CHEM-352
A term paper on INSTRUMENTAL METHODS OF ANALYSIS
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, Depar tment Of Chemistry, BUET
Date of Submission: 11th June, 2013
Partners Id: 0902038 0902040 0902044
Bangladesh University of Engineering and Technology i
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 experiments were performed in the laboratory. This term paper 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 chemistr y.
ii
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
” ” . I am thankful to my I nstrum ental M ethods of An alysis alysis
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
iii
Table of contents Topics
Preface
Page no
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
iv
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 Conductometer used in Laboratory; Laboratory; Model (CMD=750) (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
50
graph
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 K 2Cr 2O7
68 v
Graph-A.2:
dE
vs. Volume of K 2Cr 2O7
dV 2
Graph-A.3:
d E dV
vs. Volume of K 2Cr 2O7(ml)
2
69 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:
pH V 2
Graph-C.3:
vs. Volume Volume of HCl graph
pH
V
2
vs. Volume of HCl graph
79 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 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 sci ence 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 create d this predominance. Professionals in all sciences base important decisions, deci sions, 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, 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 finall y 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 instr uments 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 creat ed a new area in analytical chemistry, so knowledge for operations and handling of those instr uments 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 per formed in the laboratory. Those are as follows:
Potentiometric titration Conductometric titration
7
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 analysis by Justus von Liebig and systematized organic analysis based on the specific reactions of functional groups. The first instrumental analysis was flame emissi ve spectrometry developed by Robert Bunsen and Gustav Kirchhoff who 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 trans formed into high performance instruments. In the 1970s many of these techniques began to be used together to achieve a complete characterization of samples. Starting Starti ng 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
8
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 prett y dull field. pH meters and single si ngle wavelength spectrophotometers, and electrochemical techniques were used. Lots of titrations, gravimetric anal ysis. 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 a nd 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). Experi Experi ment ment
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 t ime 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 scie nce. 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 s tudy 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. 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 materia l with-
Organic Inorganic Both of these can be subdivided into two groups:
Qualitative analysis. Quantitative analysis
Analytical Analytical Chemistry
Quantitative Analysis
Qualitative Analysis
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Quali tative analysis: 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. Quanti tative analysis: 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
Qualitative Analysis
Volumetric Analysis
Gravimetric Analysis
Physico Chemical 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 M ethod
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. Gravimetri Gravimetri c M ethod
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 Physico-Chemical M ethod
This method involves the final measurements of some physical properties of the composition of the system as a whole.
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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 A nalytical Tar gets gets
Bioanalytical chemistry Material analysis Chemical analysis Environmental analysis Forensics By An alytical alytical M ethods thods
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. titr ations. These titrations may use different types t ypes 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. I norganic quali tative analys analysis is
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 measure ment 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 belowM olecular olecular analysis analysis
I)
Nuclear magnetic resonance ii)
iii)
UV absorption and UV fluorescence
Infrared spectroscopy
Spe Spectr ophotometr ophotometr i c An alysis
I)
X- ray diffraction and absorption 14
ii)
Radio tracer techniques
iii) Mass spectrometry
Th er mal analysis analysis
I)
Gas chromatography
ii) Electromicroscopy
El emental mental A nalysis nalysis
I)
Emission spectrograph
ii) Flame photometry
iii) Atomic absorption spectroscopy. Th e Natur e of of M easur asur ement ment
The steps involved in a chemical determination are the following: Gene Generr ation of sign al
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 li ne on a chart, pointer deflection on a meter, met er, 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. Choi ce Of M ethod Of M easur asur ement ment
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 15
Different instrumental methods areSpectroscopy
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. M ass ass Spectr Spectr ometry ometr y
Mass spectrometry measures mass-to-charge ratio of molecules using electric and magnetic fields. There are several ionization methods: electron impact, chemical ionizati on, electrospray, fast atom bombardment, matrix assisted laser desorption ionization, and others. Also, mass spectrometry is categorized by approaches of mass analyzers: magnetic-sector ,quadrupole ,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 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. El ectrochemi ctrochemical cal 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 Cur ve
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 16
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. 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 17
single molecules. This cutting-edge technology has a promising potential of leading a new revolution in science as integrated circuits circuit s 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 anal ysis, 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 tages
1. Sampling
Ex amples of proce pr ocedur dur e
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
Standardization, calibration, optimization,
instrumental factors
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 anal yses may be classified into four types as followsProxi mate analysi analysi s determines the amount of each element in a sample but not the
compounds that are present. Par Par tial analysis analysis determines selected constituents in the sample. Tr ace consti consti tuent analysis is a type of partial analysis that determines specified
components present in very minute quantity. quantit y. Complete anal anal ysi ysi s determines the proportion of each component in the sample.
On the basi , analytical methods are often classified as follows: basi s of sampl e si si ze 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 – (100 – 10000 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 instr uments 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 te chnique 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 i n 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.
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Experiments performed
POTENTIOMETRIC TITRATION
S T N E M I R E P X E
CONDUCTOMETRIC TITRATION pH METRIC METRIC TITRATION SPECTROPHOT SPECTROPHOTOMETRIC OMETRIC METHOD ELECTROGRAVIMETRIC METHOD
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Potentiometric titration 2+ SAMPLE:Fe SOLUTI ON TI TI TRATED TRATED WITH K C O O 2 2Cr r 2 7 7
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 refer ence 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 ele ctrodes which in turn depends on the activity activit y 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”. 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 n+ E = E + --------- ln a M M n F 0
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 number
Potentiometric titration based upon the quantitative relationship between the emf of a cell is given by the distribution of potential E c e l l = E r ef + E i n d i c a t o r + E j u n c t i o n 22
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):
H g 2 e- = 2 H g + 2 C l 2 C l 2 + 2 e+ + - + + + F e - e = F e , E 0 = 0.76v
Platinum (right):
C r 2 O 7 -
-
+ 14H + + 6e - = 2 C r + + + 7H 2 O
Emf of the cell is given by, E c e l l = E l e f t - E r i g 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 i s 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 = E 0 + 0 . 05 0 5 91 9 1 l o g [ F e + + ] / [ F e ]
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. 23
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 r eagent 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 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. 24
From the stock solution 10ml sample is taken in a cleaned beaker and 20 ml of 6(N) N 2SO4 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 K 2Cr 2O7 solution. Then the potential is recorded after addition of each 0.1 ml of the titrant up to 13 ml of K 2Cr 2O7 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 25 0c of 0.3358 volt, 0.2824 volt and 0.2444 volt respectively. Calomel electrode represented as Hg|Hg 2Cl2(S)/C1(C). The electrode reaction is, - H g Cl Cl 2 (S) (S) +2e = 2Hg (li quid) + 2Cl 2Cl - 2 2 2
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 t hat 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 i ndicators 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.
25
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 K 2Cr 2O7 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 K 2Cr 2O7 this ratio changes and EMF changes gradually. The critical problem in a titration tit ration 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 K 2Cr 2O7 in accurately measured increments a sigmoid curve is obtained which is similar to the theoretical titration ti tration curve for oxidation - reduction reaction. It is seen that EMF 26
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. 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 s harp 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 Δ 2E / ΔV2) becomes zero. All of the methods considered are based on the assumption that the titration ti tration 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.
27
Conductometric titration SAMPLE: H Cl ACID & CH & CH 3 COOH COOH 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 NaNO 3), 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 s olution 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 + BThus 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.
28
Chemistry of str ong acids & weak weak acids with str ong base bases
End point
Conductance
Volume of base Figure 4: Conductometric Conductomet ric titration curve for strong acid and strong base.
Reaction:
- + - + - (H + + Cl ) + (N a + OH ) = (N a + Cl ) + H 2 O O
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:
+ - + CH 3 COOH C OOH + N a + OH - = CH 3 COO COO + N a + H 2 O O
End point Conductance
Volume of base Figure 5: Conductometric Conductomet ric titration curve for weak acid and strong base. Immediately after the end point, further addition of NaOH introduces the f ast moving hydroxyl ions and the conductivity value shows a sharp increase. The point of intersection of the two curves therefore gives the end point. 29
Consider the titration of solution of weak acid, such as aceti c 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 c onductivity of a titrated solution is observed until the equivalence point is reached, at which we have a solution of sodium acetate, CH 3COONa. 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 30
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 titrated using solution of sodium hydroxide.
acetic acid mixture
–
As is seen, the Conductometric titration curve is i s 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.
31
CHEMICAL REACTIONS
Sodium Hydroxide, Acetic acid and Sulfuric acid ionize in the solution in the following forms: + NaOH NaOH = Na + OH -
HCl = H + + Cl - - + CH 3 COOH C OOH = CH 3 C 3COO OO + H
Then the following reactions occur: + N a + Cl - = NaCl
H + + OH - = H 2 O O NaOH NaOH + H Cl = NaCl NaCl + H 2 O O NaOH + CH 3 COOH C OOH = CH 3 COONa C OONa + H 2 O O 3
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
32
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 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 K can be estimated by the following equation: K
d
A
where d is d is the distance between the sensing elements and A 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. polarization. Using AC potential potential eliminates eliminates the latter problem, problem, but an AC Wheatstone Wheatstone bridge bridge minimi minimizes zes both of the above above problems. problems. In Figure Figure 13, R f1 f1 and R f2 f2 are fixed resistors of known resistance; R b is a resistance box in i n which the resistance can be b e changed very ver y precisely precis ely over a wide range; and R m is the resistance to be measured, in this case the conductivity cell with the solution of interes t.
Figure 9: Schematic diagram of Wheatstone Wheatstone bridge 33
The bridge works correctly when V 2 = V4 and therefore R m = R b(R f1 f1/R f2 f2). If the two fixed resistors are chosen to be identical, then R m is equal to the reading of the variable resistance box R b. (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
34
Method of Operation
20 ml mixture of HCl and acetic acid was taken ta ken 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. taken . A graph of specific conductance vs Volume added of NaOH was plotted From the graph the end point strength 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 forVery 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 mi xtures of HCl or any other strong s trong 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 h ydrolysis at the equivalent point very week
35
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 excessi ve 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 le vel of contamination of water or any an y other chemical species, as the conductance of pure species of diff erent substances are constant.
Result Strength of `HCl= 0.19 N Strength of CH 3COOH = 0.212 N Percentage of HCl= 35.28% Percentage of CH 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 neutralizat ion 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
36
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 CH 3COONa 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 the conducting power of highly highly ionized salt exceeds that of the the weak acid. Immediately 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 – consideration – The relative change of conductance of the solution during the re action 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. accurac y. 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 te mperature 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 e xtremely 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. 37
P
H metric titration
Sample: mple: H Cl A cid Titr ated ated With N aOH
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 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 h ydrogen 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=-log ] 10 10 [ H With the introduction of the concept of activity the correct definition is
pH= -l og a H H
Where, aH = hydrogen ion activity In acid base titration pH of the test solution changes with ever y 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 cell s 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 st em of non-pH-responsive, high resistance glass. The potential of the glass electrode is given by, E = E e .000198T T (pH -pH ) e + .000198 r
Where, pHr = pH of the reference solution contained within the glass bulb Ee =combination of junction potential T =Temperature in degree Kelvin
38
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 electr ode that is dipped in the solution. The conventional pH value may be defined in an operational manner, as follows, pH = pH s + (E-E ).000198T. s s
Where, E= emf of pH cell which cantains an unknown solution. E s= emf of pH cell which contains standard referance of known pH. 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 Na 2CO3 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 Na 2CO3, 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-
H A
K
HA
A log H log
log K
HA
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 + H 2CO3
CO2 + H 2O
39
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 end point d2 pH/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 sensor made from glass of a specific composition. It is an Ionselective electrode (ISE) with a transducer (sensor) 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 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. time basis. 40
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 or calomel electrode
41
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 electr ode 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 42
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 NaCO 3 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 Na 2CO3 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. 43
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, pK 1= 7.7 Second dissociation constant, pK 2= 5.25
Discussion The purpose of this experiment was to study the nature of Na 2CO3-HCl titration with the measurement of pH and to plot the titration ti tration curve of pH vs. milliliters milliliter s 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 Na 2CO3 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: N a 2 CO 3 + HCl = NaHCO 3 + 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 Na 2CO3. 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. NaHCO NaCl + H 2 O + CO 3 3 + H Cl = NaCl 2 2 2
This point indicates the conversion of HCO 3- to H 2CO3 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 Δ 2 pH/ΔV2 Vs V is plotted. 44
This point indicates the conversion of HCO 3- to H 2CO3 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:
salt acid
pH=pK+ log
So when 50% of the acid is neutralized and forms salt then pH= pK, so from the curve the pK 1 and pK 2 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 act ual 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 body can be measured through it.
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SPECTROPHOTOMETRIC ANALYSIS S a m p l e : F e3 + s o l u t i o n
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. spectr ophotometer. Light of a particular wavelength passes through the solution inside the cuvette and the amount of light transmitted (passed through the solution — Transmittance) 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. 46
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/I 0 0
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/I 0 ) ×100 10 0
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 = -l og 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 re ference 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 47
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 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.
48
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 that of the reflected light by Ir, than, I o = I a + I t + I r For air-glass interfaces, arising a rising 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 I o = I a + I t 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 (I / ) o I t = k t - k t I t = I o e
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: - k c I t = I o e k c I t = I o . 1 0 - k c
Where c is the concentration and k and k are constants. Combination of these two laws gives Beer-Lambert Law: 49
I t = I o 1 0 - 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.
Figure 19: Absorbance vs Wavelength graph
Figure 20: Absorbance vs Concentration 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. + 3 + SCN - = [ F e ( S C N ) ] - 3 F e
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 intensit y of light passing through a sample ( I ( I ), ), and compares it to the intensity of light before it passes through the sample ( I o). The ratio I ratio I // I o is called the transmittance, transmittance, and is usually expressed as a percentage (%T). The absorbance, A absorbance, A,, is based on the transmittance:
50
A
=
−
l og (%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 single beam or double or double beam. beam. In a single beam instrument (such as the Spectronic 20), 20), all of the light passes through the sample cell. I cell. I o 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 – lab – Bausch Bausch and Laumb spectronic-21 A spectrophotometer must consist of some essential components, such asSource Prism Screen(passing monochromatic light) Spectrometer 51
Photo meter Amplifier Recorder
Source
Prism
Screen
Spectrometer
Recorder
Photomete
Amplifier
Figure 22: Working principle of a spectrophotometer spectrophotometer
Each system will have proper slits, lenses, controls etc. curettes are r equired 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 .
52
Method of Operation Procedure mentioned below is maintained during conducting this experiment1. Ferric solution of concentration 1 to 5 ppm are prepared and then 20ml 10% KSCN and 3ml 6(N) HNO 3 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.
Figure 23: Samples of different concentration
Figure 24: Spectrophotometer in the lab
Advantages Important characteristics of photometric method include – 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.
53
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 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 la w 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 distille d 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 (Fe 3+ solution) doesn’t absorb radiation appreciably in the provided wavelength regions. It is necessar y to form an absorbing substance by reacting the sample with other reagents. That is why KSCN is added to Fe 3+ solution. This forms a complex with Fe 3+ solution. A ligand (SCN -) had been added to form color of the solution, as for colorless solutions no absorption will take place in the 54
spectrophotometer. again the Fe3+ ion concentration is so minute that it couldn’t form any color, that’s why the ligand (NH 4SCN) had to be added. The plot drown in Spectrophotometric estimation can be used for determining concentrations of other solutions of the same species whereas where as in the other methods a plot can be used only for that respective solution.
55
ELECTROGRAVIMETRIC ANALYSIS Sample: ample: Br B r ass ass Sampl ampl e Containi ng Cer Cer tain Am ount Of Coppe Copperr An d Zi nc
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 an ELECTROLYTE (such (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. 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. 56
The current I is directly proportional to the resis tance R, i.e. I = E/R 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 – occurred – Reaction at Cathode:
Reaction at Anode:
Cu
2
4OH
2e
O2
Cu
2 H 2 O 4e
When the sample was dissolved in conc.HNO 3 solution the following reactions took place++ - Cu + 4H 4H NO O O 3 3 = Cu + 2NO 3 3 + 2NO 2 2 + 2H 2 ++ - Zn + 4H 4H NO O O 3 3 = Z n + 2NO 3 3 + 2NO 2 2 + 2H 2
The beneficial effect of nitrate nitr ate ion is due to its depolarizing action at the cathode. The reaction is 57
NO3
10 H
8e
NH 4
3 H 2 O
The reduction potential of NO 3- 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 NO 3- for depolarization and also for liberation of nitrate ion i on (NO3-) as NO2 when it is warmed. The reaction is – is – NO3
NO2
1 2
O2
e
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 Electrogravimetric analysis
58
In order to carry out the electrolysis properly properl y magnetic stirrer was used for constant stirring st irring 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 – experiment – By using an electronic balance weight of brass and platinum gauge cathode were taken; Brass was then taken into a beaker and concentrated HNO 3 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 C 2H5OH 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 br ass was calculated. 59
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 s ubstance can be measured.
Limitations Some limitation of this electro gravimetric method is also present al though this process is very much advantageous. They are listed below – 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 usuall y 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 thes e 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.
60
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 NO 3- as NO2 gas, but in this case excess heat was given which oxidized Cu into CuO, Cu + ½O 2 = 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 + + 2NO CONH C ONH 2 = 2N 2 O O 2 2 + N H 2 2 + CO 2 2 + 3H 2
- by boiling the HNO HNO 3 before adding it; or -
by adding sulphamic acid to the solution, - - - H + + N O O O 2 2 + SO 3 3 N H 2 = N 2 2 + H SO 4 4 + H 2
If any of the above said measures are not performed, pe rformed, 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 H 2 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 s urface, - - - + N O O O 3 3 + 10H + 8e = N H 4 + 3H 2
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.
61
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.
62
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 determinati on of the amount of an element, find out end point of solution etc. and so on. To deal with these types t ypes 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, di gits, 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.
63
Bibliography Bahl, B. S., Tuli, G.D., Bahl, A., Esse Essenti als of of Ph ysi ysi cal Chemis Chemi str y , Millennium edition, “
”
S.Chand Publications, (2000). Besset & associates,
Vogel’s Textbook of Quantitative Inorganic Analysis , 5th Vogel’s
“
”
ed.,Addison Wesely Longman Inc,(1999). Electrochemistry” , International edition, East-West Glasstone, S., An Introduction to Electrochemistry” “
Press Private Ltd, (1942). Physi cal cal Chemi Chemis str y , 4th edition, McMillan Maron, S. H., Prutton, C. F., “ Pri ncipl es of Physi ”
Publishing Co.,(1965). Williard, H. H., Merritt, L.L., & Dean, J.A., “ Instrumental method of analysis”, 4th ed., Nostrand company, London. London.
64
Appendices
Containing Observed data Calculated data Sample calculation graphs
65
Appendix-a POTENTIOMETRIC TITRATION of iron (Fe2+) Concentration of K 2Cr 2O7 , S1= 0.1(N) Concentration of H2SO4=6.0(N) Volume of ferrous (Fe ++) solution, V2=10 ml Required volume of Kr 2Cr 2O7 , V1=4.45ml Table 1: Observed and calculated data for volume, cell EMF, first and second derivative of cell EMF Volume of K 2Cr2O7 (ml) 0 0.2 0.4 0.6 0.8 1 1.2 1.4 1.6 1.8 2 2.2 2.4 2.6 2.8 3 3.2 3.4 3.6 3.8 3.9 4 4.1 4.2 4.3 4.4 4.5 4.6 4.8 5 5.2 5.4 5.6 5.8
6
E (Volt)
dV
dE
dE/dV
d2E/dV2
0.35125 0.3575 0.36625 0.37375 0.3825 0.38875 0.39375 0.3975 0.40125 0.4075 0.41 0.41375 0.41875 0.4225 0.42875 0.4325 0.44 0.445 0.4525 0.46 0.465 0.4725 0.48 0.48625 0.49625 0.50625 0.65125 0.675 0.685 0.6925 0.6975 0.7 0.70375 0.7075 0.71125
0.2 0.2 0.2 0.2 0.2 0.2 0.2 0.2 0.2 0.2 0.2 0.2 0.2 0.2 0.2 0.2 0.2 0.2 0.2 0.1 0.1 0.1 0.1 0.1 0.1 0.1 0.1 0.2 0.2 0.2 0.2 0.2 0.2 0.2 0.2
0.00625 0.00875 0.0075 0.00875 0.00625 0.005 0.00375 0.00375 0.00625 0.0025 0.00375 0.005 0.00375 0.00625 0.00375 0.0075 0.005 0.0075 0.0075 0.005 0.0075 0.0075 0.00625 0.01 0.01 0.145 0.02375 0.01 0.0075 0.005 0.0025 0.00375 0.00375 0.00375
0.03125 0.04375 0.0375 0.04375 0.03125 0.025 0.01875 0.01875 0.03125 0.0125 0.01875 0.025 0.01875 0.03125 0.01875 0.0375 0.025 0.0375 0.0375 0.05 0.075 0.075 0.0625 0.1 0.1 1.45 0.2375 0.05 0.0375 0.025 0.0125 0.01875 0.01875 0.01875 0.118542
-0.0625 0.03125 -0.03125 0.0625 0.03125 0.03125 -1.3E-15 -0.0625 0.09375 -0.03125 -0.03125 0.03125 -0.0625 0.0625 -0.09375 0.0625 -0.0625 2.36E-15 -0.125 -0.25 -3.3E-15 0.125 -0.375 8.88E-15 -13.5 12.125 0.9375 0.0625 0.0625 0.0625 -0.03125 4.16E-16 2.78E-15 0.016632 -0.0625 66
Sample Calculation 2+ Strength trength of F e at the th e end poi nt
Concentration of K 2Cr 2O7 ,, S1 = 0.1N From graph required volume of K 2Cr 2O7 for titration,V1 = 4.45 ml Volume of ferrous (Fe ++) solution, V 2 =20 ml Concentration of Fe 2+ solution, S2 = ? Now, V1S1 = V2S2 S2 = (4.450.1) / 20 = 0.02225N
3+ Dete Deter min ation of Concentr Concentr ation of F e ion F or 4 dif ferent poin poin ts of the Graph
For V=3ml K 2Cr 2O7 ,Ecell=0.4325 V Now,Ecell=EFe2+/Fe3+ +Ec =>
EFe2+/Fe3+ = Ecell - Ec……………….(Eq-1) ……………….(Eq-1)
Here,Ec =0.244 V for saturated KCl solution So,from Eq-1, E Fe2+/Fe3+ =0.1885 V Again, E Fe2+/Fe3+ = E0Fe2+/Fe3+ -
ln
3+
2+
……..(Eq-2) ……..(Eq-2)
Here, E0Fe2+/Fe3+ = - 0.771 V R=8.314 J/(mol.K) T=320C =305 K (Room temperature) F=96500 Faraday
From Eq-2,
3+
2+
= 1.387×10-16
Strength of Fe2+ at the point where 3ml K 2Cr 2O7 is used can be calculated from equivalence relation. 4.45 ml 0.1N K 2Cr 2O7 = 0.02225 N 20 ml Fe 2+ 1ml 0.1 N K 2Cr 2O7 = 0.005N 20 ml Fe 2+ 3 ml 0.1 N K 2Cr 2O7 = 0.015 N 20 ml Fe 2 So, aFe2+ =0.015 N Now, aFe3+ =1.387×10-16 × aFe2+ =2.08×10-18 N 67
APPENDIX-B
CONDUCTOMETRIC TITRATION of an acid mixture(Hcl+CH 3COOH) Obse Obser ved ved and an d cal cul ate ated d 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) 0.0
Conductance (mS) 2.30
Volume of NaOH (ml) 5.4
Conductance (mS) 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 Burette Reading Observation No.
Volume of NaOH(ml)
01 02
10 10
IBR(ml)
0 20.2
FBR(ml)
20 40.2
Volume of Oxalic Acid(ml)
20 20
Sample Calculation Standardization Standardization of NaOH by H 2C2O4
Here, volume of NaOH, V 1 = 10 ml Volume of H2C2O4 , V2 = 20 ml 71
Strength of H 2C2O4 , S2 = 0.1N Strength of NaOH = S 1 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, V 1 = 9.5ml (from graph) Strength of NaOH, S 1 = 0.2N Volume of HCl solution, V 2 = 10ml Strength of HCl = S 2 Again, V1 × S1 = V2 × S2 S2 = (9.5 × 0.2)/ 10 = 0.19N Now, Equivalent weight of HCl, e = 36.5 Weight of HCl = W 1 Again, W 1 = S2 e V2 = (0.19 × 36.5 × 10)/ 1000 = 0.06935g
Calculation of CH3COOH
Here, volume of NaOH for neutralizing CH 3COOH, V1 = (20.1-9.5) = 10.6ml Strength of NaOH, S 1 = 0.2N Volume of CH3COOH solution, V 2 = 10ml Strength of CH 3COOH = S2 Again, V1 × S1 = V2 × S2 S2 = (10.6 × 0.2)/ 10 = 0.212N Now, Equivalent weight of CH3COOH, e = 60 Weight of CH3COOH = W 2 Again, W 2 = S2 e V2 = (0.212 × 60 × 10)/ 1000 = 0.1272g Percentage (%) of HCl = [W1/ (W1+ W2)] × 100% 72
= [0.06935/ (0.1272 + 0.06935)] × 100% = 35.28%
Percentage (%) of CH 3COOH = [W2/ (W1+ W2)] × 100% = [0.1272/ (0.1272 + 0.06935)] × 100% = 64.72%
73
APPENDIX-C pH
METRIC TITRATION
Observed and calculated data Table 4: Observed and Calculated data for pH-matric titration Volume(V)
pH
0
9.93
0.8
d2(pH)/dV2
dV
d(pH)
d(pH)/dV
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
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
1
Sample Calculation Concentration of Na 2CO3, S1 = 0.05 N From graph, Required volume of HCl acid for complete titration, V 2= 2.78 ml (from graph) Volume of ferrous Na2CO3 solution, V1= 10 ml Now, V1S1 = V2S2 S2 =
×. .
= 0.1798 N
Calculations for dissociation constants
Volume of HCl for converting Na 2CO3 to NaHCO3, V1 = 1.17 ml Dissociation constant of Na 2CO3, pk 1 = the corresponding pH for V1/2 ml of HCl = 7.7 (From Graph) 75
Volume of HCl for neutralizing NaHCO 3, V2 = 2.78-1.17=1.61 ml Dissociation constant of NaHCO3, pk 2 = the corresponding pH for (V1 + V2 /2) ml of HCl = 5.25(From Graph)
76
APPENDIX-D
Spectrophotometric Determination of Iron Observed Data Peak Pick Wavelength, λ max max = 478.5nm Absorbance = 0.4187
Table 5: Table for observed data of absorbance of different concentration 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
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
81
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%
83
