UN
NESCO-NIG
GERIA TEC
CHNICAL &
VOCATIO
ONAL EDUC
CATION
REVIITALISATIION PROJE
ECT-PHASE II
NATIIONAL
L DIPLO
OMA IN
I
SCIENC
S
CE LAB
BORATORY TECH
HNOLO
OGY
Phyysical chem
mistrry
COURSE CODE:
C
STC 1222
YE
EAR I- SE
S MES
STER III
EXPE
ERIMEN
NTS
V
Version
1:: Decembeer 2008
Deceember 20008
1
TABLE OF CONTENTS
WEEK 1……………………………………………………………………..3
WEEK 2……………………………………………………………………..5
WEEK 3……………………………………………………………………..7
WEEK 4……………………………………………………………………..13
WEEK 5……………………………………………………………………..16
WEEK 6……………………………………………………………………..18
WEEK 7……………………………………………………………………..21
WEEK 8……………………………………………………………………..23
WEEK 9……………………………………………………………………..25
WEEK 10……………………………………………………………………27
WEEK 11……………………………………………………………………30
WEEK 12……………………………………………………………………32
WEEK 13……………………………………………………………………34
WEEK 14……………………………………………………………………36
WEEK 15……………………………………………………………………38
2
WEEK
K 1. EXPE
ERIMENT
T 1.
The Effeect of tempeerature on rate
r
of reacttion
AIM: Too investigatee the effect of temperaturre on the ratee of reactionn between soodium
thiosulphhate and hyd
drochloric acid.
SAFETY
Y: Wear eyee protection. Take care not to inhale fumes
Materialls:
• 250 cm3 conical flasks (xx2)
T
r
• Thermometer
• Bunsen
B
burneer or hot platte
• trripod stand
• 100 cm3meassuring cylindder
uring cylindeer
• 10 cm3 measu
• White
W
paper with
w a cross written by a pen (black or blue)
• Stop clock
• Distilled
D
wateer
• 1 mol dm-3 so
olution of sodium thiosuulphate
-3
• 2 mol dm so
olution of hyydrochloric acid
a
T RECOR
RD:
WHAT TO
Record your
y
resultss in the table
Initial
Final
F
temperature
t
e of
temperaature of
the mixture in
the
t mixture in
k
°C
the
t flask °C
the flask
Averrage
temp
perature of
the mixture
m
in
the flask
f
°C
3
ken for
Time tak
the crosss to
disappeaar (sec)
(sec-1)
PROCEDURE:
1. Put 10 cm3 of 1 mol dm-3 sodium thiosulphate solution and 40 cm-3 of distilled water into
a conical flask.
2. Measure 5 cm3 of 2 mol dm-3 hydrochloric acid solution in 10 cm-3 measuring cylinder.
3. Warm the thiosulphate solution in the flask if necessary to bring it to the required
temperature. The object is to repeat the experiment five times with temperatures in
the range 15-65 °C.
4. Put the conical flask over a piece of paper with a cross drawn on it.
5. Add the acid and start the clock. Swirl the flask to mix the solutions and place it on a
piece of white paper marked with a cross. Record the initial temperature of the mixture.
6. Look down at the cross from above. When the cross disappears, stop the clock and record
the time taken. Record the final temperature of the mixture in the flask.
7. As soon as possible, pour the solution down the sink (in the fume cupboard if possible)
and wash away.
Questions
1. For each set of results, calculate the value of 1/time. (This value can be taken as a
measure of the rate of reaction for this experiment).
2. Plot a graph of 1/time on the vertical (y) axis and average temperature on the
horizontal (x) axis.
3. From your graph say: (i) how rate varies with temperature
(ii) describe the shape of the curve
4
WEEK
K 2, EXPER
RIMENT 2.
The effecct of concen
ntration on rate of reacction
AIM: Too investigatee the effect of concentration of sodiuum thiosulphate on the raate of reactioon
between sodium thiosulphate andd hydrochlorric acid.
SAFETY
Y: Wear eyee protection. Take care not to inhale fumes
Materialls:
• 250 cm3 conical flasks (xx2)
T
r
• Thermometer
• Bunsen
B
burneer or hot platte
• trripod stand
• 100 cm3meassuring cylindder
uring cylindeer
• 10 cm3 measu
• White
W
paper with
w a cross written by a pen (black or blue)
• Stop clock
• Distilled
D
wateer
• 1 mol dm-3 so
olution of sodium thiosuulphate
-3
• 2 mol dm so
olution of hyydrochloric acid
a
T RECOR
RD:
WHAT TO
Complette the table
Volume of
Volume
V
of
sodium
water
w
(cm3)
thiosulph
hate
3
(cm )
50
40
30
20
10
0
10
0
20
0
30
0
40
0
Time taaken for O
Original
cross too
c
concentratio
on
disappeear (sec) of
o sodium
th
hiosulphatee
(ggdm-3)
5
50
5
PROCEDURE:
1. Put 50 cm3 of 1 mol dm-3 sodium thiosulphate solution in the conical flask.
2. Measure 5 cm3 of 2 mol dm-3 hydrochloric acid in a 10 cm-3 measuring cylinder.
3. Add the acid to the flask and immediately start the clock. Swirl the flask to mix the solutions
and place it on a piece of paper marked with a cross.
4. Look down at the cross from above. When the cross disappears stop the clock and
record the time in the table.
5. Repeat the experiment using different concentrations of sodium thiosulphate solution. Make
up 50 cm3 of each solution. Mix different volumes of the sodium thiosulphate solution
with water as shown in the table.
6. As soon as possible, pour the solution down the sink (in the fume cupboard if
possible) and wash away.
QUESTIONS:
1. Calculate the concentration of sodium thiosulphate in the flask at the start of each
experiment. Record the results in the table.
2. For each set of results, calculate the value of 1/time. (This value can be taken as a
measure of the rate of reaction).
3. Plot a graph of 1/time taken on the vertical (y) axis and concentration on the
horizontal (x) axis.
4. From your graph say: (i) how rate varies with concentration
(ii) describe the shape of the curve
6
WEEK 3, EXPERIMENT 3
The iodine clock reaction
AIM: To determine the rate constant and order of reaction by iodine clock reaction
SAFETY: Wear eye protection. Take care not to inhale fumes
MATERIALS:
• 0.005mol dm-3 Sodium thiosulphate
• 0.2mol dm-3 potassium iodide
• 0.1mol dm-3 ammonium potassium peroxodisulphate (vi)
• 1% Starch solution
• Thermometer
• Stop clock
• Test tubes
• 100cm3 Burettes
• 100cm3graduated cylinder (x3)
• 10cm3graduated cylinder (x3)
• 250- cm3 erlenmeyer flask
Procedure:
The Preparation Table below shows the reaction mixtures No.1-5 to be studied in order to
determine the rate constant and reaction orders for the reaction of iodide ion with peroxodisulfate
ion. The KI and (NH4)2S2O8 solutions should be carefully measured with graduated cylinders,
and Na2S2O3 solutions should be dispensed with a burette (or a 10.00 cm3 volumetric pipette);
the KCl and (NH4)2SO4 solutions may be measured out in graduated cylinders.
Table for the Preparation of Experimental Runs
Reaction
0.2mol
0.2mol
0.1mol dm-3
No.
dm-3
dm-3
(NH4)2S2O8
KI
KI
(cm3)
(cm3)
(cm3)
0.1mol
dm-3
(NH4)2SO4
(cm3)
0.005mol
dm-3
Na2S2O3
(cm3)
1
2
0
0
10
10
20
10
0
10
20
20
7
Temp. Time
t
(°C)
(s)
3
4
5
20
20
15
0
0
5
10
5
15
10
15
5
10
10
10
Procedure
Carefully measure out the KI into a 50-cm3 beaker (I) and the (NH4)2S2O8 into a similar beaker
(II) for reaction No.1. The 10.00 cm3 of Na2S2O3 solution and 3-4 drops of starch solution are
placed in a 250- cm3 erlenmeyer flask (along with the required volumes of KCl and (NH4)2SO4
solutions in reaction Nos. 2-5). These last two components are added merely to dilute each
reaction mixture to a total volume of 50.0 cm3 while keeping the ionic strength constant. Start the
reaction by simultaneously pouring the solution from beakers I and II into the erlenmeyer flask
and swirling thoroughly. At the instant the reactants are mixed, the stop watch should be started.
Stop timing when the first sudden appearance of the blue-black colour occurs. Record the time
required for the reaction. Also, record the final temperature of the reaction mixture. Repeat the
procedure above for each of the four other reaction mixtures.
Calculations:
For the rate law,
∆
∆
use the method of initial rates to evaluate and record both the reaction "orders" (m and n) and
evaluate and record the apparent rate constants (
the average value of
and show its units.
Determination of Rate
8
) from the data of Reactions No. 1-5. Report
Since the [
] original is the same in each of the reactions and since it is assumed to be
completely used up by the time ( ) the color change occurs, the following relationship can be
shown to be true:
Rate of loss of reactant,
Where
∆
.
∆
= time required for the clock to "tick". The value 5.00 x 10-3 mol dm-3 is the same for
each beaker, because it depends on the limiting reactant, the
.
Determination of Rate Law
If you have been careful in solution preparation, the rate law can be determined for your
reaction:
∆
∆
m and n are called the reaction orders for [I-] and
respectively.
Determination of m and n is made by inspection (more accurately by mathematics) using the
method of initial rates.
Method of Initial Rates
The method of initial rates measures the rate of a reaction just after the reactants are brought
together and allowed to interact. To determine the orders of the reaction, individual reactions are
compared to see how rate is affected by changing the initial concentrations of one reactant, while
holding the concentration of the other reactant constant. Consider the following simplified
possibilities:
If the [A]init. in Reaction No.1 is doubled in Reaction No.2 while [B]init. is the same in Reactions
No.1 and 2, the rate of the reaction may:
1. remain unchanged, if so the reaction order for [A] is 0
2. be doubled, if so the reaction order for [A] is 1
9
3. be quadrupled, if so the reaction order for [A] is 2
Thus, for the purpose of this experiment, the orders of reactant (m and n) can have values of 0, 1,
or 2.
EXAMPLE
Reaction No.
[A] moldm3
[B] moldm3
Rate of loss of A
(mols-1)
1
0.15
0.15
4.0 x 10-2
2
0.15
0.3
4.0 x 10-2
3
0.3
0.15
1.6 x 10-1
•
Between Reaction No.1 and Reaction No.2, [ A ] is constant while [ B ] is doubled (i.e.
0.15 mol dm-3
0.30 mol dm-3). The rate of the reaction is unchanged, so the order of
reaction for reactant B is n = 0.
•
Between Reaction No.1 and Reaction No.3, [ B ] is constant while [ A ] is doubled (i.e.
0.15 mol dm-3
0.30 mol dm-3 ). The rate of the reaction is quadrupled, so the order of
reaction for reactant A is m = 2.
The rate law is therefore: Rate of Loss of A = k[A]2 [B]0
or more simply Rate = k[A]2.
10
The rate constant (k) can be solved for by algebraic substitution into the equation: k = Rate / [A]2
For example: using the data from Reaction No.1 gives k = 9.0 x 10-4 M-1s -1. For a given reaction,
at a given temperature, k should be a constant that is independent of each of reactant
concentrations.
Caculation results
Reaction No.
[I]-init
Rate
Rate constant
(mol dm-3s-1)
k
1
2
3
4
5
Average value for k =
m=
n=
Rate Law:
Rate =
Questions:
1. What effect does doubling [ S2O8-2.] have on the rate? By what factor (experimental, not
rounded to a whole number) does the rate increase? Hint: Rate 2 / Rate 1 = ?
2. What effect does doubling [I.] have on the rate? By what factor (not rounded) does the
rate increase? Hint: Rate 3 / Rate 1 = ?
11
3. Why are the coupled reactions called an iodine clock reaction?
4. What interaction is responsible for the dark color?
5. Why does the dark color appear suddenly, rather than a gradual darkening?
6. Based on your experimentally determined rate law, predict a value for rate of reaction if
[I-]init. = [S2O82-]init = 0.50 M.
12
WEEK 4, EXPERIMENT 4
Hydrolysis of p-nitrophenylacetate
AIM: To determine pseudo first order rate constant and true rate constant
Safety: Wear eye protection and don’t inhale fumes.
Materials:
•
p-nitrophenylacetate
•
buffer solutions
•
UV-spectrophotometer
•
glasswares
Procedure:
We require a spectrophotometer and means of controlling and measuring hydroxyl ion
concentration. By observing the reaction at several different concentrations of hydroxyl ion, we
will be able to test the validity of the second order rate law which is thought to represent this
reaction. UV-Vis Spectrophotometer will be used to collect kinetic data, which will be obtained
by following the absorbance of the reaction solution at fixed wavelength. A pH meter, a buffer,
and dilute aqueous NaOH are used to measure and control pH; careful measurement of pH is
critical to the success of this project.
Preparation of solutions
Prepare a solution of 9 - 12 mg of p-nitrophenylacetate in 10 cm3 of acetone in a 10 cm3
volumetric flask; this solution must be prepared at the start of the laboratory period.
A small amount of dilute aqueous NaOH should be added to water to obtain the desired initial
pH, near 9.5. The stirred solution will be monitored continuously with a pH meter. A pH range
13
from about 9.5 to 11 will provide a suitable data set. Particular pH values are obtained by adding
small amounts of dilute NaOH to the solution being monitored with the pH meter. Be certain that
the pH readings are stable before starting any runs.
Turn on the UV and allow the instrument to warm up for at least two hours prior to the
experiment. Set the wavelength to 400 nm and zero the instrument with respect to the blank. Use
2.00 cm3 of the aqueous NaOH solution for a reference. Place the blank in the front
compartment, close the lid, and record the absorbance. Remove the cell containing the reference
solution.
To start the reaction: To prepare the reaction mixture, add 15 to 25 microliters of the ester
solution (via a Hamilton syringe) to 2.00 cm3 of aqueous NaOH in a spectrophotometric cell.
Cap the cell and mix by inverting several times. Place the cell in the front compartment of the
instrument, replace the cell compartment lid, and record the absorbance. This sequence should be
done with calm dispatch, to minimize the time between the start of the reaction and the start of
data collection. Remember, data are invalid if the absorbance exceeds 1.0.
For the next data set, adjust the NaOH solution to the desired pH by addition of dilute NaOH
solution, re-zero the UV, and repeat the absorbance/time measurement. At least four different pH
values are required for an adequate analysis.
For runs at or below pH=10, make measurements for 30 minutes. For runs above pH=10, make
runs for about 15 minutes. Set the stop time feature accordingly. You may stop a measurement if
the absorbance seems to be constant over a period of about 10 measurements. Be sure that the
14
absorbance does not exceed 1.0; if it does, redo the run, and use a little less of the ester solution
(see below).
Exercise
1. Using the equation ln
ln
,
ln
.
A straight line indicates first order reaction with respect to that reactant.
The slope = The intercept on the ordinate axis = ln
2.
Plot
A straight line indicates second order reaction with respect to that reactant.
The slope = k
The intercept on the ordinate axis =
3
From the two graphs, state the order of the reaction
15
WEEK
K 5, EXPER
RIMENT 5
Quantitaative electro
olysis
AIM: Too relate the amount
a
of metal removedd from an electrode to thhe electric cuurrent and thhe
time the current flow
ws.
Safety: Wear
W eye pro
otection.
APPARA
ATUS:
• Power supply
y
A
• Ammeter
• Weighing
W
ballance
• Beaker
B
• Copper
C
anodee
• Copper
C
catho
ode
-3
• 1 mol dm so
olution copper sulphate
E
paper
• Emery
What to record
T masses of
o the electroodes before electrolysis
e
(
(identify
thee electrodes by
b writing on
• The
thhem with a sharp
s
iron naail).
• The
T masses of
o the electroodes after eleectrolysis.
• The
T current flowing.
f
• The
T time the current flow
ws.
PROCED
DURE:
1. Clean the electrodes with emerry paper (avvoid inhalingg any dust).
2. Weighh the anode.
3. Immerrse the electrrodes to a deepth of 3–4 cm
c in 1 mol dm-3 solutioon of copper((II) sulfate.
4. Allow about 0.4 A to pass for about 30 miin.
5. Removve the anodee, wash careffully in wateer and dry geently with a tissue
t
paper..
6. Reweiigh the anodee.
Question
ns
1. Calcullate the numb
ber of moless of copper that
t have beeen removed from the
anode. (M
Mass lost in g / 63.5)
16
2. Calculate the charge that has flowed through the circuit using the relationship
charge (in Coulombs) = current (in amps) x time (in seconds).
3. Using the answers to questions 1 and 2, calculate the number of Coulombs
required to remove one mole of copper.
4. 193,000 (2 x 96,500) Coulombs is required to remove one mole of copper. The
difference between this and the answer to question 3 is due to errors in the
experiment. What are the main sources of error in this experiment?
17
WEEK 6, EXPERIMENT 6
Effect of catalyst on the rate of chemical reaction
AIM: To determine the effects of the enzyme catalase and Hydrogen peroxide concentrations on
the rate of decomposition of Hydrogen peroxide.
Safety: Wear eye protection.
Materials:
•
Weighing balance
•
50 cm3 conical flask
•
Splinter
•
Potato
•
Liver
•
Hydrogen peroxide
•
100cm-3 measuring cylinder
Introduction
Enzymes are biological catalysts, they increase the speed of a chemical reaction. They are large
protein molecules and these enzymes are very specific to certain reactions. Hydrogen peroxide
decomposes slowly in light to produce oxygen and water. There is an enzyme called catalase that
can speed up (catalyse) this reaction.
18
What to record
•
W you obsserve
What
•
W
Which
enzym
me source makes
m
the moost effective catalyst?
What to do
1.
Usinng a measuring cylinder, put 25 cm3 of hydrogenn peroxide soolution into a conical flaask.
2.
Add a small piecce of liver.
3.
Test the gas giveen off with a glowing splint.
4.
Disppose of this mixture,
m
inclluding the livver, into a buucket, and puut another 25 cm3 of
hydrrogen peroxiide solution in the flask.
5.
Add a small piecce of potato.
6.
Test the gas giveen off with a glowing splint.
7.
Repeeat this expeeriment with 20cm3 of yeeast suspension preparedd by suspendding 2g of
powdered yeast in 160 cm3 of
o water andd the mixturee left for 4hrss to aerate.
19
Questions
1.
Which enzyme source produces the fastest reaction (liver, potato or yeast suspension)?
2.
Write a word equation for this reaction.
3.
Describe how the volume of gas produced in the reaction may be measured.
3.
How could the rate of gas production be measured?
20
WEEK 7, EXPERIMENT 7
Construction of an Electrochemical cell (EMF)
AIM: To measure the Electromotive force generated by an electrochemical cell.
APPARATUS:
• High resistance voltmeter
• Crocodile clips
• Potassium nitrate solution
• Stripes of zinc
• Copper
• Iron.
WHAT TO RECORD: The EMF generated (voltmeter reading)
Procedure:
Copper-Zinc Voltaic Cell
The apparatus shown here is used to channel
electron flow in this reaction and is called a
voltaic cell. The zinc metal is on the right side
and it is providing electrons to the blue copper
ions in the beaker to the left. However, to get
there the electrons have to travel through a wire.
We can see that they are doing so because the
voltmeter that they must pass through shows a
reading.
Couple a voltaic cell with copper metal and copper sulphate solution in the left hand beaker; zinc
metal and zinc sulfate in the right hand beaker; with filter paper soaked in potassium nitrate
solution as the salt bridge. With the wires connected to the voltmeter, you can see that there is a
voltage reading; electricity is flowing through this cell. It's a voltaic cell because the chemical
reaction is causing the flow of electric current. Read and record the voltmeter. Repeat the
experiment by using the following:
™ ALL SOLUTIONS MUST BE 1 mol dm-3
™ Copper metal/copper nitrate solution (left); lead metal/lead nitrate (right).
™ Copper metal/copper sulphate (left); Iron metal/iron sulphate (right)
The importance of the salt bridge:
Remove the salt bridge and confirm that the voltage reading is zero, showing there is no current
flowing.
™ Replace the voltmeter with a bulb and observe the glowing bulb
21
22
WEEK 8, EXPERIMENT 8
Construction of an Electrochemical cell (EMF) with alternative salt bridge and a wider
variety of metal electrodes
AIM: To measure the Electromotive force generated by an electrochemical cell.
APPARATUS:
• Voltmeter
• crocodile clips
• sodium chloride solution
• stripes of zinc
• Copper
• Lead
• Iron
• Magnesium.
WHAT TO RECORD: The EMF generated (voltmeter reading)
Procedure:
Copper-Zinc Voltaic Cell
The apparatus shown here is used to channel
electron flow in this reaction and is called a
voltaic cell. The zinc metal is on the right side
and it is providing electrons to the blue copper
ions in the beaker to the left. However, to get
there the electrons have to travel through a wire.
We can see that they are doing so because the
voltmeter that they must pass through shows a
reading.
Couple a voltaic cell with copper metal and copper sulphate solution in the left hand beaker; zinc
metal and zinc sulfate in the right hand beaker; with sodium chloride solution in the salt bridge.
With the wires connected to the voltmeter, you can see that there is a voltage reading; electricity
is flowing through this cell. It's a voltaic cell because the chemical reaction is causing the flow
of electric current. Read and record the voltmeter. Repeat the experiment by using the
following:
™ ALL SOLUTIONS MUST BE 1 mol dm-3
™ Copper metal/copper nitrate solution (left); lead metal/lead nitrate (right).
™ Copper metal/copper sulphate (left); Iron metal/iron sulphate (right)
™ Copper metal/copper sulphate (left); magnesium metal/magnesium sulphate
The importance of the salt bridge: Remove the salt bridge and confirm that the voltage reading
is zero, showing there is no current flowing.
Exercise:
1. Use the standard electrode potential table and calculate the potential difference in each of
the cells. Arrange the metals in order of reactivity.
23
Example:
Zinc metal is reacting with the copper ion to form zinc ion and copper metal. The zinc metal is
on one side and the copper ion in solution is on the other side in the other beaker. Consequently,
the two chemicals are not in direct contact with one another.
Even so, the tendency for the zinc metal to lose (or transfer) electrons to the copper ion still
exists. That transfer is made possible by connecting a wire between the zinc metal and the
copper metal. The electrons go from the zinc over to the copper metal where they can react with
the copper ions in solution.
Over a period of time, the zinc electrode will dissolve and increase the concentration of the zinc
ion solution. The copper ion will plate onto the copper electrode and thus the concentration of
the copper ion in solution will decrease. As this happens, the reaction slows down and the
voltage decreases.
Zinc is the anode because that is where
oxidation is occurring. The oxidation halfreaction is Zn
Zn2+ + 2e-. Those electrons
go over to the copper side where they react
with copper ion and change it into copper metal
Cu), which is the reduction half(Cu2+ + 2ereaction that takes place at the cathode.
Zn
Zn2+ + 2e0.76v (right)
2+
Cu + 2e
Cu +0.34v (left)
____________________________
Zn + Cu2+
Zn2+ + Cu 1.10v (over all)
2. Explain why the calculated EMF may be different from the measured values.
24
WEEK9, EXPERIMENT 9
Determination of Molar Mass by application of Boiling-Point Elevation
AIM: To determine Molar Mass by application of Boiling-Point Elevation
SAFETY: Wear eye protection and hand gloves
APPARATUS:
•
250 cm3 Beakers •
250cm3conical flask •
Thermometer PROCEDURE:
Experimental Determination of Kb (boiling point elevation constant or the ebullioscopic
constant or simply boiling point constant)
1. Boil water on a hot plate in a clean 250 cm3 beaker. Measure the boiling temperature.
2. Experimentally determine the mass of 50.0 cm3 of distilled water (in your procedure include
how you determined this).
3. Dissolve about 12.5g of lactose in the 50.0 cm3 of distilled water using a 100 cm3 beaker.
4. Measure the boiling temperature of the solution using the thermometer used in step 1.
Calculation: Determine an experimental value for Kb this value will be used as the Kb in the
molar mass determination. Determine the % error of this Kb value.
Lactose is a hydrate so when determining the molality be sure to include to find the molar mass
of (C12 H22 O11. H2O).
Determination of Molar Mass of an Unknown
1. Experimentally determine the mass of 50.0 cm3 of distilled water (in your procedure include
how you determined this).
25
2. Dissolve about 7.5 g of Glucose (C12 H22 O11) in the 50.0 cm3 of distilled water using a 100
cm3 beaker.
3. Measure the boiling temperature using the thermometer used in experimental determination
of Kb.
Calculation:
Use the equation below to determine an experimental value for the molar mass of the glucose.
∆
10
MB = unknown molar mass of the solute
Ebullioscopic constant = Mass of solvent in grams = Mass of solute in grams ∆ = Increase in boiling point Questions:
1. Why was the use of high concentrations of solute required?
2. If tap water were used instead of distilled water in the experiment, what would be the
effect on the boiling point elevation (would it be larger, smaller or be the same). Why?
(assume the tap water is uniform throughout)
26
WEEK 10, EXPERIMENT 10
Determination of the freezing point of glacial acetic acid and molecular weight of an
unknown organic compound
AIM: To determine the freezing point of glacial acetic acid and molecular weight of an
unknown compound.
SAFETY: Wear eye protection and don’t inhale fumes
Materials:
•
50cm3Teflon beaker •
100cm3Teflon beaker •
500cm3 beaker •
Test tube (20 x 150 mm) •
Thermometer •
Bunsen flame •
Stirrer •
Glacial acetic acid •
Unknown organic compound Part I. Determination of the freezing point of the glacial acetic acid.
1. Prepare an ice bath by filling a 100cm3Teflon beaker 2/3 full with crushed ice and a
small amount of water. You can use two beakers, one inserted inside the other, for
better insulation, and you may want to place the beaker in a beaker (500 cm3) for
better support.
27
2.
Obtain an empty, dry large test tube (20 x 150 mm) and add 10.0 cm3 (density =
1.049 g/ cm3) of glacial acetic acid using a 10 cm3 pipette. Place a thermometer in the
test
tube and record the temperature under time = 0 seconds on the data sheet.
3. Insert the test tube deep into the ice bath, carefully stirring the solution continuously
with the thermometer, and record the temperature every 30 seconds on the data sheet.
Note the initial temperature at which solid first appears in the glacial acetic acid. If
you need to repeat the process, you may warm up the glacial acetic acid in a water
bath and go to step 3. After recording the temperature for 300 seconds (5 minutes),
warm the test tube back to room temperature with a water bath and save for Part
II.Using graph paper, plot temperature (Y) versus time (X) and determine the
freezing point of the glacial acetic acid using the equation below.
∆
10
= mass of solute in grams
MB = Molar mass
10
= is the mass of solvent in Kg
ΔTf =depression of freezing
Kf = cryoscopic constant
Part II. Determination of the Molecular Weight of an Unknown Compound
1. Obtain a sample of unknown solid from your laboratory instructor and record its identification number on the report form. The unknown is a molecular organic compound (not ionic) so we know that n, the Van’t Hoff factor = 1 in the Van’t Hoff equation. Weigh about 0.500 g of the unknown on a weighing paper and record its weight on the report sheet. 28
2. Carefully add it to the test tube containing the glacial acetic acid (re‐warmed back to room temperature) from Part I. Stir the mixture with glass rod until the entire unknown has dissolved. Replace the thermometer and record the temperature (time = 0) on the data sheet. 3. Insert the test tube deep into the ice‐water mixture, stirring the mixture continuously, and record the temperature every 30 seconds on the data sheet as before. Mark the initial temperature at which solid first appears in the mixture. To repeat the process you may warm up the mixture in a water bath back to room temperature and start at step 3. 4. Determine the molecular mass of the unknown from the equation above. NOTE ‐ Some supercooling may occur before freezing. Supercooling results when a solution cools below its normal freezing point without freezing. To prevent supercooling, steady stirring of the liquid is required. Once solid begins to form, further crystallization usually occurs rapidly without supercooling. 29
WEEK 11, EXPERIMENT 11
Effect of concentration on Boiling Point Elevation
AIM: To determine the effect of concentration of solution on the elevation of boiling point
INTRODUCTION:
Dissolving a solute in a liquid can change its properties. In this experiment you will
discover what effect, increasing concentrations of salt (NaCl) have on the boiling
temperature of water. To do this you will need to make various solutions and then test
their boiling temperatures. Make sure you are careful; it’s easy to make little mistakes.
SAFETY: Wear eye protection and hand gloves
Materials:
•
Bunsen burner set-up
•
250 cm3 beaker
•
Salt (NaCl)
•
Tap water
•
100 cm3volumetric flask
•
Tongs
•
Thermometer
•
Glass funnel
PROCEDURE:
i. Make 100 cm3 of 1 mol dm-3 NaCl solution using distilled water
ii. Heat to boil the solution
iii. Record the temperature when boiling starts
iv. Make 100 cm3 of 0.5 mol dm-3 NaCl solution
30
v. Boil and record the temperature
vi. Make 100 cm3 of 1.5 mol dm-3 NaCl solution
vii. Boil and record the temperature
viii. Make 100 cm3 of 2.0 mol dm-3 NaCl solution
ix. Boil and record the temperature
x. Make 100 cm3of 2.5 mol dm-3 NaCl solution
xi. Boil and record the temperature
xii. Make 100 cm3 of 3.0 mol dm-3 NaCl solution
xiii. Boil and record the temperature
xiv. Boil distilled water and record the temperature
DATA TABLE:
Make a data table to record the data from the procedure.
EXERCISE
1. Plot concentration on the y-axis against temperature on x-axis
2. From your graph explain the effect of concentration on the boiling points of the
different solutions
31
WEEK 12, EXPERIMENT 12
Determination of the effect of different solutes on the Boiling Point Elevation of solutions
in a particular solvent
AIM: To determine the effect of different solutes on the Boiling Point Elevation of solutions in a
particular solvent.
SAFETY: Wear eye protection and hand gloves
Materials:
•
Distilled water
•
Calcium chloride
•
Sucrose
•
Sodium chloride
•
Bunsen burner set-up or heating mantle
•
250 cm3 beaker
•
100 cm3volumetric flask
•
Tongs
•
Thermometer
•
Glass funnel
WHAT TO RECORD:
Make a data table to record the data from the procedure. PROCEDURE:
i. Dissolve 10g of calcium chloride in 100g of distilled water in a 250 cm3 beaker. ii. Heat to boil the solution. iii. Record the temperature when boiling starts. 32
iv. Repeat steps 1 to 3 using Sucrose and Sodium chloride. v. Measure 100g of the distilled water into 250 cm3 beaker. vi. Heat to boil the water. vii. Record the temperature when boiling starts. EXERCISE:
1. Determine the elevation of boiling point for each of the solutions. 2. Arrange the solutes in order of increasing effect on the elevation of boiling point of distilled water. 33
WEEK 13, EXPERIMENT 13
Determination of the effect of different solutes on depression of freezing point of solutions
in a particular solvent
AIM: To determine the effect of different solutes on the depression of freezing point of solutions
in a particular solvent
Materials:
•
50cm3Teflon beaker •
100cm3Teflon beaker •
500cm3 beaker •
Test tube (20 x 150 mm) •
10 cm3 pipette •
Thermometer •
Bunsen flame •
Stirrer •
Distilled water
•
Calcium chloride
•
Sucrose
•
Sodium chloride
Procrdure:
1. Prepare an ice bath by filling a 100cm3Teflon beaker 2/3 full with crushed ice and a
small amount of water. You can use two beakers, one inserted inside the other, for
better insulation, and you may want to place the beaker in a beaker (500 cm3) for
better support.
34
2.
Obtain an empty, dry large test tube (20 x 150 mm) and add 10.0 cm3 of distilled
water using a 10 cm3 pipette. Place a thermometer in the test tube and record the
temperature under time = 0 seconds on the data sheet.
3. Insert the test tube deep into the ice bath, carefully stirring the solution continuously
with the thermometer, and record the temperature every 30 seconds on the data sheet.
Note the initial temperature at which solid first appears in the distilled water. If you
need to repeat the process, you may warm up the distilled water in a water bath and
go to step 3. After recording the temperature for 300 seconds (5 minutes), warm the
test tube back to room temperature with a water bath.Using graph paper, plot
temperature (Y) versus time (X) and determine the freezing point of the different
solutions using the equation below.
∆
10
= mass of solute in grams
MB = Molar mass
10
= is the mass of solvent in Kg
ΔTf =depression of freezing
Kf = cryoscopic constant
EXERCISE:
viii. Explain the effect of the different solutes on the depression of freezing point of water. ix. Arrange the solutes in order of increasing effect on the depression of freezing point of distilled water. 35
WEEK 14, EXPERIMENT 14
Determination of molecular weight of naphthalene by Rast’s method
AIM: To determine the molecular weight of naphthalene from the freezing point constant for
camphor
SAFETY: Wear eye protection and hand gloves
MATERIALS:
•
Ignition tubes
•
Thermometer
•
Olive oil
•
Camphor
•
Naphthalene
Thermometer
Retort stand
Ignition tube
Boiling tube
Olive oil
Heat
PROCEDURE:
Weigh an ignition tube empty and weigh 1g of camphor into it. Add 0.1g of naphthalene to the
tube and weigh. Insert a thermometer with a piece of rubber tubing to act as a cork. Place the
36
tube in another boiling tube containing olive oil and warm gently until the whole mass in the
ignition tube melts. Remove the ignition tube and adjust the thermometer so that it dips into the
molten contents. Replace the tube and allow to cool while in the oil. Record the temperature
when the camphor becomes cloudy and take this as the freezing point. Repeat the experiment in
another ignition tube using only camphor.
EXERCISE:
1. Calculate the molecular weight of naphthalene from the equation below given the
freezing point constant for camphor as 400 :
∆
′
10
= mass of solute in grams
MB = Molar mass
′
10
= is the mass of solvent in Kg
ΔTf =depression of freezing
Kf = cryoscopic constant = 400
2. What is the advantage of the high value for the freezing point constant of camphor in this
experiment?
37
WEEK 15, EXPERIMENT 15
Determination of the boiling point constant (Ebullioscopic constant) for water by
Landsberger method.
AIM: To determine the boiling point constant for water by Landsberger method.
MATERIALS:
•
Landsberger apparatus
•
Boiling cans or 500cm3 beaker
•
Glazed paper
•
Cane sugar (urea may be used)
Thermometer
Steam
Wooden cork
Rubber bung
Boiling tube
Gas- jar
Glass tube
38
PROCEDURE:
The modified Landsberger apparatus consists of a boiling tube with a small hole blown
about 2cm from the top and fitted loosely with a bung which carries a glass tube and
thermometer. The whole fits into a gas jar, a cork holding the boiling tube in place.
Weigh accurately 7g of cane sugar on a weighed piece of glazed paper. Fill the boiling
tube half way with water and pass steam via a delivery tube. Record the temperature
when the temperature becomes steady. This is the boiling point of water at the pressure of
the atmosphere. The bung in the boiling tube must be loosely fitted throughout (to
prevent the cracking of the tube). Remove this bung, slide in the sugar cane and heat to
determine the new boiling point. Weigh the solution immediately and subtract the weight
of solute to calculate the weight of water.
EXERCISE: Calculate the boiling point constant from the equation below:
′
∆
′
MB = unknown molar mass of the solute
Ebullioscopic constant (boiling point constant)
′
= Mass of solvent in grams
′
= Mass of solute in grams
∆ = Increase in boiling point
39
10