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MIDLANDS STATE UNIVERSITY
DEPARTMENT OF CHEMICAL TECHNOLOGY
PHYSICAL CHEMISTRY (CT201) PRACTICALS
Compiled by Mambanda Isaac (Chief Technician) & Gonzo Muriel (Snr Lab Technician)
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EXPERIMENT 1: CRITICALMICELLE CONCENTRATION
INTRODUCTION
Ionic surface active materials, such as sodium dodecylsulphate (SDS) consist of two species of ion, an ionic head
group attached to a hydrocarbon tail (dodecyl-sulphate ion) and a counter ion (sodium ion).The way in which the
physical properties of a solution of a typical colloidal electrolyte, such as SDS, vary with increasing concentration,
alters when a particular concentration is reached. This concentration is known as the critical micelle concentration,
or c.m.c .The accepted explanation of this is that in the region of the c.m.c., aggregation of the long chain
electrolyte ions into quite large, charged units begins to occur. These charged units are typically made out up of
about 100 surface active ions forming a spherical unit, with the charges situated at the periphery and the
hydrocarbon tails in the centre .The adsorption of a considerable number of counter ions reduces the overall
charge and also the coulombic repulsion between the head groups.
APPARATUS
Stock solution (2x10-2 moldm-3) of sodium Dodecylsulphate glassware:
Either (a) surface tension (torsion balance)
Or
(b) conductance bridge,
METHOD
Prepare a series of solutions from the stock solution provided over the concentration range 10 -3 moldm-3 to
2x10-3moldm-3.To measure the c.m.c either measure the surface tension (γ) or the conductance (G) .If conductance
is to be measured also measure the conductance of the water used to prepare the dilutions. Subtract the
conductance of the water from the conductance of each of the test solutions.
RESULTS
(a)
Plot a graph of γ against c
γ
C.m.c
ℓ
(b) Conductivity= K= Cell Constant x C
A= molar conductivity = K/C
= cell constant x G/C
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Hence in (b) it is sufficient to plot G/C against √ c
G/C
c.m.c
√c
Estimate the c.m.c using your results and give a future for the likely experimental error.
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EXPERIMENT 2: THE SOLUBILITY OF A SOLID SOLUTE IN AN IDEAL SOLUTION.
Aim
To construct the solubility curve of a near ideal system.
Theory
For a solid solute saturated solution, the equilibrium constant K, is just the mole fraction of solute in the saturated
solution. Hence the variation of with temperature is given by the van Hoff`s equation;
In
x2
=
(N2)
-∆Ho
+
constant
RT
Where x2 is the solute mole fraction. Now ∆Ho can be split up into the sum of the enthalpy changes for two
processes:
Solid solute
liquid solute
saturated solution
∆Ho for the first step is ∆Ho for the second step should be zero if the system is ideal, because in the mixing of two
ideal liquids there should be no energy changes.
Therefore, In x2 = -∆Hofusion/RT + constant
When x2 = 1, i.e. pure solute, T is the melting point of pure solute, To
Therefore, In x2 = -∆Hfo (1/T – 1/To)
(1)
R
In this experiment the system; naphthalene in toluene is to be investigated and the data compared with the
theoretical results assuming ideal behavior.
Procedure
Weigh between 14, 9 and 15,1g of naphthalene flakes into a boiling tube and add 6,00ml of toluene from a
burette. Immerse the boiling tube in a beaker of hot water (70oC; note that the boiling point of toluene is 110 oC)
and stir with a wire loop stirrer until the solid has dissolved. Place a thermometer in the tube, remove the tube
from the bath and allow it to cool in air, stirring fairly vigorously to prevent super cooling. Note the temperature at
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which crystals first appear in the solution. Add a further 2,00ml of toluene to the test tube and repeat the above
sequence. Continue the process until eight crystallization temperatures have been recorded.
CAUTION- owing to the toxic nature of toluene, this experiment should be performed in the fume hood, and care
taken that no vapour is inhaled or any liquid spilt on the skin.
Calculations
1. Calculate x2 for each solution taking the density of toluene at room temperatures as 0,864g/ml.
2. Plot In x2 versus 1/T (K).
3. Draw on the same graph the straight line of theoretical slope -∆Hf/R
For naphthalene, ∆Hf = 19,079J/mol. This line will pass through the point corresponding to the melting of pure
naphthalene (x2 = 1) which is 80,2oC.
4. (a) Obtain ∆Hf and Tf for naphthalene from your experimental curve and
(b) Comment on the extent to which the naphthalene/toluene system approached ideal behavior and, using
your knowledge of the molecular structure of these compounds, where the results could have been expected.
Questions
How would such data be used in the construction of a phase diagram for a two component system?
*∆Hf = ∆Hfusion ; Tf = Melting point of naphthalene.
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EXPERIMENT 3: THE THERMODYNAMICS OF ELECTROCHEMICAL CELLS
Objective
To measure the e.m.f of an electrochemical cell at a variety of temperatures and hence determine the
thermodynamic functions ∆G®, ∆S®, and ∆H® for the cell.
Introduction
It is readily shown from thermodynamics that
E® = -∆G®/nF
and (δE®/δT)p = ∆S®/nF
Consequently electrochemical potentials and their temperature variation provide an invaluable experimental
method of determining ∆G® and ∆S®, and hence K and ∆H®.
The cell used here is a Gerke Cell :
Ag(s) | AgCl(s) | Cl- (aq) | Hg2Cl2(s) | Hg (l)
which has two distinct advantages: there is no liquid junction between the electrodes , and the e.m.f. does not
depend on any potentially varying activity ,but is always equal to E®. The cell reaction is :
2Ag(s) + Hg2Cl2
2Hg(l) + 2AgCl(s)
Note that the sign of the e.m.f for the cell as written is the polarity of the mercury electrode.
Method
Set up an ice-water bath into which the cell is placed , so that the bath liquid is above the liquid inside the cell
.Leave to cool to about 0°C , adding more ice as needed: it may take one hour or longer before the cell is at the
bath temperature . Stir the cooling bath from time to time .
Once the experiment has equilibrated , take three readings of the cell potential at one minute intervals, noting the
polarity of the cell and temperature .As the sign displayed by the voltmeter belongs to the electrode connected to
the Red/High/V terminal (sometimes confusingly labelled +), this enables the polarity to be determined .Mix some
of the ice-cold water with tap water to achieve a temperature a little above 10°C and re-immerse the cell in the
bath . Leave for about 20minutes , stirring from time to time. Carry out the determination of temperature , cell
potential and polarity as before .Repeat the procedure at temperatures of about 20° ,30° ,40° and 50°C , mixing
colder and hotter water to achieve the required temperatures , and allowing about 20 minutes for equilibration at
each temperature (add more hot water if necessary).
Note: only connect the voltmeter when taking the readings: disconnect when equilibrating.
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Results and discussion
1. Deduce the electrode reactions and hence the cell reaction and it`s Nernst equation, proving that invariably E =
E®.
2. Plot a graph of e.m.f with temperature and determine the slope and value of E at 25oC. Calculate ∆G® and ∆S®
and hence ∆H® at 25oC. Compare these with values calculated from thermodynamic tables.
3. Deduce the direction of spontaneous change at 25oC.
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EXPERIMENT 4: MEASUREMENT OF THE SURFACE TENSION OF A LIQUID BY THE BUBBLE
PRESSURE METHOD.
Theory:
When a bubble is formed slowly at the end of a capillary tube, radius r, which dips into a liquid to a depth d, the
pressure in the tube increases to a maximum which corresponds with a hemispherical bubble and then decreases
as the bubbles break away.
The maximum bubble pressure is given by
P = dℓg + 2γ/r
(1)
Where ρ is the density of the liquid to be measured, g is the acceleration due to gravity, and γ is the surface
tension of the liquid. Finally P is the excess pressure.
Method: the excess bubble pressure is measured using the apparatus shown in Figure 1.
Step 1; The tap should be open to the atmosphere. Open the clamp sufficiently to allow the rubber bulb to be filled
with air through the side valve. The liquid in the two tubes in the manometer should now be equal in height. Note
the manometric liquid is dibutyl phthalate which has a very small vapour pressure; its density is 1,046 g/cm 3 at
20oC.
Now, fill the clean sample tube to 3/4 with double-distilled water, and place on the lab jack under the capillary
tube. The sample tube should be centrally placed. Raise the lab jack until the capillary just touches the surface of
the water (h = 0,00 cm).
Step 2; Close the tap, and open the main valve on the bulb and tighten the clamp sufficiently tight so that a bubble
forms at the end of the liquid. Measure the difference in height of the liquid (dibutyl phthalate) in the two
manometer arms. The excess pressure is given by
P = ∆hℓ®g
(2)
Where ∆h is the difference in height ℓ®is the density of dibutyl phthalate.
Step 3; Open the tap, and loosen the clamp and refill the bulb by opening the side valve. Now raise the lab jack so
that the capillary dips into the liquid to a depth d. Use the capillary scale to determine this depth. (Note the scale is
not metric and must be calibrated using a metric metal rule and the conversion factor obtained at the end of the
experiment.)
Now repeat steps 1 and 2.
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At least six measurements of d and h should be obtained.
Treatment of results
Use double-distilled water as the first liquid. From equations (1) and (2) we obtain
∆hℓ®g = dℓg + 2γ/r
or
∆h = dℓ/ℓ®
+ 2γ/ ℓℓ®g
(3)
Using (3), plot ∆h (y-axis) against d (x-axis ). The gradient as (ℓ/ℓ®) and the intercept on the ∆h axis is (2γ/rℓ®g).
Substitute for the γ value for water to obtain the constant (2/rℓ®g).
Repeat the method using the two unknown liquids. The constant, 2/rℓ®g obtained for water can now be used to
obtain γ from the intercept.
Precautions: 1. Ensure that the capillary is clean by placing the tip of it in a small sample bottle of acetone.
2. When changing the liquid ensure that the capillary tube tip is washed with liquid prior to measurement. Discard
this liquid and replace with fresh sample.
3. On no account should you allow liquid to be drawn up the capillary tube. This can be avoided by following
instructions carefully.
4. Do not touch the tip of the capillary tube as this will contaminate the tube with grease and affect your results.
Compiled by Mambanda Isaac (Chief Technician) & Gonzo Muriel (Snr Lab Technician)
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EXPERIMENT 5: THE EMF`S AND USES OF CONCENTRATION CELLS AND THE
THERMODYNAMICS OF ELECTROCHEMICAL CELLS.
Objective:
To construct various concentration cells and measure their e.m.f.`s, and to deduce activity coefficients and
equilibrium constants from them.
Introduction
When two electrodes of the same chemical nature are combined to form a cell, an e.m.f. is produced if there is
some difference in the activity between two electrodes. Such a cell is called a concentration cell, and the
difference in activity may be brought about by having similar electrodes in contact with solutions of different ionic
activities.
As with any other electrochemical cell, the e.m.f. of the cell will be equal to the difference between the potentials
of the individual electrodes. The two electrodes may be conveniently connected by means of a salt bridge, which
usually consists of a glass tube containing a concentrated solution of a univalent electrolyte in which the transport
numbers of the cat ion and anion are very nearly equal. This has the consequence that the liquid junction potential
between the electrodes is approximately zero. In this experiment, the following typical concentration cell is
investigated (a1 anda2 are activities, and indicates a salt bridge):
Ag
Ag+, a1
Ag+, a2
Ag
; a1 = γ1 c1 and a2 = γ2 c2
The equilibrium at a silver electrode is
Ag+(aq) + e-
Ag(s)
And the potential, E2, of the right hand electrode in the above cell will be given by the Nernst equation as
E2 = EAg Θ + RT/F x In a2
(1)
Where a2 is the activity of Ag+ ions in the electrode compartment.
The e.m.f, E, of the concentration cell is given by E = E 2 – E1, E being positive or negative according to the
experimentally determined sign of the right-hand electrode, E2. Using equation (1),
E = RT/F x In a2/a1
(2)
If the activity in one electrode compartment is known then that in the other can be found. The second activity can
be controlled by a chemical equilibrium, hence enabling the equilibrium activities and constant to be found.
Also investigated here is the cell;
Ag Ag+; aAg+
Cl-, aCl-
AgCl Ag
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where the right-hand electrode is a AgCl/Ag electrode, which is a reversible Cl- electrode:
AgCl(s) + e-
Ag(s) + Cl-
Consequently the overall cell reaction is
AgCl(s)
Ag+(aq) + Cl-(aq)
and its Nernst equation is
E = EΘ - (RT/F)In(aAg+)(aCl-)
At equilibrium, E = 0, so
EΘ = (RT/F)InKsp
The measurement of cell e.m.f`s has a number of complications. Firstly the potential should be measured when
there is no current flowing: ideally a potentiometer is used, but an electronic voltmeter is also good. Secondly
there should be a state of equilibrium at the electrodes, and thirdly the electrode should be reversible – the rates
of the forwards and backwards electrode reactions should be high. Clean metal/ion electrodes can achieve the last
two conditions provided the ion activities are not determined by slow reactions.
Method
Prepare accurately using graduated flasks 0,005 and 0,001 mol/L silver nitrate solutions from the stock solution.
Next prepare the silver electrodes: rub gently with emery paper, dip them briefly in dilute nitric acid and then rinse
them with distilled water and store them in distilled water. Put them into a beaker containing some 0,01mol/L
silver nitrate solution and measure and record the cell potential and polarity. If it exceeds 0,005V obtain
replacement electrodes and check them.
Set up each of the following cells, rinsing silver electrodes with a little of their respective solutions first. Measure
their e.m.f`s and polarities: allow two minutes for equilibration and take three readings at one minute intervals. As
the sign displayed by the voltmeter belongs to the electrode connected to the Red/High/V terminal (sometimes
confusingly labeled +), this enables the polarity to be determined. Record the temperatures of the left-hand
electrodes.
1. Ag Ag+, 0,01M Ag+, a2 Ag
with a2 for 0,005 and 0,001M solutions
2. Ag Ag+, 0,01M Cl-, 0,01M AgCl Ag. (Obtain the AgCl/Ag electrode from the technicians)
3. Ag Ag+, 0,01M Ag+, a2 Ag
one drop of AgNO3.
within the right-hand electrode compartment 0,01M KCl, to which is added
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4. Ag Ag+, 0,01M Ag+, a2 Ag
within the right –hand electrode compartment a mixture of 25 ml of each
of 0,01M AgNO3 and 1,0 M NH3 solutions. Standardize the NH3 solution by diluting 25 ml to 250ml and titrating 25
ml aliquots with 0,1 M HCl solution using screened methyl orange.
Treatment of results
Where appropriate, correct the measured potentials for the zero-error of the silver electrodes, taking care to
account for the the polarities correctly. Use 0,892 and 0,902 as the activity coefficients of 0,01M Ag + and Clrespectively.
1) For each of the cell 1, deduce the ratio of the activities and of the activity coefficients for the Ag+ solutions, and
hence calculate the activity coefficients for 0,005 and 0,001 M solutions.
2) For cell 2 deduce EΘ, and hence calculate Ksp for AgCl. Why is only one drop of silver nitrate solution required in
the right-hand compartment?
3) For cell 3 deduce a2 and hence Ksp for AgCl. By expressing E for each of the cells 2 and 3 in terms of Ksp.
4) For cell 4 deduce a2 and hence the equilibrium concentrations of the species and K for the formation constant
of the complex Ag (NH3)2+. Take the activity coefficient of the complex ion to be the same as that for the same
concentration of Ag+ about 0, 92.
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EXPERIMENT 6: STUDY OF POTENTIOMETRIC AND INDICATOR END-POINTS
Discussion
The objectives of this experiment are to correlate electrometric and indicator end-points, to distinguish between
equivalence points and end-points and to investigate any indicator blanks and hence the choice of indicator.
The titration selected is that of phosphoric acid and sodium hydroxide. Phosphoric acid is a tribasic acid with three
successive ionizations:
H3PO4 = H2PO-4 + H+
H2PO4- = HPO42- + H+
HPO42- = PO43- + H+
In this experiment three titrations of phosphoric acid with sodium hydroxide are carried out using three different
indicators. A fourth titration is performed, when the pH is recorded using a pH meter and a suitable glass and
calomel electrode assembly.
Apparatus and chemicals
Bench model pH meter, glass electrode, dip-type calomel electrode, 0, 2 mol/dm3 phosphoric acid, 0, 5 mol/dm3
sodium hydroxide, methyl red, phenolphthalein and thymolphthalein indicators.
Method
10 cm3 of the phosphoric acid and 50 cm3 of distilled water are placed in a 250 cm3 conical flask and titrated with
sodium hydroxide to the methyl red end point. The titration is repeated using phenolphthalein and
thymolphthalein. Again 10 cm3 of the phosphoric acid are placed in a 250cm3 beaker and sufficient distilled water
added so that the glass and calomel electrodes are just covered. The electrodes are then connected to the pH
meter. A preliminary titration is then performed to determine the region where the pH changes most rapidly. In
subsequent titrations more frequent observations may then be made in these regions. The solution should be
stirred after each addition of alkali. Upon the completion of the titrations, the electrode system is immersed in
distilled water.
The results obtained are recorded in tabular form under the headings:
Compiled by Mambanda Isaac (Chief Technician) & Gonzo Muriel (Snr Lab Technician)