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Index

S.No. Name of the experiment Page No.

1. Reaction kinetics - Decomposition of benzene-diazonium

chloride

2

2. Enzyme catalysis using UV-Vis spectrophotometer 4

3. Kinetics of iodination of acetone 6

4. Determination of pKausing UV-Vis spectrophotometry 8

5. Heat of solution by solubility measurement 10

6. Estimation of excess thermodynamic properties 12

7. Estimation of thermodynamic functions from EMF data 15

8. Determination of stability constant of silver-ammonia

complex

17

9. Electrochemical oxidation of L-cystine to L-cysteic acid 19

10. Verification of Beer-Lambert Law using gold nanoparticles 21

11. Estimation of free energy of protein denaturation usingintrinsic fluorescence of protein

25

12. Fluorescence quenching of fluorescein by inorganic anions:

Determination of quenching constants and estimation of free

energy change and activation energy for the quenching process

27

13. Critical micelle concentration of CTAB 29

14. Acidity determination in zeolites by temperature programmed

desorption (TPD) of ammonia

31


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1. REACTION KINETICS - DECOMPOSITION OF BENZENE-

DIAZONIUM CHLORIDE

1.1 Aim

To determine the first order rate constant for the decomposition of benzene-diazonium chloride.

1.2 Chemicals required

Aniline, conc. HCl, sodium nitrite (NaNO2)

1.3 Apparatus required

5 mL measuring flask, 200 mL standard flask, ice water, stop watch and dropping funnels

1.4 Principle

Guggenheim method is useful to obtain the first order rate constant when the initial amount of

reactant cannot be determined by employing a physical method such as measuring the change inpressure or volume in the course of a reaction. In this method a series of readings of the

concentration of product (reading Y) at times t that are spread over an interval which is two to

three times the half life of the reaction. A second series of readings (Y) can also be made, each

at time (t+ ), where is an exactlyconstant interval after the time of the corresponding reading

Y. The period must be at least two to three times the half life of the reaction.

Integrated first order rate law: We can also write

ln ( ) = When benzene diazonium chloride (BDC) is heated in acidic aqueous solution it decomposes,following first order kinetics, to yield phenol, HCl and N2. The rate constant is determined as a

function of temperature so that the activation energy for the reaction can be calculated.

NaNO2+ HCl HNO2+ NaClThe nitrous acid formed is reacted with aniline to give benzenediazonium chloride:

C6H5NH2+ HNO2+ HCl [C6H5N+N][Cl

]

1.5 Procedure

1. Dissolve 1.3 mL of aniline in 5 mL of conc. HCl

2. Cool the solution in an ice bath.

3. Prepare a solution of sodium nitrite by dissolving 2 g of the salt in ~ 30 mL of water.

4. Add slowly with stirring the sodium nitrite solution to the aniline solution maintaining the

temperature at 0 C.
http://en.wikipedia.org/wiki/Sodium_nitritehttp://en.wikipedia.org/wiki/Sodium_nitritehttp://en.wikipedia.org/wiki/Sodium_nitritehttp://en.wikipedia.org/wiki/Sodium_nitrite


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5. Dilute the resulting solution to 200 mL with ice cold water.

6. Transfer 50 mL of solution to the reaction vessel and put it into the constant temperature bath

maintained at 45 C.

7. After 10 minutes connect the T- tube into the burette filled with water. When water level in

the burette is zero note this initial reading as Vo. Start the stop watch.

8.

Take a series of 40 readings, Vi------Vn, at times ti-------tnat fixed intervals of time (30 sec).

9. Choose a time interval (t) such that it is about 2-3 times the time of the half reaction. For

every Vtat time t note the corresponding Vt+tat t+tas in table below

S. No. Time

(t) sec

Volume

(vt) mL

S. No. Time

(t+ t)sec

Volume

(Vt+t) mL

Log (Vt+tVt)

10.Plot Log (Vt+tVt) versus t. Find the slope for the straight line. 2.303 slope = -k

Table and graphs:

Result

The first order rate constant of the decomposition of benzene diazonium chloride is k = 0.961 x

104sec

-1(literature value)

References

1. P.W. Atkins, Physical Chemistry, 6th ed. pp. 767-769 (1998);

2. B. Viswanathan and P.S. Ragavan, Practical physical chemistry, pp. 132-134 (2005)

Log10

(V1-V)

Time (t)


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2. ENZYME CATALYSIS USING UV-VIS SPECTROPHOTOMETER

2.1 Aim

To verify MichaelisMenten equation

2.2 Chemicals required100 mL of 0.1 % Starch, 100 mL of 0.0006 N Iodine, enzyme (10 mg - amylase in 100 mL

water)


UV-Visible spectrophotometer, 12 boiling test tubes, four 50 mL burettes

2.4 Principle

12

1

k

k

max M0

M 0 max max

1 2max 2 0 2

1

Michaelis Menten Mechanism

E + S ES E + product

V [S] K1 1 1Initial velocity, V = = +

K +[S] V V S

where V = k [E] ; . k is called turnover number

k

M

V

k kK

k

2.5 Procedure

1. Prepare stock solutions of 100 mL 0.1% starch solution, 100 mL 0.0006 N Iodine and

enzyme (10 mg - amylase in 100 mL water). Fill them in 50 mL burettes.

2. Take 12 boiling test tubes and prepare the following solutions

Set 1: Blank solutionsS.No. 0.1 % starch (mL) Iodine (mL) Water (mL) Conc of starch %

1 3 1 6 0.03

2 4 1 5 0.04

3 5 1 4 0.05

4 6 1 3 0.06

5 7 1 2 0.07

6 8 1 1 0.08

Set 2: Test solutions to which - amylase is to be added while starting the reaction

S.No. 0.1 % starch (mL) Iodine (mL) Water (mL) Conc of starch %

7 3 1 5 0.03

8 4 1 4 0.04

9 5 1 3 0.05

10 6 1 2 0.06

11 7 1 1 0.07

12 8 1 0 0.08


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3. Set the spectrophotometer at 590 nm and take absorbances of first six solutions (Set 1).

4. To the 7th

solution (Set 2) add 1 mL of enzyme solution while starting the stop watch. Mix

solution well transfer into a cuvette. Monitor the reaction for about 3 min by measuring the

absorbance at every 10 sec.

5. Repeat the step 4 with solutions 8. 9. , 12. Tabulate the reading as follows

S. No. Timesec

0.03%

0.04%

0.05%

0.06%

0.07%

0.08%

Blank

10

20

30

180

6. Plot absorbance of iodine versus time. Obtain the initial rate [V]0for every concentration of

starch, [S]7. Plot [V]0versus [S] and verify the Michaelis-Menten equation.

8. Plot 1/ [V]0 versus 1/[S] and obtain the MichaelisMenten constant KM, according to the

linear equation given above.

Results

1. Michaelis-Menten equation is verified.

2. The Michaelis-Menten constant KMis

References

1. K. J. Laidler, Chemical Kinetics, 5th

edition.

2.

P.W. Atkins, Physical Chemistry 6thedition (1998).

1/V

1/[S]

V

[S]


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3. KINETICS OF IODINATION OF ACETONE

3.1 Aim

To study the kinetics of iodination of acetone using a UV-visible spectrophotometer

3.2 Chemicals requiredIodine (0.005 M), 0.5 M acetone and 0.5 M HCl


UV-Visible spectrophotometer, 25 mL standard flasks (6), 5 mL, 10 mL, 25 mL burettes, 5 mL

pipette

3.4 Principle

The experiment is conducted with a view to study the kinetics of iodination of acetone. The

balanced equation cannot predict the rate law. The rate law can only be obtained by experimental

determination. The form of rate law would help to reach conclusions about the mechanism of thereaction.

The chemical equation for this reaction is (CH3)2CO + I2 CH3-CO-CH2I + I +H+

The reaction rate increases with the concentration of H+ in acid solutions and with OH

concentration in basic solutions. The rate of halogenation of acetone is independent of the

concentration of the halogen except at very high acid concentration.

3.5 Procedure

1. Suitable concentrations of iodine, acetone and HCl solutions are prepared (0.5 M each)

2.

Different compositions of the reaction mixtures are prepared in 25 mL standard flasks asshown in the chart below. Iodine, water and acetone are taken in 25 mL standard flask. HCl

should be added only at the time of starting the reaction.

S. No Volume of

I2

(mL)

Volume of

HCl

(mL)

Volume of

H2O

(mL)

Volume of

Acetone

(mL)

1 10 1.25 1.25 12.52 5 1.25 6.25 12.5

3 7.5 2.5 2.5 12.5

4 7.5 2.5 12.5 2.5

5 7.5 0.5 4.5 12.5

3. When half of the HCl has been added, start the stopwatch. Mix the solution well and transfer

to the cuvette and place in the UV-Vis spectrophotometer. Measure the absorbance at 450 nm

at various intervals of time for 1520 minutes.


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3.6 Calculations

1. In each case absorbance is plotted as a function of time.

2. From the plot of optical density versus time, the initial slope is obtained which gives the

initial rate of the reaction.

3. From (1) and (2) the order with respect to iodine is determined.

4.

From (3) and (4) the order with respect to acetone is determined.

5. From (3) and (5) the order with respect to acid can be found.

Table and graphs:

Results

1. The order with respect to iodine is

2. The order with respect to acetone is

3. The order with respect to acid is

References

1. S.W. Benson, The Foundations of Chemical Kinetics, McGraw-Hill, pp 569-573, (1960).2. A.R. Knight, Introductory Physical Chemistry, Prentice-Hall, Chapter 6, (1970).

3. J.L. Latham, Elementary Reaction Kinetics, Butterworths (1964).

4. G.M. Harris, Chemical Kinetics, D.C.Heatch and Company (1966).

Absorbance Time

Ab

sorbance

Time


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Table and graph

Results

The pKavalue of the given weak acid is 5.0 (literature value)

(i) Calculated

(ii) From graph

Reference

1. Hideo Yamazaki, R. P. Sperline, and Henry Freiser,Analytical Chemistry, Vol.64, 2720-

2725 (1992).

S. No. pH Abase Aacid Abase/Aacid Log[Abase/Aacid]

LogA

base

Aacid

pH


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5. HEAT OF SOLUTION BY SOLUBILITY MEASUREMENT

5.1 Aim

To determine H, G

, S

of solution for potassium dichromate by solubility measurements


Starch, K2Cr2O7, KI, H2SO4 (2 N), sodium thiosulphate

5.3 Glassware Required

Pipette, 20 mL measuring cylinder, conical flask, weighing bottle, beaker and thermometer

5.4 Reactions involved

K2Cr2O7 + 6KI + 7H2SO4 Cr2(SO4)3 + 4K2SO4 + 7H2O + 3I2

I2 + 2Na2S2O3 2NaI + Na2S4O6

5.5 Principle

Solubility of substances, in general, increases with temperature. There are a few exceptions to

this observation. In all cases, the effect of temperature on the solubility of a substance can be

predicted from its heat of solution if no chemical reaction occurs during dissolution.

Le Chateliers principle indicates that the substances which absorb heat during dissolution

process (H is positive), dissolve to a greater extent at higher temperature. In the same

way it can be shown that the substances, which evolve heat (His negative) during dissolution,

dissolve to a greater extent at low temperatures.

From thermodynamics, the following relationship can be derived

where S1and S2are the solubilities of the substances in 100 g of water at temperatures T1and T2

in K and R is the gas constant in JK-1

mol-1

. T2is greater than T1.

5.6 Procedure

Prepare saturated solutions of potassium dichromate around 20 C, 30 C (room

temperature) and 40 C by keeping in the ice cold water bath and hot water bath as the case may

be. Keep them immersed in bath for around 5 minutes for equilibrium. Note down the roomtemperature.

Quickly pipette out 1 mL of the saturated solution at 20oC and deliver into a weighing

bottle previously weighed. Allow sometime for the solution to attain the room temperature.Wipe out the outside by means of filter paper gently to remove the condensed moisture and

weigh again. From these two weights, calculate the density of the saturated solution at 20 C.

By the same procedure, find out the density of the saturated solution of K2Cr2O7at 40 C and atthe room temperature.


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Solubility determinations

1. Pipette out 1 mL of saturated solution of K2Cr2O7at 20 C into a conical flask.

2. Add about 15 mL of 2 N H2SO4and a little solid KI.

3. Titrate the liberated I2against standard 0.2 N thiosulphate solution using starch as indicator.4. From the normality of potassium dichromate solution calculate the weight of K2Cr2O7 in 1

mL of the solution.5. Let it be X. If X1is the weight of 1 mL of the saturated solution at 20 C the solubility S1

at 20 C is given by 6. Calculate the solubility S2of K2Cr2O7at 30 C and S340 C in a similar fashion.

7. From the equations described earlier, calculate the H, G

, S

of solution, using the

solubilitysat 20 C, 30 C and 40 C.

Results

Solubility (in gm/ 100 gm water) of potassium dichromate at 20C, 30C and 40C respectively

are 12.3, 18.1, 26.3

The H, G

, S

of solution are

Reference

1. F. Daniels and R.A. Alberty, Physical Chemistry, Wiley pub., 3rd Ed. (1996)


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6. ESTIMATION OF EXCESS THERMODYNAMIC PROPERTIES

6.1 Aim

Estimation of excess thermodynamic properties of binary liquid mixture from density and

ultrasonic velocity measurements


Benzene and cyclohexane

6.3 Glassware required

Beaker, Pipette (5 mL, 10 mL) and pycnometer

6.4 Principle

Raoult's Law for the mixture of two volatile liquids states that the total vapour pressure ofthe mixture is equal to the sum of the individual partial pressures for ideal solution.

Total Vapour Pressure = PA+ PB

Many pairs of liquids are present in which there is no uniformity of attractive forces i.e.the adhesive & cohesive forces of attraction are not uniform between the two liquids, so that they

show deviation from the Raoult's law.

Deviation from ideal behaviour: When adhesive forces between molecules of A & B are greaterthan the cohesive forces between A & A or B & B, then the vapor pressure of the solution is lessthan the expected vapor pressure from Raoult's law. This is called as negative deviation from

Raoult's law. e.g. chloroform and acetone show such an attraction by formation of a hydrogen

bond.

When the cohesive forces between like molecules are greater than the adhesive forces,the dissimilarities of polarity or internal pressure will lead both components to escape solution

more easily. Therefore, the vapor pressure will be greater than the expected from the Raoult's

law, showing positive deviation. If the deviation is large, then the vapor pressure curve will showa maximum at a particular composition, e.g. benzene & ethyl alcohol, carbon disulfide &

acetone, chloroform & ethanol.

Excess thermodynamic properties:Knowledge of excess thermodynamic properties of the multi-component systems is essential for the design calculations involving separations, heat transfer,

mass transfer, and fluid flow. Further, these properties have relevance in the theoretical research

to get the insight into intermolecular forces and microscopic structure of liquids.The excess molar volumes and excess isentropic compressibility are the properties very

sensitive to different kinds of association in the pure components and in the mixtures and often

they are related to local order. They have been used to investigate the molecular packing,

molecular motion and various types of intermolecular interactions and their strength, influenced

by the size, shape and the chemical nature of component molecules.

6.5 Procedure

1. Find the weight of the dry pycnometer.

2. Prepare 10 mL binary mixture of benzene and cyclohexane with 0.2, 0.4, 0.6 and 0.8 mole

fractions by narrow-mouth specific gravity bottles / pycnometer taking due precautions tominimize the evaporation loses.


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3. For preparation of binary, measure the volume as specified in table below. As exact volume

is difficult to measure so after transferring the required volume (v1, v2) measure the

corresponding weight (w1, w2)4. The table given below is of standard value and is merely to guide you the amount of volume

you have to take. Make your own table as below only excluding v1, v2column.

Benzene + cyclohexane mixture

S.No. x1 x2 v1 v2 w1 w2

1 0.0 1.0 0.0 10.0 0.000 7.790

2 0.2 0.8 1.7 8.3 1.498 6.458

3 0.4 0.6 3.5 6.5 3.110 5.026

4 0.6 0.4 5.5 4.5 4.847 3.482

5 0.8 0.2 7.7 2.3 6.726 1.8126 1.0 0.0 10.0 0.0 8.765 0.000

28.4 31.6 24.946 24.568

5. Calculate X1and X2by the given formula

X1=

X2= 1-X1Where wi,Mi are weight and molar weight respectively

6. Calculate density by weighing the pycnometer with sample for each solution.

7. Measure u for each binary solution and pure component with the help of Ultrasonic

interferometer.8. Measure ultrasonic velocity of two pure liquids and four binary mixtures by using the formula u

= where is wavelength and is frequency of ultrasonic wave (both and are

obtained from interferometer)9. Repeat the procedure number 6, 7 and 8 at 50

oC using water bath

10.From the density and ultrasonic velocity data of two pure liquids and four mixtures calculate the

excess thermodynamic properties using following equations:

Excess molar volume: Where xi, Miand irepresents mole fraction, molar weight and density ofi

thcomponent in

mixtures respectively

Isentropic Compressibility: or S= 1/ (.u 2)

Excess Isentropic Compressibility (Eor S

E):

{

}


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Vi, iand Cp,iare the molar volume, isobaric thermal expansion coefficient and molar isobaric

heat capacity respectively of pure component iand i= xiVi/xjVjis the volume fraction of iin

the mixture, stated in terms of the unmixed components.

7. Draw the graph between and T for each solution and find out the slope. Calculate i bydividing the slope with density as given by the equation

= - Results

Plot of excess molar volume and excess isentropic compressibility of binary mixture versus

mole fraction of component A are shown

References

1. S.L. Oswal, M.M. Maisuria1, R.L. Gardas , Journal of Molecular Liquids, vol.109, 155166

(2004).

2. Rowlinson, and F.L. Swinton, Liquid and Liquid Mixtures, 3rd Ed. Butterworth, London,

(1981).


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7. ESTIMATION OF THERMODYNAMIC FUNCTIONS FROM EMF

DATA

7.1 Aim

I.

To determine the standard electrode potential of Cu

2+/

Cu half cell.II.

To calculate the thermodynamic functions G, S

, H

.


Copper sulphate, distilled water


Copper electrode, 100 mL standard flask, 25 mL standard flask (4), 50 mL beaker, salt bridge,

calomel electrode, potentiometer, water bath and thermometer

7.4 Principle

The electrode potential is defined as the tendency of the electrode to loose or gain electrons.Here the potential of the electrode is determined using a reference electrode (calomel electrode).

When a chemical reaction is carried out reversibly and isothermally in a galvanic cell, the

decrease in free energy is the net electrical work done by the system ie G

= nFE(joules/mole) where n is the number of electrons involved in the redox reaction, F is the faraday

constant (96497 coulombs/equivalent) and E is the reversible EMF of the cell.

7.5 Procedure

1. Prepare 25 mL of each 0.1 M, 0.05 M 0.01 M 0.005 M of copper sulphate solution.

2. For determining the standard electrode potential, set up a galvanic cell by dipping a copperelectrode into 0.005 M solution of copper sulphate solution taken in a beaker and combine

this with the reference electrode (calomel electrode) through salt bridge.3. Connect the positive end of the potentiometer to the copper electrode and negative end of the

potentiometer to the calomel electrode.4. Switch on the potentiometer and note the reading.

5. The cell is placed in the water bath and the EMF is measured at different temperatures (20

C, 25 C (RT) and 30oC)

6. Repeat the experiment with each of the above indicated concentration of copper sulphate.

7.6 Calculations

Write the half cell representations of the both the half cells.

Calculate the standard electrode potential of copper electrode from the measured EMF values,

molarity of copper sulphate.Eobs= E1E2E2= EMF of the standard calomel electrode (0.2422 V vs SHE) and E

0can be calculated from

is the activity of copper ions (use concentration in place of activity for calculationpurpose).


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From the potential data calculate the G, S

, H

from the below equations.

G= nFE1

H

= G

+T S

Results

1.

Concentration of copper

sulphate (M)

EMF of the concentration cell

taken at (temp)

EMF of the concentration cell

taken at (temp)

0.1

0.05

0.01

0.005

2.Values of G, S

, H

(for any two concentrations)

Concentration ofcopper sulphate (M)

H (kJ mol- ) S (J K - mol- ) G (kJ mol- )

0.1

0.01

References

1. D. E. Smith, J. Chemical Education, vol.60, 1983, 299.2. P. A. Rock, J. Chemical Education, vol.52, 1975, 787.

3. C. A. Vincent J. Chemical Education, vol.47, 1970, 365.


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8. DETERMINATION OF STABILITY CONSTANT OF SILVER -

AMMONIA COMPLEX

8.1 Aim

To determine the stability constant of silver- ammonia complex according to the reaction

Ag++ 2 NH3 [Ag(NH3)2]+

8.2Chemicals required

0.1 N AgNO3solution, 4 N ammonia solution


50 mL beaker, silver electrode, salt bridge, 100 mL standard flask, 50 mL standard flasks (5),

potentiometer.

8.4 Principle

The determination of stability constant of a complex is based upon the determination of the

activity of the metal ion in a dilute solution containing the complex. The determination of themetal ion activity is as usual done by constructing a suitable galvanic or concentration cell. The

stability constant of the equilibrium Ag++ 2 NH3 [Ag(NH3)2]

+is given by

Where x = conc. of complex at equilibrium

a = initial conc. of AgNO3b = initial conc. of ammonia

Using the Nernst equation, from the conc. cell we will obtain conc. of free ions in the complex

solution i.e., C1. Therefore Substituting x in eqn1 we get

() 8.5 Procedure

1. The following concentration cell is set up

Ag| 0.025 N AgNO3in 1 N NH3|| KNO3(sat) || 0.01 N AgNO3| Ag

Where in the pure 0.01 N AgNO3 is taken as reference.

2. Two silver electrodes are dipped in the respective half cells and the silver electrode placed in

the solution of the complex is connected to the negative end of the potentiometer while the

silver electrode is placed in silver nitrate solution is connected to the positive end of the

potentiometer. The voltage reading is noted.3. The experiment is repeated with different concentrations of AgNO3. (i) 0.025 N in 1 N NH3

(ii) 0.0125 N AgNO3in 1 N NH3(iii) 0.006 N AgNO3 in 1 N NH3(iv) 0.0125 N AgNO3in 2

N NH3(v) 0.0125 N AgNO3in 1.5 N NH3.

8.6 Calculations

The observed EMF of the cell is given by the equation


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{

}= { }

From eqn3, C1i.e., the conc. of free Ag+ions in the complex equilibrium, is calculated.

Substituting a, b, C1in eqn2 we obtain the equilibrium constant K.

For Example in (i) 0.025 N AgNO3 in 1 N NH3; a = 0.025, b = 1, and C1is given from the eqn3substituting these in eqn2 we get the equilibrium constant K.

Result

The stability constant of the silver ammonia complex is 1.67 x107(literature value)

References

1. V. D. Athawale, Experimental Physical Chemistry, New Age International Ltd, page 208.

2. R. W. Ramette, J. Chemical Education vol.49, 1972, 423.


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9. ELECTROCHEMICAL OXIDATION OF L-CYSTINE TO L-

CYSTEIC ACID

9.1 Aim

To determine the mechanism of electrochemical oxidation of L-cystine to L-cysteic acid and

hence to find the diffusion coefficient of Br-ion by using Randels-Sevcik equation


10 mM L-cystine, 10 mM HBr, 0.5 M H2SO4


Electrochemical cell, glassy carbon electrode, platinum electrode and Calomel electrode

9.4 Principle

The oxidation of L-Cystine to L-Cysteic acid is performed by electrochemically generated

bromine using HBr solution.

The effect of peak current on scan rate and concentration in cyclic voltammetry is given by

Randels-Sevcik equation which is given by

Where ip= peak current (A), n= no of electrons, A= area of the electrode (cm2)

, D- diffusion

coefficient (cm2s

-1), C- concentration of L-Cystine (mol cm

-3), -scan rate (Vs

-1) .

9.5 Procedure

1. The electrochemical cell and the three electrodes are rinsed with distilled water.

2. 10 mL of 0.5 M sulphuric acid is taken in the cell. The electrode is immersed in the solutionand connected to the potentiostat.

3. The Cyclic Voltammetry is run with the following parameters

a. Initial potential 0.0 V and final potential 1.8 V vs SCE.

b. The current range is assigned in the milliampere range.

c.

Scan rate is chosen as 0.05 V/s.4. The CV is started by giving the run command. The data is then stored.

5. To the 10 mL of 0.5 M sulphuric acid add 10 mM in 10 mL of L-cystine. Record CV with theabove given parameters.

6. In another cell again 10 mL of 0.5 M sulphuric acid is taken and 10 mM in 10 mL of HBr is

added. Record the CV with the same parameters. Save the recorded CV7. To the solution of step 6 add 10 mM in 10 mL L-cystine and run the experiment.

8. Repeat the experiment with different scan rates.


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9. Repeat the steps 6 to 8.

10.Record the CV with different concentrations of L-cystine from 10 mM to 50 mM.

11.Save all the recorded CVs.

9.6 Calculations

1. A graph is drawn by taking concentrations of L-Cystine on the x-axis and peak current of

different concentration of L-cystine.2. A plot of peak current of L-Cystine and square root of scan rates is made.

3. From the slopes calculate the diffusion coefficient using Randels-Sevcik equation

The typical cyclic voltammogram recorded looks like the following

Figure1. cyclic Voltammograms of L-Cystine oxidation at glassy carbon electrode surface(a) 0.5 M H2SO4, (b) 10 mM of L-Cystine in 0.5 M H2SO4, (c) 10 mM of HBr in 0.5 M

H2SO4 and (d) 10mM of HBr in 0.5 M H2SO4with different concentrations of L-Cystine

from 10 mM to 50 mM.

Results

1. The electrochemical oxidation of L-cystine is performed and the mechanism is elucidated.

2. The linearity of the obtained from the two graphs verifies the Randels-sevcik equation.3. Diffusion coefficient obtained from Randels-Sevcik equation is 2.08 cm

-2sec

-1(literature)

References

1. Wang X, Zhao, Chemistry Letters, vol.33, 2004, 332.

2. G. Sanchez-cano, V. Montiel, A. Aldaz, Tetrahedron, vol.47,1991, 877.


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10. VERIFICATION OF BEER-LAMBERT LAW USING GOLD

NANOPARTICLES

10.1 Aim

I. Preparation of gold nano particles using the citrate method and verification of Beer-

Lambert law.II. Preparation of gold nanorods of different aspect ratios and to determine the effect of

concentration of NaBH4on the position of LSP.


Tetrachloroauric acid trihydrate (HAuCl43H2O), sodium borohydride (NaBH4), silver nitrate,cetyltrimethylammonium bromide (CTAB), ascorbic acid, trisodium citrate (Na3C6H5O7)


50 mL flat bottomed flask, 5 mL standard flasks (10), 50 mL, 100 mL, 250 mL standard flasks, 1

mL, 100 L pipettes, 25 mL measuring jar, heating mantle and cuvettes

10.4 PrincipleThe Beer-Lambert law, also known as the Beer's law or the Beer-Lambert-Bouguer, is the

linear relation between the absorbance and concentration of an absorbing species. According to

this law, when light passes though a material, there exists a linear dependence between the

absorbance (A) of light and the product of absorption coefficient () with the path length ( l) of

the light (distance though which light travels in the material). Absorption coefficient of a

material in turn depends on the molar absorptivity or molar extinction coefficient or molar

absorption coefficient () and the concentration (c) of the absorbing species in the material. Thus,according to Beer-Lambert law, absorbance is given by,

where is constant, for a given wavelength of light, and can be obtained from the absorbancevalues of solutions with different concentrations. Absorbance is unitless, while the units of , l, care L.mol

-1.cm

-1, mol.L

-1and cm

-1, respectively.

Appearance of wine-red color in step 6 indicates the formation of colloidal gold or gold

nanoparticles. The recorded UV-Vis spectrum of the solution (typically 400-1100 nm window),

in step 10, shows the absorbance maximum at a wavelength (max) of ~520 nm, as shown in Fig.1a. These features are attributed to the plasmon resonance of gold nanoparticles. The intensity of

the peak is reduced with increase in dilution.

500 600 700

0.3

0.6

0.9

1.2

Absorbance

Wavelength (cm-1

)

parent AuNPs

4 mL AuNPs + 1 mL water




500 600 700

0.3

0.6

0.9

1.2

Absorbance(a

.u.)

Wavelength (cm-1

)

B


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22

Fig.1a. UV-Vis spectra of gold nanoparticles of different dilutions. Fig.1b. Plot of absorbance vs.

concentration of gold nanoparticles

Fig. 1b shows the plot of absorbance vs. concentration for the gold nanoparticles of different

dilutions. It can be observed that all the data points fit to a straight line, thus conforming to theBeer-Lambert law. Slope of the above plot gives the absorption coefficient of the gold

nanoparticles. For the calculation of absorption coefficient, path length is being taken as equal tothe width of the cuvette used for recording the UV-Vis spectrum (we have taken it as 1 cm).From the above data, the value of absorption coefficient is calculated as ~3.36 L.mol

-1.cm

-1. By

making use of the above plot, we can calculate the concentration of gold nanoparticle of

unknown dilution.

Nanorods belong to the class of non-spherical or anisotropic nanoparticles. Such particlesrequire more than one parameter to describe them, such as length and width in the case of rod

shaped particles. Most of the synthesized nanorods are cylindrical in shape. They are

characterized by their aspect ratio which is the ratio of the major dimension to the minor

dimension. Nanorods are often classified as 1-D nanostructures and have less aspect ratio thanother class of 1-D nanostructures (e.g., nanowires or nanotubes). Size of a nanorod ranges within

1-100 nm and aspect ratio of a nanorod is greater than 1 but less than 20, with typical values

between 3 and 5. Gold nanorods belong to an important class of anisotropic nanomaterials.

Chemical methods to synthesize gold nanorods are very recent.

Absorption spectrum of gold nanorod is characterized by two surface plasmons. They are

transverse surface plasmon (TSP) resonance and longitudinal surface plasmon (LSP) resonance.

TSP of a nanorod is observed in the range of ~505-520 nm and LSP occurs at higher wavelength

(600-1300 nm). There is negligible blue shift in TSP with increase in aspect ratio, whereas LSPis very sensitive and increases with increase the aspect ratio of nanorods. In this experiment, we

will find a linear dependence between the position of LSP and aspect ratio.

Chemistry of gold nanorods is considerably different from colloidal gold due to its surface

structure and the presence of the {110} facet on nanorod surface owing to the use ofcetyltrimethylammonium bromide (CTAB) as the capping agent. CTAB stabilizes the {110}

facet, although it is of higher energy. In a nanocrystal, different planes have different surface

energies, which can be related to the stability (hence reactivity) of the surfaces.

Fig.2c. shows the TEM images of gold nanorods prepared this way at different magnifications.Aspect ratio of gold nanorods can be obtained from such a TEM image. There are also

theoretical ways of obtained this by calculating the absorption spectrum for a particle of specific

shape.

Fig.2a. shows the absorption spectrum of the gold nanorods having two surface plasmons at 511

and 796 nm corresponding TSP and LSP, respectively, suggesting an aspect ratio as 3.4. This is

obtained from an optimized synthesis. Aspect ratio can be tuned by varying the volume ofNaBH4added to the growth solution, as shown in Fig. 12.4. It is observed that the aspect ratio

decreased with increase in the volume of NaBH4used, finally resulting in spherical nanoparticles(aspect ratio, 1) at 0.5 mL (keeping all other parameters the same). Under optimum conditions,

gold nanorods are formed have an aspect ratio of 3.4, by the addition of 50 L of 1.67 mM

NaBH4(trace a). With further increase in the amount of NaBH 4(500 L of 1.67 mM NaBH4),the LSP merged with the TSP and a single absorption peak was obtained, characteristic of

spherical particles (trace e).


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23

50 nm 5 nm

A plot of the position of LSP and the concentration of NaBH4 is shown in Fig.2b. An

approximate estimate of residual NaBH4available in the volume of seed solution added to the

growth solution is calculated from the fact that one mole of BH4-can reduce eight moles of Au

1+

existing in the growth solution.

Fig.2A. Absorption spectrum of the gold nanorods having two surface plasmons. Fig.2B. UV-Vis spectra of gold nanorods formed with increasing amount of NaBH4: (a) 50 L, (b) 100 L,

(c) 200 L, (d) 300 L, and (e) 500 L. Fig.2C. Position of LSP versus effective NaBH 4

concentration in growth the solution

10.5 Procedure -I

1. Prepare 10 mM HAuCl4 solution in a 5 mL standard flask by dissolving 17 mg of

HAuCl43H2O in water. HAuCl43H2O is hygroscopic and a standard solution will be given

to you which may be diluted to make the required solution.

2. Prepare ~0.5% solution of trisodium citrate in a 10 mL sample bottle flask by dissolving 25mg of Na3C6H5O7in 5 mL of water.

3. Take 0.5 mL of 10 mM HAuCl4solution in a 50 mL flat bottomed flask using a 1 mL pipette

and add 13 mL of distilled water to it using a 25 mL measuring jar.

4. Heat the solution over a heating mantle and bring the solution just to a boil. (marked by the

appearance of bubbles from the colourless solution).

5. To the boiling solution, add 1 mL of ~0.5 % trisodium citrate solution using a 1 mL pipette.

6. Continue heating (with boiling) till the colour turns to wine-red (this may take about 1-2

minutes).

7. Remove the flask from the mantle and keep it for cooling in air, for about 15-20 minutes.

8. To account for loss of water during boiling, make the solution to 14.5 mL (i.e. to the original

volume) in a 25 mL measuring jar.

9. Calculate the concentration of gold nanoparticles in the solution in terms of gold.

10.Prepare the following solutions of different concentrations.

5 mL prepared solution

4 mL prepared solution + 1 mL distilled water


450 600 750 900 10500.0

0.4

0.8

1.2

Absorbance(a.

u.)

Wavelength (nm)

e

c

d

b

a

A

0.0 0.2 0.4 0.6650

700

750

800

850

PositionofLSP(nm)

Effective NaBH4(

mole)

B


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24



11.Collect the UV-Vis spectrum of undiluted solution along with those of above diluted ones,with distilled water as the reference, using UV-Vis absorption spectroscopy.

12.

Note the value of absorbance maximum and the corresponding wavelength (max) of the eachspectrum.

13.Plot a graph of absorbance vs. concentration using the values obtained. Slope of the aboveplot gives the absorption coefficient of the gold nanoparticles.

10.6 Procedure - II

1. Prepare the growth solution containing 5 mL of 100 mM CTAB, 250 L of 10 mM

HAuCl43H2O, 32.5 L of 10 mM AgNO3, and 35 L of 100 mM AA.

2. Prepare 10 mL of 1.67 mM of NaBH4in ice-cold water.

3. Prepare various samples of the following proportions:

a.

50 L of freshly prepared ice-cold NaBH4+ growth solution of (~5 mL)

b. 100 L of freshly prepared ice-cold NaBH4+ growth solution of (~5 mL)

c. 200 L of freshly prepared ice-cold NaBH4+ growth solution of (~5 mL)

d. 300 L of freshly prepared ice-cold NaBH4+ growth solution of (~5 mL)

e. 400 L of freshly prepared ice-cold NaBH4+ growth solution of (~5 mL)

f. 500 L of freshly prepared ice-cold NaBH4+ growth solution of (~5 mL)

2. Collect the absorption spectrum of each sample.

3. Plot position of LSP vs conc. of NaBH4

Results

1. Beer lamberts law is verified.

2. The absorption coefficient of the gold nanoparticles is ~3.36 L.mol-1

.cm-1

.

3. The effect of NaBH4is plotted.

Reference

1. J. P. Juste, I. P. Santos, L. M. Liz-Marzan, P. Mulvaney. Coord. Chem. Rev., vol. 249,

2005, 18701901.

2. C. J. Murphy, T. K. Sau, A. M. Gole, C. J. Orendorff, J. Gao, L. Gou, S. E. Hunyadi, T.

Li, J. Physical ChemistryB, vol.109, 2005, 13857-13870.

3.

T. S. Sreeprasad, A. K. Samal, T. Pradeep, Langmuir, Vol. 23, 2007, 9463-9471.4. A. K. Samal, T. S. Sreeprasad, T. Pradeep, J. Nanopart. Reserarch, vol.12, 2010, 1777-

1786.


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11. ESTIMATION OF FREE ENERGY OF PROTEIN DENATURATION

USING THE INTRINSIC FLUORESCENCE OF THE PROTIEN

11.1 Aim

To determination of free energy of protein denaturation using the intrinsic fluorescence of the

protein (bovine serum albumin / human serum albumin / lysozyme).

11.2 Chemicals Required

Protein (bovine serum albumin), urea, guanidine hydrochloride, citrate buffer and HCl tris buffer


5 mL, 10 mL, 25 mL, 50 mL, 100 mL standard flasks, 0.1 mL, 1 mL, 5 mL, 10 mL pipettes, 25

mL, 50 mL beakers, fluorimeter and pH meter.

11.4 Principle

Denaturation is a process in which proteins lose their secondary structure and/or tertiary structure

by application of external stress like heat or chemicals such as urea, guanidine hydrochloride.

Proteins have an intrinsic fluorescence due to tryptophan, tyrosine, and/or phenylalanine. Mostof the emissions are due to excitation of tryptophan residues, with a few emissions due to

tyrosine and phenylalanine. The table summarizes the fluorescence characteristics of the three

aromatic residues:

Lifetime Absorption Fluorescence

(nm) (M-1

cm-1

) (nm) Quantum

Tryptophan 2.6 280 5,600 348 0.20

Tyrosine 3.6 274 1,400 303 0.14

Phenylalanine 6.4 257 200 282 . 0.04

The fluorescence of the aromatic residues varies in somewhat unpredictable manner in variousproteins. Compared to the unfolded state, the quantum yield may be either increased or decreased

by the folding. Accordingly, a folded protein can have either greater or lesser fluorescence thanthe unfolded form. Fluorescence intensity, wavelength of the emitted light or a combination of

both can serve as a probe for studying perturbations of the folded state or denaturation.

Tryptophan residues that are exposed to water, have maximal fluorescence at a wavelength ofabout 340-350 nm, whereas totally buried residues fluoresce at about 330 nm.

An assumption made in this experiment is that at any concentration of the denaturant, lysozyme

will exist in only two states, the native (N) and completely unfolded or denatured (D) states.These two states are assumed to be in equilibrium, as shown in Scheme I.

From the measured parameters and their calculated baseline and maximum values, the fraction ofprotein in the denatured state at any denaturant concentration may be calculated, assuming a two

state model, shown in the following equation Yis the measured parameter (fluorescence intensity, I), Yois the baseline value of the parameter

at low denaturant concentration, Ymaxthe maximum value of the parameter when the protein is


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26

completely unfolded. The apparent equilibrium constant,KD, for the denaturation process may

then be calculated as shown in following equation

FromKD, the free energy change may be calculated from GD= -RT lnKD.

From these G values, the G value at zero denaturant concentration may be obtained using thelinear extrapolation method.

11.5 Procedure

1. Buffer 1: Prepare 0.050 M citrate buffer, pH = 3.

2. Solution 1: Prepare 25 mL solution of 10 M urea in buffer 1.

3. Stock 1: Prepare 5 mL, 0.12 mM solution of the protein in buffer 1.

4. Sample set 1: Prepare a series of solutions containing 3M protein each and varyingconcentrations of urea (3.0 9.0 M) from stock 1. (Nt: The final volume of each sample

should be 4.0 mL).

5. The solutions were allowed to sit for at least 1 h

6.

Measure the fluorescence intensity. (ex= 280 nm, em= 300-460 nm).

7. Calculate the GD from the equations given above.8.

S. No [Urea]

in M

Vol. of Urea sol.

in mL

Vol. Protein

in mL

Buffer

in mL

1 0

2 3

3 4

4 5

5 6

6 7

7 8

8 9

Result

The value of free energy of protein denaturation (GD) is 17-21kJM-1

(literature value)

Reference

1. R. F. Greene, Jr. and C. N. Pace, Journal of Biological Chemistry, vol. 249, 1974, 5388

5393.


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27

12. FLUORESCENCE QUENCHING OF FLUORESCEIN BY

INORGANIC ANIONS: DETERMINATION OF QUENCHING

CONSTANTS, ESTIMATION OF FREE ENERGY CHANGE AND

ACTIVATION ENERGY FOR THE QUENCHING PROCESS

12.1 Aim

To determine the fluorescence quenching constant and calculate the activation energy as well as

free energy change of the quenching process in fluorescein by selected anions.


Fluorescein,Quenchers: NaBr, NaIand Na2SO4

12.3 Appartus required

5 mL, 10 mL, 25 mL standars flasks, 5 mL, 10 mL pipettes, 1 mL graduated pipette, 25 mL, 50mL beakers

12.4 Principle

The fluorescence quenching constant from fluorescein was analyzed using Stern-Volmer

kinetics. []

According to equation 1, a plot of I0/I vs[Q] will be a straight line with KSVas the slope, where

I0and I are the intensity of emission in the absence and presence of Q, respectively.

KSV=kq, where kq is the quenching constant and is the excited state lifetime of the fluorophore.

The free energy change in the electron transfer process can be calculated from Treinin and

Hayon equation 2. Where, Ecttsis the charge transfer to solvent transition energy of the quenchers, E 1/2is the redox

potential and ESis the singlet transition state energy for fluorescein. All the units are in electron

volt.

A plot of logKqand G will show a linear plot, which suggests that electron transfer is the rate

determining step in the process. The equation for the linear plot will be given by equation 3. Since kETis the rate determining step, equation 3 can be re-written as

From Arrhenius equation, Where A is the collisional frequency for which a value of 10

10can be assumed. Substituting the

value of kETfrom equation 5 in equation 4, and assuming that kdiff~ k-diff, From equations 4 and 6, the following expression can be obtained for room temperature

experiments. 12.5 Procedure

1. The quenchers are NaBr, NaIand Na2SO4. A stock solution 25 mL of fluorescein (2 x 10-6

M) in water was prepared.


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28

2. Six separate solutions of the fluorescein ex= 484, (10-6

M) at identical concentrations were

prepared containing increasing amounts of quencher concentrations (0, 0.5 M).(Nt: The final

volume of each sample should be 10 mL).3. The emissions intensity at 520 nm from each solution is measured and each I0/I values was

plotted against [Q] to determine the fluorescence quenching constants with all the quenchers

given above.4. Using the calculated G values, the corresponding activation energy is calculated usingequation7.

5. Plot IO/I= Ratio of fluorescence intensity without or with quencher vs [Q]= Quencher

concentration

S.No [Q]

in M

Vol. of NaI

(mL)

Vol. fluorescein

(mL)

Water

(mL)

1 0

2 0.1

3 0.24 0.3

5 0.4

6 0.5

Results

1. The fluorescence quenching constant in fluoroscein derivatives by selected anions is

2. Free Energy Change (G) for the Quenching Process is -50 eV

3. Activation Energy (Ea) for the Quenching Process is 1.125 eVmol-1

References

1. P.K. Behera, T. Mukherjee and A.K. Mishra, Indian Journal of Chemistry:A, vol.34(A),

1995, 419-422.

2. Lakowiz J. R. Principles of fluorescence Spectroscopy, 1999.

Io

/I

Conc. Of [KI] (M)


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13. CRITICAL MICELLE CONCENTRATION OF CTAB

13.1 Aim

To study the aggregation behaviour and determination of the critical micelle formation

concentration of cationic surfactant in water by surface tension measurement.

13.2 Chemical required

Hexadecyltrimethylammonium bromide(CTAB)

13.3 Glassware required

Beaker (50 mL), Tubing with a balloon, Weighing bottle

13.4 Principle

The surfactant concentration at which micelle formation begins is known as the critical micelleformation concentration (CMC). Micelles are spherical or ellipsoid structures on whose surface

the hydrophilic heads of the surfactant molecules are gathered together whereas the hydrophobic

tails project inwards.

The critical micelle formation concentration (CMC) can be determined by carrying out surfacetension measurements on a series of different surfactant concentrations. If the liquid with known

surface tension is used for one of them (water), the surface tension of the other liquid can be

calculated from the equation: Surfactants exhibit a specific surface tension curve as a function of the concentration. Initially

the surfactant molecules increasingly enrich themselves at the water surface. During this phase

the surface tension decreases linearly with the logarithm of the surfactant concentration. When

the CMC is reached, i.e. when the surface is saturated with surfactant molecules, a furtherincrease in surfactant concentration no longer has any appreciable influence on the surface

tension.Determination of the critical micelle formation concentration

This means that in order to determine the CMC the two linear sections formed by the measuring

points obtained from the series of different concentrations must be determined. The CMC isobtained from the intersection of the straight lines for the linear concentration-dependent section


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30

and the concentration-independent section. The surface excess concentration of surfactant ions

s and the area per molecule was calculated from the slope of the straight line in the surface

tension plot (d/d ln C) below CMC, using appropriate form of Gibbs adsorption equation:

13.5 Procedure

1. Mount the clean and dry stalagmometer on the vertical stand.2. Weigh the mass of the weighing bottle m0.

3. Fill the beaker with distilled water. Mount the tubing with balloon on the top end ofstalagmometer. Immerse the bottom end of stalagmometer into water and fill it up, such that

the water level is above the wide part of stalagmometer.

4. Remove the balloon and collect 20 water drops into the weighing bottle.

5. Weigh the mass of the weighing bottle with water and determine the mass of 20 drops.6. Empty the weighing bottle and stalagmometer, dry them and prepare for the next

measurement.

7.

Repeat steps 2-6 for liquids with the unknown surface tension.

8. Knowing the temperature in laboratory, determine the water surface tension using valuesfrom the table 1, and calculate the surface tensions of studied liquids according to the

equation.

Table: The temperature dependence of the surface tension of distilled water

TempoC

20 21 22 23 24 25 30

Surface tensionN/m

0.07275 0.07259 0.07244 0.07228 0.07213 0.07197 0.07118

Result

1. The critical micelle concentration of CTAB is

References

1. A. Domnguez, A. Fernndez, N. Gonzlez, E. Iglesias and L. Montenegro, Journal of

Chemical Education. Vol. 74, 1997, 1227-1231.

2. Peter Atkins, Physical Chemistry, 9th

Ed., W. H. Freeman and Co., (2009)


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14. ACIDITY DETERMINATION IN ZEOLITES BY AMMONIA

TEMPERATURE PROGRAMMED DESORPTION

14.1 Aim

To analyse various zeolite samples for its qualitative and quantitative determination of acidic

sites by Temperature Programmed Desorption of ammonia

14.2 Chemicals required: ZSM-5 Zeolite

14.3 Apparatus required: TPD

14.4 Principle

Zeolites are microporous aluminosilicate structures, important for a variety of applicationsincluding selective reactions and adsorptive separations. It is important to estimate the acidic

(both Brnsted or Lewis) sites present in zeolites. Temperature-programmed desorption (also

called thermal desorption spectroscopy, TDS) of ammonia is the most widely used technique for

characterizing acidity in zeolites. Using this method the number of acid sites can be determined

from the amount of desorbed ammonia molecules adsorbed directly on the acid sites. TPDexperiment typically involves saturation of the surface with ammonia under a set of adsorption

conditions, followed by linear ramping of the temperature of the sample in a flowing inert gas

stream. Ammonia concentration in the effluent gas can be measured by absorption/titration ormass spectroscopy. The experiment can also be carried out in a microbalance and changes in

sample mass can be recorded continuously as a function of temperature. The amount of ammonia

desorbing above some characteristic temperature is taken as the acid-site concentration, and thepeak desorption temperatures can be used to calculate heats of adsorption. This method provides

information on the number of acid sites and their strength.

14.5 Procedure

1. 100 mg of dry zeolite sample pre-treated at 500 C in 25 mL/min Pure Helium gas flow for

30 min.2. Temperature cooled to 100 C and then 10% NH3in Helium gas mixture (25 mL/min) is

passed through the sample for 30 min.

3. Physisorbed Ammonia gas is removed by flushing with Helium gas for 15 minutes.4. Temperature programmed heating started from 100 C to 500 C @ 10 C per min and NH3

evolution recorded.

5. The peak area will give the volume of NH3desorbed which is directly proportional to the

acidity at that peak temperature.

S. No. Temp.C

TCD Signal(a.u.)

0 100 200 300 400 500 600

0.124

0.128

0.132

0.136

0.140

TCDs

ignal(a.u.)

Temperature (oC)

Documents

Msc Pgp Modified April 10