.07 0.35 3.5

2.4 5.68 0.2 0.37 1.85 -0.72222

2.7 5.21 0.3 0.47 1.566667 59.66667

2.8 4.45 0.1 0.76 7.6 -29

2.9 3.98 0.1 0.47 4.7 -16

3 3.75 0.1 0.23 2.3 -4.84

3.5 3.38 0.5 0.37 0.74 -0.92

4 3.22 0.5 0.16 0.32 0.56

4.5 3.08 0.5 0.14 0.28

76

Volume of HCl for neutralizing NaHCO3, V2 = 2.78-1.17=1.61 ml

Dissociation constant of NaHCO3, pk2 = the corresponding pH for (V1 + V2 /2) ml of HCl

= 5.25(From Graph)

81

APPENDIX-D

Spectrophotometric Determination of Iron

Observed Data

Peak Pick

Wavelength, λmax = 478.5nm

Absorbance = 0.4187

Table 5: Table for observed data of absorbance of different concentration

Calculated Data

From the following Graph at Absorbance = 0.505

Concentration of the known solution = 3.52ppm

And also from the graph, the equation of the line is,

y = 0.1427x

x = Concentration & y = Absorbance

At y = 0.505, From the Equation x =3.52

No. of observation Concentration (ppm) Absorbance

01 Reference cell 0.0588

02 1 0.1874

03 2 0.320

04 3 0.440

05 4 0.553

06 5 0.698

unknown - 0.503

83

APPENDIX-E

Electrogravimetric Determination Of Copper In A Brass Sample

Observed Data

Initial weight of the platinum electrode = 18.3889gm

Weight of the electrode after deposition = 18.5135gm

Weight of brass = 0.2950gm

Sample Calculation

Weight of the copper deposited = (18.3889 – 18.5135)gm

= 0.2246gm

Percentage recovery = (0.2246/0.2950) × 100%

= 76.13%