Experiment 1
Charles’s Law: Develop Your Own Procedure
Objectives
To learn how to design an experiment and develop your own
procedure
To investigate the relationship expressed in Charles‘s law
Principle
The relationship between the volume of a gas and the absolute
temperature at constant pressure is known as Charles‘s Law. The law
states that the volume of a fixed quantity of gas at constant pressure
increases linear with the absolute temperature.
The experiment involving Charles’s Law has been chosen since the
relationship between volume and temperature is easily understood and
virtually no chemicals are necessary. The students are asked to write their
own procedure and execute it in the Lab. They should learn the meaning
of variables and controls within an experiment. Then the students
examine several simple experiments for validity by posing the following
questions:
(1) Are the controls adequate and the variables changed one at a time?
(2) Are enough data collected?
(3) Is the purpose achieved?
Handout
(1) Balloon---spherical shape
(2) String
(3) Metric ruler
1
(4) Ice
(5) Hot plate
(6) Thermometer
(7) 2 L beaker
(8) Ring stand
(9) Ring clamp
(10) Wire gauze
(11) Tongs
(12) 100 ml Graduated cylinder
(13) 1 liter beaker
(14) 125 ml Erlenmeyer flask
(15) Rubber stopper
(16) Glass tube
(17) rubber tube
(18) Clamp
Questions
According to the following questions to design your procedure:
(1) What container will hold the gas molecules?
(2) How will the number of molecules of gas be held constant?
(3) How can the volume of the gas be measured?
(4) Recall that the formula for the volume of a sphere is
V=4πr3/3
(5) How will the temperature of the gas be varied?
(6) How many volume-temperature measurements should be made?
(7) Make a V vs. T plot.
2
(8) What kind of information you can obtain?
(9) What is your experimental error?
(10)What is the most difficult obstacle in this experiment?
(11)How does the pressure be kept constant?
NOTE : Before beginning the lab, the students must have a
completed procedure and data table.
Procedures
For this experiment, it is important to control variables such as
pressure and the number of molecules so that the volume-temperature
relationship can be studied.
(1) Students made four to six volume measurements from 0 - 100°C.
(2) Measure the balloon’s circumference.
(3) Present the data in the table and plot temperature vs volume of gas in
the squarepaper.
(4) Discussion and conclusion.
Reference
Assemble the apparatus shown in Fig. 1.
Figure 1
3
The bottom of the gas-delivery tube should extend about 1 in. below the
bottom of the Erlenmeyer flask. The flask and delivery tube must be
arranged so that both may be placed into a 1-liter beaker, and a 50-ml
graduated cylinder placed over the opening of the delivery tube, as
shown in Fig. 1. Fill the 1-liter beaker with water while the flask is
immersed, so as to have the proper quantity of water in the beaker for
later use. Dry the flask by warming it gently over a low flame, allow the
flask to cool to room temperature, insert the stopper bearing the delivery
tube and thermometer, and support the flask as shown in Fig. 1. While
the flask is cooled down to room temperature, heat the water in the
beaker to about 75℃. Fill the 50-ml graduated cylinder with water
previously boiled to expel dissolved air. Invert the cylinder and place its
mouth at the bottom of the beaker while the water in the beaker is being
heated, so that the cylinder, its contents, and the water in the beaker may
come to the same temperature.
Read to the nearest degree the temperature of the air in the
Erlenmeyer flask and record it on the data sheet. Quickly lower the flask
into the 1-liter beaker to a position such that the open end of the delivery
tube is covered by the hot water to depth of about 1-in. At once, have
another student or the instructor place the cylinder over the opening of
the delivery tube, taking care that no air gets into the cylinder from the
outside. Now lower the flask as far as possible, so that the apparatus is in
the condition represented by Fig. 1. Since the air in the flask is heated by
the hot water, the air expands rapidly and it is necessary to collect in the
cylinder all of the air that is expelled from the Erlenmeyer flask when its
temperature is raised. If any air is lost, the experiment must be started
over again.
Measure the temperature of the hot water and record it as the final
temperature. When no more air flows into the cylinder, lift it gently from
the opening of the delivery tube, but keep the mouth of the cylinder
below the level of the water in the beaker. Raise or lower the inverted
cylinder until the level of the water in it is the same as that in the beaker,
and holding the cylinder in this position, read the volume of gas in the
cylinder. Record this value on the data sheet. Determine the original
volume of air by marking with a gummed label the position of the
4
bottom of the rubber stopper in the Erlenmeyer flask, and measure with a
100-ml graduated cylinder The volume of water required to fill the flask
to this mark.
Since the gas collected in the cylinder consists of a mixture of air and
water vapor, the volume occupied by the air alone must be calculated.
The instructor will provide the correct value for the prevailing barometric
pressure. Values for the vapor pressure of water at different temperatures
are tabulated in the Appendix. The resulting volume plus the original
volume of air contained in the flask at the beginning of the experiments is
equal to the final volume of dry air at the higher temperature.
5
Experiment 2
Determination of the Acid Content of Vinegars
and Wines
Principle
NEUTRALIZATION TITRATIONS
Neutralization titrations are performed with standard solutions of
strong acids or bases.
A single solution of either acid or base is sufficient for the titration of a
given type of analyte, but to locate end points more exactly with back
titration, it is convenient to have standard solutions of both acid and base.
The concentration of one solution is established by titration against a
primary standard; the concentration of the other is then determined from
the acid/base ratio (that is, the volume of acid needed to neutralize
1.000mL of the base ).
The total acid content of a vinegar or a wine is readily determined by
titration with a standard base. It is customary to report the acid content of
vinegar in terms of acetic acid, the principal acidic constituent, even
though other acids are present. Similarly, the acid content of a wine is
expressed as percent tartaric acid, even though there are other acids in
the sample. Most vinegars contain about 5% acid (w/v) expressed as
acetic acid; wines ordinarily contain somewhat under 1% acid (w/v)
expressed as tartaric acid.
A. The Effect of Atmospheric Carbon Dioxide on Neutralization
Titrations
When water is in equilibrium with the atmosphere, the concentration
5
of carbonic acid is about 1 10 M.
6
CO2(g) + H2O H2CO3(aq)
At this concentration level, the amount of 0.1 M base consumed by the
carbonic acid in a typical titration is negligible. With more dilute reagents
( < 0.05 M ), however, the water used as a solvent for the analyte and in
the preparation of reagents must be freed of carbonic acid by boiling for a
brief period.
Water that has been purified by distillation rather than by
deionization is often supersaturated with carbon dioxide and may thus
contain sufficient acid to affect the results of an analysis(Note 1). The
instructions that follow are based upon the assumption that the amount of
carbon dioxide in the water supply can be neglected without causing
serious error. For further discussion on the effects of carbon dioxide in
neutralization titrations, see textbook.
B. Preparation of Indicator Solutions for Neutralization Titrations
Stock solutions ordinarily contain between 0.5 and 1.0 g of indicator per
liter. (One liter of indicator is sufficient for hundreds of titrations. )
Phenolphthalein ---- Dissolve the solid Phenolphthalein in a solution
consisting of 800 mL ethanol and 200 mL of distilled or deionized water
Note 1:
Water that is to be used for neutralization titrations can be tested by
adding 5 drops of phenolphthalein to a 500-mL portion. Less than 0.2 to
0.3 mL of 0.1 M OH—should suffice to produce the first faint pink color
of the indicator. If a larger volume is needed, the water should be boiled
and cooled before it is used to prepare standard solutions or to dissolve
samples.
C. Preparation of Carbonate-Free Sodium Hydroxide
The concentration of solutions of sodium hydroxide decreases
slowly ( 0.1% to 0.3% per week ) when the base is stored in glass bottles.
7
The loss in strength is caused by the reaction of the base with the glass to
form sodium silicates. For this reason, standard solutions of base should
not be stored for extended periods ( longer than 1 or 2 weeks ) in glass
containers. In addition, bases should never be kept in glass-stoppered
containers because the reaction between the base and the stopper may
cause the latter to “ freeze ” after a brief period. Finally, to avoid the same
type of freezing, burets with glass stopcocks should be promptly drained
and thoroughly rinsed with water after use with standard base solutions.
This problem is avoided with burets equipped with Teflon stopcocks.
If so directed by the instructor, prepare a bottle for protected storage.
Transfer 1 L of distilled water to the storage bottle. Decant 4 to 5 mL of
50% NaOH into a small container, add it to the water, and mix thoroughly.
Use extreme care in handling 50% NaOH, which is highly corrosive. If
the reagent comes into contact with skin, immediately flush the area with
copious amounts of water. Protect the solution from unnecessary contact
with the atmosphere.
Apparatus and Reagents
1. 25 ml pipet --- 1
2. 250 ml volumetric flask --- 2
3. Buret --- 1
4. Phenolphthalein solution
5. NaOH
6. Vinegar --- 25 ml
7. Wine --- 50 ml
Before beginning the lab, you must have a completed-design procedure
flow chart and Data table.
Procedure
(a) If the unknown is a vinegar ( Note 1 ), pipet 25.00 mL into a 250-mL
volumetric flask and dilute to the mark with distilled water. Mix
thoroughly, and pipet 50.00-mL aliquots into 250-mL flasks. Add
about 50 mL of water and 2 drops of phenolphthalein ( Note 2 ) to
each, and titrate with standard 0.1 M NaOH to the first permanent (
30 s ) pink color.
Report the acidity of the vinegar as percent (w/v) CH 3 COOH (60.053
8
g/mol).
(b) If the unknown is a wine, pipet 50.00-mL aliquots into 250-mL
conical flasks, add about 50 mL of distilled water and 2 drops of
phenolphthalein to each ( Note 2 ), and titrate to the first permanent
( 30 s ) pink color.
Express the acidity of the sample as percent (w/v) tartaric acid
C2 H4 O2 ( COOH)2 (150.09 g/mol ). (Note 3)
Notes:
1. The acidity of bottled vinegar tends to decrease on exposure to air. It
is recommended that unknowns be stored in individual vials with
snug covers.
2. The amount of indicator used should be increased as necessary to
make the color change visible in colored samples.
3. Tartaric acid has two acidic hydrogens, both of which are titrated at a
phenolphthalein end point.
Report
Your designed procedure flow chart and Data table
Result and Discussion
1. Calculate the percentage of acetic acid in vinegar and tartaric acid in
wine.
2. Define titration error.
9
Experiment 3
Vitamin C Content of commercial Orange
Juices
It has become customary at the University for the first year,
nonchemistry major students, to conduct a survey involving the analysis
of a consumer item and to perform a detailed statistical evaluation of the
results.
This project involves a study of the magnitude and stability of the
ascorbic acid content of commercial orange juices. This study had two
aims: firstly to confirm that newly purchased juice contained sufficient
ascorbic acid to meet government standards ( at present 40 mg/ 100 ml of
juice ), and secondly to establish the rate of aerial oxidation of this
ascorbic acid when the juice was stored in a refrigerator. Students were
interested in ascertaining if the juice represented a suitable source of
Vitamin C after one or two weeks of storage.
Principle
In view of the background of the students, we selected a redox
titration procedure. The most satisfactory methods for determination of
ascorbic acid in orange juice were those using the reagents 2,6
dichlorophenolindophenol ( 2,6 DPIP ) and N-bromosuccinimide ( NBS ).
In order to provide comparison data for use in the statistical analysis
section of the project, both of these reagents were used concurrently.
The reaction of 2,6 DPIP with ascorbic acid is shown below.
10
2,6 DPIP is blue in neutral or alkaline solution and red in acid solution
while its reduced form is colorless. Thus when titrating an acidic solution
of ascorbic acid against blue 2,6 DPIP, the blue reagent will turn colorless
while ascorbic acid is present, but when all the ascorbic acid has been
consumed, any excess 2,6 DPIP added will turn the solution pink, which
is where the endpoint is.
NBS reacts with ascorbic as follows
11
Use of a starch iodide indicator provides a color change of colorless
to blue at the endpoint.
Apparatus and Reagents
1. Orange juice ( Five brands )
2. 2.6 DPIP 1.25g
3. NaHCO3 1.05g
4. NBS 1g
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5. Celite or Supercel ( Filter aid )
6. Oxalic acid 0.2g , 2g
7. Ether
8. ascorbic acid 50mg
9. 4% KI
10. 10% acetic acid
11. starch indicator
12. Buchner funnel
13. flask -250ml
14. separating funnel
15. Cylinder
16. Buret
Experimental
Samples and Reagents
Five brands of juice were used for the experiment and 2 l of each
were purchased ( this was sufficient for 50 students ). In selection of the
brands, consideration was given to the different preservatives allowed
( sulfur dioxide, sorbic acid, and benzoic acid ) and also to the different
forms of packages available, so that the samples chosen were
representative of the range sold on the market. In addition, to minimize
sample variation, the juice was obtained in 1-l containers and then
combined to produce the volume required for the project.
Each brand of juice was stored in a separate sealed plastic container
at 0C and was lightly exposed to the air daily by shaking and pouring the
contents into another container. This process was intended to simulate
normal household treatment of orange juice. Two liters of this juice were
supplied on each practical day attended by the group which had been
allotted that particular brand for analysis.
2,6 DPIP is prepared daily by dissolving 1.25g of 2,6 DPIP and
1.05g NaHCO 3 in 5 l distilled water. NBS containing 1.0 g NBS in 5 l
distilled water is prepared fresh daily. These quantities were sufficient for
30 students.
13
General
The analysis may be conveniently divided into three stages, these
being preparation of sample, standardization of oxidizing agent, and
titration procedure.
While it is possible to perform the titration directly on orange juice,
the color of the juice can sometimes obscure the endpoint. For this reason,
it is advisable to either filter off the pulpy components of the juice or to
extract the orange coloring matter using ether. Orange juice is very
difficult to filter because the suspended matter quickly clogs the pores of
the filter paper; use of a filter aid such as celite or supercel is therefore
necessary.
The stock solutions of NBS ( ~0.001 M ) and 2,6 DPIP ( ~0.001 M )
are unstable and it is therefore necessary to standardize them before use.
This standardization may be accomplished using many common primary
standards; however it is more convenient to use pure ascorbic acid as
standard.
Procedures
Each group select three brands of juice to do the following
experiment.
Preparation of Samples --- choose one of following two methods
(1) Take about 30 - 40 ml of orange juice, add 0.2 g solid oxalic acid
( CARE! Oxalic acid is poisonous ), and swirl until the oxalic acid
has dissolved. The purpose of the oxalic acid is to stabilize the
ascorbic acid in the juice, which could otherwise be oxidized during
the filtration process, and to acidify the titration solution.
Add sufficient filter aid to the orange juice to form a thin paste and
filter this paste through a Buchner funnel into a clean, dry flask. The
filtration should proceed rapidly, but if this is not the case, remove the
juice and add a little more filter aid. The resulting filtered solution
should be only pale range in color and quite transparent. This solution
is now ready for analysis.
(2) An alternative method of sample preparation is to extract the orange
color into ether. Take about 30 - 40 ml of orange juice, add 0.2 g solid
14
oxalic acid and swirl until the oxalic acid has dissolved. Transfer this
solution to a separating funnel and add about 20 ml ether from a
measuring cylinder ( CARE! Ether is highly inflammable - ensure
that there are no naked flames in the laboratory ). Gently shake the
funnel - too much shaking will result in the formation of an emulsion
which will take a long time to separate. The orange color will be
transferred to the ether ( upper ) layer and most of the pulp will tend
to collect at the phase boundary between the aqueous and organic
layers. Drain off the aqueous phase leaving as much as possible of
the pulp in the separating funnel. The aqueous phase is now ready for
analysis.
Standardization of Oxidizing Agent.
Weigh accurately about 50 mg of ascorbic acid and quantitatively
transfer it to a 250-ml volumetric flask containing about 2 g solid
oxalic acid. Fill the flask to the mark and shake gently until all solid
has dissolved. This solution may be used to standardize the
N-bromosuccinimide and 2,6 DPIP solutions provided using the
procedures described below.
NBS: Add a 25-ml aliquot of the standard ascorbic acid solution
prepared above to a flask containing 5 ml of 4% potassium iodide
solution, 2 ml of 10% acetic acid solution, and 3 drops of starch
indicator. Dilute with about 30 ml of distilled water and titrate with
the NBS solution provided. The endpoint is marked by the
appearance of a permanent blue color.
2,6 DPIP: Add a 25-ml aliquot of the standard ascorbic acid
solution prepared above to a conical flask, dilute with about 30 ml
of distilled water, and titrate with the 2,6 DPIP solution provided.
The endpoint is marked by the appearance of the first permanent
pink color.
Titration Procedure
Using a 5-ml aliquot of prepared orange juice, proceed as outlines
above in “standardization of oxidizing agent” with the orange juice
aliquot replacing the 25-ml aliquot of standard ascorbic acid
15
solution. If a titer of less than 10 ml is obtained, the volume of
orange juice used should be increased to 10 ml, or further if
necessary. The titration should be repeated until consistent results
are obtained.
Result and Discussion
The students should present the results graphically.
Discuss in detail all possible sources of error.
16
EXPERIMENT 4
Crystal Growth of Water Soluble Substances in
Gel Media
Objectives
To understand the major concepts of nonlinear dynamics in far from
equilibrium condition
To become acquainted with the phenomenon of crystallization in gel
media
To learn about the laws of probability from the study of real fractal
phenomena in nature
Principle
Pattern formation processes under nonequilibrium conditions has
been stimulated by several recent developments. Biology has potential for
the application of fractal concepts. Biological growth almost invariably
leads to the formation of complex shapes, forms, and patterns, such as the
growth of nerve cells and blood vessels.
The growing patterns have been classified into two groups: Diffusion
limited aggregation (DLA), whose pattern is irregular and fractal, and
dendritic crystallization, that forms regular growth patterns.
In this experiment, we will investigate a variety of systems that
exhibit different types of morphology when allowed to grow the thin
films of solutions containing a denser matrix such as agar-agar or PVA
Polymer. Irregular, fractal-like growth of KCrO4, SrCl2, KCl in agar-agar
gel and nutrient broth; tree-like geometry in the growth of KCr2O7; and
bands and highly branched morphology of SrCl2 when crystallized from
aqueous solution containing agar-agar and PVA polymer are investigated
17
here.
Reagents
potassium chromate
potassium dichromate
strontium chloride
potassium chloride
agar-agar
PVA (polyvinyl alcohol polymer)
slides
incubator
microscope
Procedure
1. For the development of patterns of water soluble substances in
denser matrices, known amounts should be dissolved in double
distilled water containing agar-agar gel or PVA of known
concentration.
2. A fixed volume of the solution was poured and spread
uniformly over microslides
3. The slides are placed in an incubator at fixed temperature (room
temperature) or in some cases evaporated in presence of
sunlight.
4. After the solvent is completely evaporated, slides are observed
with the help of a microscope are taken.
5. Conditions:
[SrCl2]=0.0025 M in 0.01％agar-agar -water--- 1ml
[SrCl2]=0.01 M in 0.01％agar-agar -water ---1ml
[SrCl2]=0.01 M in 0.1％agar-agar -water ---1ml
18
[SrCl2]=0.02 M in 0.01％PVA-water ---1ml
[KCl]=0.01 M in 0.1％agar-agar-water ---1ml
[KCl]=0.01 M in 0.1％PVA -water --- 1ml
[K2CrO4]=0.01M in 0.2％agar-agar-water --- 1ml
[K2Cr2O7]=0.01M in 0.2％agar-agar -water ---1ml
0.5ml [K2CrO4]＋0.5ml [K2Cr2O7] --- 1ml
6. In the case of strontium chloride, it is performed at various
conditions differing in solute concentrations and volume of the
solution spread over the slide, which are grown in agar-agar or
PVA matrix.
Discussion
1. How do you identify CrO4-2 and Cr2O7-2 ions from this experiment?
2. Describe the growth patterns for those conditions in procedure 5 and
draw the graphs.
3. How do SrCl2 and KCl crystallize depend on the nature of the gel
matrix and concentration of the solute in the gel?
19
EXPERIMENT 5
Chemical Oscillations and Spiral
Waves----kinetics experiment
Part I : Chemical Oscillation—Traffic Light
Objectives
To study the behavior of the in-homogeneous system
With a qualitative study of the influence of chemical and physical
variables on the rate of a specific reaction
To understand nonlinear dynamics and far from equilibrium
thermodynamics and the phenomena of temporal and spatial
self-organization
Principle
The reaction system used in this experiment is given below.
HOOC-CH2-COOH + 6 Ce4+ + 2H2O
2CO2 + HCOOH + 6 Ce3+ + 6H+ (1)
10 Ce3+ + 2 BrO3- + 12 H+
10 Ce4+ + Br2 + 6 H2O (2)
The concentration of Ce4+ in solution periodically increases and decreases
with time and the change in the concentration of the Ce4+ is followed by
the repetitive color changes induced in the redox indicator Ferroin
(VIOLET to BLUE).
20
Reagents
Prepare the following solutions:
Table 1. Stock Reagent Solutions
Solution identification Concentration (molar)
A 0.00045 M Ce(NH4)4(SO4)4‧2H2O
in 3N H2SO4
B 0.090 M KBrO3 in 3N H2SO4
C 0.30M Malonic acid in 3N H2SO4
Ferroin indicator 0.1M Ferroin in 3N H2SO4
Table 2. Concentrations of Experimental Solutions
Solution Ce4+ BrO3- Malonic acid Acidity
1 0.00015M 0.030 0.10 3N H2SO4
2 0.00010 0.020 0.07 3N H2SO4
3 0.00008 0.015 0.05 3N H2SO4
4 0.00005 0.010 0.03 3N H2SO4
5 0.00005 0.010 0.03 2N H2SO4
6 0.00005 0.010 0.03 1N H2SO4
Procedure
1. prepare the reagents in Table 2.
2. mix the reagents for each solutions in a beaker and record the
oscillation time in second. YOU NEED TO SHAKE THE
SOLUTIONS.
3. (Temperature effect) Solution 1 is used for the investigation of the
influence of temperature on the period of oscillation. It is cooled to
15oC and the period of oscillation is recorded at 5o intervals as the
solution is warmed until the period of oscillation is too short for
precise observation (about 50oC).
4. The dilution study consists of noting the period of oscillation for
Solutions 1, 2, 3, and 4. A plot of the period of oscillation as a
function of the concentration of any one of the components, malonic
acid, Ce4+ or BrO3- is then made.
5. In the acidity study, record the period of oscillation of Solution 4, 5,
and 6 in the following table
21
Solution H+ concentration time of
oscillation
4 3N
5 2N
6 1N
Questions and Disscusion
(1) After several periods of oscillations, add 2 ml of a 25 mM ferroin
solution. Ferroin (tris(1,10-phenanthroline)iron(II) sulfate) is a redox
indicator. As the [Ce(IV)] increases, it can oxidize the Fe in ferroin
from Fe(II) to Fe(III). The Fe(II) complex is red, and the Fe(III) is
blue. Consequently, the color changes as the potential changes, What
effect does the ferroin have on the period and amplitude of the
oscillations? Is it really acting just as an indicator?
(2) Test the effect of Cl- by adding a small grain of NaCl. What happens?
How long does it take the system to recover? Add a pinch. What
happens?
(3) Test the effects of oxygen on the system by turning the stirrer to a
high speed. Turn it up full speed. Are the oscillations affected?
(4) Add a drop of acrylonitrile to an oscillating reaction and explain the
relevance of the white precipitate.
(5) Observe oscillations for a half-hour. Note changes in the period as a
function of time. (What is the relevance of this change to the second
law of thermodynamics?)
Part II: Spiral Waves ----Unstirred BZ system
Objective
To test this model by determining the sensitivity of front velocity to
concentration of ferroin, [BrO3-] and [H+]
To observe the phenomenology of waves and spirals
If an autocatalytic reaction is left unstirred, extremely interesting
behavior can be observed. Diffusion, instead of causing concentration
gradients to disappear, can couple to the autocatalytic reactions to
produce reactions that propagate through the medium.
22
Reagents
Note: All reagents andydrous; all solutions aqueous.
18 M H2SO4
NaBrO3
NaBr
Malonic acid
25 mM, 10 ml standard Ferroin--- 0.0695g FeSO4‧7H2O ( Mwt= 278)
+0.04955g 1,10-phenanthroline (Mwt=198.2) add 1 N H2SO4 till 10
ml
Triton X-100 surfactant
Test tube --- 1 tube
90 mm Petri dish --- 1 piece
Erlenmeyer flask --- 1 piece
Procedure
Stock solutions:
Solution A: NaBr (1g/10ml) in water
Solution B: Malonic acid (1 g/10 ml) in water
Solution C: 25mM Ferroin
Solution D: Triton X-100 surfactant (1 g/1000 ml)
2). Students prepare the following solutions:
Solution 1: 67 ml distilled water +2 ml, 18 M H2SO4 +5 g
NaBrO3
Solution 2: Take 6 ml of solution 1 in a test tube + 0.5 ml NaBr
(Solution A)+ 1 ml malonic acid (Solution B)
3). (DO NOT STIR IN THIS STEP) Solution 2 will produce Br2 (Brown
color). Wait for Brown color to vanish, then add 1 ml 25 mM Ferroin
and 1 drop Triton X-100 surfactant into solution 2, finally pour this
mixture into a 90 mm petri dish ( glass not plastic).
4). (DO NOT STIR IN THIS STEP) The reaction may oscillate between
red and blue, but ignore this. Cover the dish and wait.
Questions and Discussion
1. You will notice small rings of blue forming in the red solution. Notice
how rapidly the blue color spreads. Calculate how long a molecule
23
would take to diffuse just 1 cm ( D=10-5cm2/s, distance=α(Dt)1/2).
2. What happens when two waves collide? How is this different from
water waves?
3. What happens when a wave encounters a barrier? Does it reflect?
4. To make a spiral, slowly move a pipet tip through the center of a
target pattern. In which direction does the spiral rotate?
5. Repeat the above experiment. Use graph paper to measure the change
in the radial distance from the center of a target pattern as a function
of time. The slope of the line drawn through this data will provide the
wave speed. Also measure the wavelength (the distance between
fronts).
6. Does the velocity remain constant with time? If not, why?
7. Repeat these measurements for the following solutions.
8. Use 0.5 ml of ferroin instead of 1 ml in step 3. Is there any effect?
9. Does the velocity depend on the ferroin concentration?
24
EXPERIMENT 6
Determination of Thermodynamic Quantities:
, S and H
Objectives
To determine the thermodynamic values, S , G , and H from
thermodynamics experiment
Principle
This experiment is based on the measuring of the difference in
voltage between the high and low temperature, which can be related to
the change in entropy, S , by the relationship
G
S (1)
T P
where G is the Gibbs free energy and T is the absolute temperature, at
constant pressure, P. This equation can be written
G
S (2)
T
since individual data points will be measured, and the atmospheric
pressure will be assumed to be constant. Using the relationship between
Gibbs free energy and the voltage of an electrochemical cell
G nF (3)
substitution into eq.(2) yields
S nF (4)
T
where F is Faraday‘s constant and n is the number of moles of electrons
transferred in the cell reaction.
The battery consists of Pb/Pb+2 and Cu/Cu+2 half-cells connected by a
salt bridge. When equal concentrations of the Pb +2 and Cu+2 solutions are
used in both cells the second part of the Nernst equation becomes zero.
Assuming the concentrations (activities) of the Pb +2 and Cu+2 are equal,
and assuming that room temperature is 25℃, the standard voltage of the
25
cell will be equal to the measured voltage and will be 0.463V.
A temperature change of 55-60℃ should produce a voltage change
of 25-30 mV, a shift which is sufficiently large to be measured on the
millivolt scale of a standard digital voltmeter. The student should plot ε
versus T, and determine S from the slope of the line, / T , and
eq.(4).
Reagent and Apparatus
Note: The solutions are prepared with distilled water.
0.5 M CuSO4 --- 50 ml
0.5 M Pb(NO3)2 --- 50 ml
Pb and Cu electrode
600 ml Beaker
Digital millivoltage
2 Alligator clips
U-tube
Agar-KNO3 Salt bridges ---- A 50 ml electrolytic solution is prepared by
dissolving agar (2％ by weight) and KNO3 ( 10% by weight) in 38 ml of
distilled water, and heating by water bath until the mixture thinned,
thickened, then thinned again. The hot solution is poured into wet glass
U-Tubes, and set up in 10-15 min for cooling.
Procedure
1). Set up Agar-salt bridges and prepare solutions 0.5M CuSO4 and 0.5
M Pb(NO3)2
2). Pour 25-40 ml of 0.5M CuSO4 and 0.5 M Pb(NO3)2 solution each
separately into a 50 ml test tube and place Cu and Pb electrode into
the test tube appropriately
3). The test tubes are placed in a 600 ml beaker, the salt bridges are
added, and the millivoltage is attached to the electrodes with alligator
clips. The salt bridges connect the electrolytic solution to both the
Pb/Pb+2 and Cu/Cu+2 half-cells.
4). Measureε at room temperature and record it on the data sheet.
26
5). Add ice and water to the 600 ml beaker to cool the cell until the
temperature has stabilized 5-10 min, and measureεand T.
6). Warm the solution to about 70℃, and measureεand T.
7). The value ofεat room temperature (25℃) is used in eq.(3) to
determine a value for G for the reaction, so .
8). The value of εat high and low temperatures are used in eq.(4) to
determine a value for S .
Data sheet
T℃ 10 20 25 30 40 50 60 70 80
ε(mV)
9). Plot εversus T, and determine S from the slope of the line,
/ T , and eq.(4).
10). Repeat procedure 1 to 9 again to make sure the error is within 10%.
Discussion
1. Calculate the thermodynamic values, S , G , and H from
thermodynamics experiment.
2. Write a thoughtful and reasonably complete list of the important
principles and viewpoints you have gained from this experiment.
27
EXPERIMENT 7
Temperature Dependence of Equilibrium
Objectives
To familiarize thermodynamic concept
To study the temperature effect on the thermodynamic
equilibrium and equilibrium constant
Principle
One of the most useful results of thermodynamic analysis is the
expression for the temperature dependence of the equilibrium
constant:
K2 H 1 1
ln ( )
K1 R T2 T1
where K2 and K1 are the equilibrium constants at the temperatures
T1 and T2, H is the standard enthalpy change for the reaction,
and R is the gas constant. This valuable equation applies not only to
chemical equilibria; it also forms the basis for understanding the
temperature dependence of vapor pressure, boiling point elevation,
and freezing point depression phenomena.
We relate the equilibrium constant to the standard free energy
change, E , the standard entropy change, S , and the standard
enthalpy change
E H TS
E RT ln K
we rewrite into
28
H S
ln K
RT R
This equation says that ln K is a linear function of 1/T if H and
S are independent of temperature.
The reaction of dissolution of naphthalene in diphenylamine can
be written as
naphthalene(pure solid) naphthalene (solution, concentration
X)
the equilibrium constant is K=X, where X is the mole fraction of
naphthalene in a solution which is in equilibrium with pure solid
naphthalene. The temperature dependence of the equilibrium can be
found by weighing out naphthalene and diphenylamine to form
mixtures of known composition, melting them in a water bath, and
determining the temperature at which naphthalene precipitates by
the cooling curve method. A plot of log X as a function of the 1/T
yields a straight line of slope H / 2.3R . H is the enthalpy of
fusion of naphthalene and can be determined by this way.
If the reaction is written as
pure solid naphthalene pure liquid solution
the enthalpy change of the first step is the enthalpy of fusion of
naphthalene, while the enthalpy change of the second step is zero,
since the solution is ideal.
Note: for such an experiment we require a reaction which involves
“ideal” reagents, and whose equilibrium constant can be determined
rapidly at several temperatures without using thermostat.
29
Reagent
Naphthalene, Diphenylamine
Procedure
1. Prepare the pure naphthalene (solvent) 20 g ±0.10 g .
2. Prepare six solutions of diphenylamine (solute) in naphthalene
in the range of 1 to 0.2 mole fraction naphthalene. The addition
of successive solute to the same amount of solvent makes the
naphthalene mole fraction is about the value.
3. Plot each cooling curves, temperature vs. time, (temperature
readings every 30 seconds).
4. The values of log X as a function of 1/T are plotted and the
enthalpy of fusion of naphthalene is evaluated from the slope.
Discussion
1. According to the experiment, describe the difference between
E and E .
2. How does the temperature affect the equilibrium ?
3. By using the molecular interpretation of entropy, compare the
freedom of movement for the molecules in the pure liquid, solid
and solution.
When X increases, how does H change? S change?
30
Experiment 8
pH Titration of Sodium Carbonate in an Impure
Sample
`PRINCIPLE.
The impure sample is titrated with standard HCl using a
potentiometric (pH) end point measured with a pH meter using a pH glass
electrode-saturated calomel reference electrode combination. The
end-point breaks are compared with indicator color changes.
EQUATIONS. CO 3 2 H HCO 3 (phenolphthalein end point)
HCO 3 H H 2O CO 2 (methy1 orange end point)
Note that between the first and second end points, a gradual decrease in
pH due to the HCO 3 /CO 2 buffer system will occur. This will give a poor
visual end point, unless the buffer couple is destroyed. In practice, the
visual titration used for standardization is continued until the
methylorange end point is reached, at which time, the solution is gently
boiled to remove the CO 2 , leaving only the remaining HCO 3 ,which is
then titrated to completion.
SOLUTIONS AND CHEMICAL REQUIRED.
Provided. 0.2% phenolphthalein in 95% ethanol, 0.1% methylorange
in water, primary standard pH 7 buffer. Standard 0.1 M HCl solution.
THINGS TO DO BEFORE THE DAY OF THE EXPERIMENT.
Prepare and standardize the HCl solution. This will require prior
drying of primary standard Na2CO 3 .
31
Obtain the unknown sample from your instructor and dry for at least
two hours at.160 C .Cool at least 30 minutes in a desiccator before
weighing.
PROCEDURE
The glass electrode to be used for pH measurements should have been
soaked and stored in distilled water for at least one day prior to its use.
Always store the electrode in water when not in use. Calibrate the pH
meter as described by your instructor, using the pH 7 standard buffer.
This will consist essentially of adjusting the meter to read pH 7.00 with
the electrodes immersed in the buffer solution. If only small quantities of
buffer are used, it would be better to discard it rather than to chance
contamination of the entire supply.
1. Trial Titration. The purpose of this titration is to locate quickly and
approximately the two end points. Weigh accurately by difference a
dried sample of unknown sample (0.2-0.3g) and add it to a 400-ml
beaker containing a magnetic stirring bar. Add approximately 50 ml
of water and a few drops of phenolphthalein indicator. The indicators
are for the purpose of making a comparison between the
potentiometric end points and the indicator color changes. Place the
beaker on a magnetic stirrer, immerse the electrodes and start the
stirrer, being careful not to touch the electrodes to the stirring bar.
Titrate with standard HCl, taking readings about every 2 ml. After the
phenolphthalein color disappears, add a few drops of methy1 orange
indicator and titrate at 2-ml increments until the second end point is
reached. Add a few increments beyond the end point. The correct
color for the second end point can be determined by comparison with
the color of a few drops of the indicator in a solution of 0.20g
potassium acid phthalate in 100 ml of water. Plot a curve of pH versus
volume of HCl and locate the approximate end points.
2. Final titration. Weigh accurately another sample of the unknown
and titrate as before, but make pH readings every 5 ml to within 3ml
of each end point (both sides of end point). Then make readings at
1-ml interval within 1 ml of the end point. Near the end point, take
readings as quickly as possible because the pH will tend to drift as
CO 2 escapes from the solution. Note and record the points at which
32
the indicators change color.
3. Plot a curve of pH (on the ordinate) versus volume of HCl (on the
abscissa) and indicate of this curve the range in which the indicators
change color. Determine the end point from the second inflection
point of the curve. Repeat the titration on two more portions of the
unknown. Be sure to rinse the electrodes between titrations.
4. CALCULATIONS. Calculate and report the percent of Na2CO 3
% Na2CO 3 =
M HCl ( mmoles / ml ) ml HCl ( mmoles Na2CO 3 / mmoles HCl ) f .w .Na2CO3 ( mg / mmole )
mg sample
100%
Formula weight Na2CO 3 =105.99;
33
EXPERIMENT 9
Determination of the order and Rate constant of
a chemical Reaction
Objectives
To study the kinetics and reaction mechanism
To determine the reaction order and rate constant
To study the rate of the reaction at different temperatures in
order to find the activation energy
Principle
The rate at which a chemical reaction occurs depends on several
factors: the nature of the reaction, the concentrations of the reactants, the
temperature, and the presence of possible catalysts. All of these factors
can markedly influence the observed rate of reaction.
Some reactions at a given temperature are very slow indeed; Other
reactions are essentially instantaneous; the precipitation of silver chloride
when solutions containing silver ions and chloride ions are mixed and the
formation of water when acidic and basic solutions are mixed are
examples of extremely rapid reactions. In this experiment we will study a
reaction which, in the vicinity of room temperature, proceeds at a
moderate, relatively easily measured rate.
For a given reaction, the rate typically increases with an increase in
the concentration of any reactant. For the reaction
aA + bB cC
The rate can usually be expressed by the equation
rate = k(A) m ( B) n (1)
where m and n are generally, but not always, integers, 0, 1, 2, or possibly
34
3; (A) and (B) are the concentrations of A and B ( ordinarily in moles per
liter ); and k is the rate constant of the reaction, which makes the relation
quantitatively correct. The numbers m and n are called the orders of the
reaction with respect to A and B. If m is l the reaction is said to be first
order with respect to the reactant A. If n is 2 the reaction is second order
with respect to reactant B. The overall order is the sum of m and n. In this
example the reaction would be third order overall.
The rate of a reaction is significantly dependent on the temperature at
which the reaction occurs. An increase in temperature increases the rate,
an often cited rule being that a 10°C rise in temperature will double the
rate. This rule is only approximately correct; nevertheless, it is clear that
an arise of temperature of 100°C could change the rate of a reaction
appreciably.
As with the concentration, there is a quantitative relation between
reaction rate and temperature, but here the relation is somewhat more
complicated. This relation is based on the idea that in order to react, the
reactant species must have a certain minimum amount of energy present
at the time the reactants collide in the reaction step; this amount of energy,
which is typically provided by the kinetic energy of motion of the species
present, is called the activation energy for the reaction. The equation
relating the rate constant k to the absolute temperature T and the
activation energy Ea is
Ea
log10 k constant (2)
2.30 RT
where R is the gas constant ( 8.31 joules/mole K for Ea in joule per
mole ). By measuring k at different temperatures we can determine
graphically the activation energy for a reaction.
In this experiment we will study the kinetic of the reaction between
iodine and acetone (see Appendix):
O O
|| ||
CH3 - C - CH3 ( aq ) + I 2 ( aq ) CH3 - C - CH 2I( aq ) + H + ( aq ) + I- ( aq )
The rate of this reaction is found to depend on the concentration of
hydrogen ion in the solution as well as presumably on the concentrations
of the two reactants. By Eq. (1), the rate law for this reaction is
35
rate = k ( actone)m ( I 2 )n ( H + ) P (3)
where m, n, and p are the orders of the reaction with respect to acetone,
iodine, and hydrogen ion respectively, and k is the rate constant for the
reaction.
The rate of this reaction can be expressed as the ( small ) change in
the concentration of I 2 , (I2 ) , divided by the time interval t required
for the change:
- (I 2 )
rate = (4)
t
The minus sign is to make the rate positive [ (I2 ) is negative].
Ordinarily, since rate varies as the concentrations of the reactants
according to Equation 3, in a rate study it would be necessary to measure,
directly or indirectly, the concentration of each reactant as a function of
time; the rate would typically vary markedly with time, decreasing to
very low values as the concentration of at least one reactant become very
low. This makes reaction rate studies relatively difficult to carry out and
introduces mathematical complexities that are difficult for beginning
students to understand.
The iodination of acetone is a rather typical reaction, in that it can be
easily investigated experimentally. First of all, iodine has color, so that
one can readily follow changes in iodine concentration visually. A second
and very important characteristic of this reaction is that it turns out to be
zero order in I 2 concentration. This means (see Equation 3 ) that the rate
of the reaction does not depend on ( I 2 ) at all; ( I 2 )°=1, no matter what the
value of ( I 2 ) is, as long as it is not itself zero.
Since the rate of the reaction does not depend on ( I 2 ), we can study
the rate by simply making I 2 the limiting reagent present in a large
36
excess of acetone and H + ion. We then measure the time required for a
known initial concentration of I 2 to be completely used up. If both
acetone and H + are present at much higher concentrations than that of
I 2 , their concentrations will not change appreciably during the course of
the reaction, and the rate will remain, by Equation 3, effectively constant
until all the iodine is gone, at which time the reaction will stop. Under
such circumstances, if it takes t seconds for the color of a solution having
an initial concentration of I 2 equal to ( I 2 ) . To disappear, the rate of the
reaction, by Equation 4, would be
-(I 2 ) (I )
rate = = 2
t t
(5)
Although the rate of the reaction is constant during its course under
the conditions we have set up, we can vary it by changing the initial
concentrations of acetone and H + ion. If, for example, we should double
the initial concentration of acetone over that in Mixture 1, keeping ( H + )
and ( I 2 ) at the same values they had previously, then the rate of Mixture 2
would, according to Equation 3, be different from that in Mixture 1:
rate 2 = k ( 2A )m ( I 2 ) ( H + )p (6a)
rate 1 = k (A )m ( I 2 ) ( H + )p
a (6b)
Dividing the first equation by the second, we see that the k’s cancel,
as do the terms in the iodine and hydrogen ion concentrations, since they
have the same values in both reactions, and we obtain simply
rate 2 ( 2A )m 2A m
= m
= ( ) = 2m (6)
rate 1 ( A) A
37
Having measured both rate 2 and rate 1 by Equation 5, we can find
their ratio, which must be equal to 2 m . We can then solve for m either by
inspection or using logarithms and so find order of the reaction with
respect to acetone.
By a similar procedure we can measure the order of the fact that the
reaction is zero order with respect to I 2 . Having found the order with
respect to each reactant, we can then evaluate k, the rate constant for the
reaction.
The determination of the orders m and p, the confirmation of the fact
that n, the order with respect to I 2 , equals zero, and the evaluation of the
rate constant k for the reaction at room temperature comprise your
assignment in this experiment. You will be provided with standard
solutions of acetone, iodine, and hydrogen ion, and with the composition
of one solution that will give a reasonable rate. The rest of the planning
and the execution of the experiment will be your responsibility.
Apparatus and Reagents
watch glass
4 M Acetone-water
1 M HCl-water
0.005 M I2-water
Be careful not to spill the iodine solution on your hands or clothes.
Procedure
1. Select two test tubes and fill with distilled water. (When you
view them down the tubes against a white background, they
should have identical color. Make mark 3/4 equal height for
each test tube.)
38
2. Draw 50 ml of each of the following solutions into clean dry 100
ml beakers, one solution to a beaker: 4M acetone-water, 1M
HCl-water, and 0.005M I 2 -water ( Note that I 2 is difficult to be
dissolved in water. Add some KI until I 2 is totally dissolved.) .
Cover each beaker with a watch glass.
3. Prepare Mixture Solutions : (keep the total volume at 50 ml)
see table 1. Cover each beaker with a watch glass.
4. The temperature should be kept within about a degree in each
run.
5. Record the time (ti) to one second, when you pour the iodine
solution into the Erlenmeyer flask and quickly swirl the flask to
thoroughly mix the reagents. (The reaction mixture will appear
yellow because of the presence of the iodine, and the color will
fade slowly as the iodine reacts with the acetone.)
Table 1
Mixture 4M 1M H2O 0.005M t t f t i Rate=
HCl (ml) I2 (ml) (I 2 )
acetone( (ml) t
ml)
1 10 10 10 20
2 20 10 10 10
3 10 20 10 10
4 10 10 5 25
6. For each mixture, fill one test tube 3/4 full with the reaction
mixture, and fill the other test tube to the same depth with
distilled water (as reference color). Record the time (tf) when
the color of the iodine just disappears and measure the
temperature of the mixture in the test tube.
7. Repeat the experiment, using as a reference the reacted
solution instead of distilled water. The amount of time required
in the two runs should agree within about 20 seconds.
(I 2 )
8. Calculate Rate = for the reaction mixture. Since the
t
reaction is zero order in I 2 , and since both acetone and H +
39
ion are present in great excess, the rate is constant throughout
the reaction and the concentrations of both acetone and H+
remain essentially at their initial values in the reaction mixture.
9. Carry out the reaction twice with each mixture; the times should
not differ by more than about 15 seconds.
10. Calculate the order of the reaction with respect to acetone, H + ,
I2 by using a relation similar to Equation 6.
11. Evaluate the rate constant k for the reaction from the rate and
concentration data in each of the mixtures. If the temperatures
at which the reactions were run are all equal to within a degree
or two, k should be about the same for each mixture.
12. (Optional) As a final reaction, make up a mixture using reactant
volumes that you did not use in any previous experiments.
Using Equation 3, the values of concentrations in the mixtures,
the orders, and the rate constant you calculated from your
experimental data, predict how long it will take for the I 2 color
to disappear from your mixture. Measure the time for the
reaction and compare it with your prediction. This part is also to
study the rate of this reaction at different temperatures in order
to find its activation energy. The procedure here would be to
study the rate of reaction in one of the mixtures, at room
temperature and at two other temperatures, one above and one
below room temperature. Knowing the rates, and hence the k’s,
at the three temperatures, you can the find Ea, the energy of
activation for the reaction, by plotting log k vs. 1/T. the slope of
the resultant straight line, by Equation 2, must be -Ea/2.30R.
Discussion
1. What changes in composition you might make to decrease the
time and increase the rate of reaction?
40
2. How could you change the composition to allow you to
determine how the rate depends upon acetone concentration ?
3. How do you keep that the concentrations of H + and I 2 are
the same as in the experiment ?
4. Calculate the reaction order and rate constant for the reaction.
Appendix: Enolization ( acid-catalyzed halogenation of ketones)
Rate-determining reactions racemization. The rate determining
reaction here is the formation of the enol, which involves two steps: rapid,
reversible protonation (step 1) of the carbonyl oxygen, followed by the
slow loss of an -hydrogen (step 2).
(1) CH3-C-CH3 + H : B CH3-C-CH3 + : B fast
:O: +OH
(2) CH3-C-CH3 + : B CH3-C= CH2 + H : B slow
+OH OH
(3) CH3-C=CH2 + X - X CH3-C-CH2X + X- fast
OH +OH
(4) CH3-C-CH2X + : B CH3-C-CH2X + H : B fast
+OH O
41
Experiment 10
Titration of Calcium and Magnesium in Milk
with EDTA
Complexometric titrations using EDTA are invariably encountered in
the undergraduate analytical chemistry laboratory.
The most common experiment (where samples other than typical
commercial unknowns are used) is determination of water hardness.
Although it is interesting for students to collect his own water sample,
exciting samples of complicating waters are not always available, and the
analysis of calcium and magnesium in ordinary water is not so
challenging.
A sample which we have found to be interesting and somewhat
challenging to students is milk.
Principle
The titration of calcium and magnesium in milk and dairy products is
seriously complicated by the presence of proteins and orthophosphate
ions, which precipitate calcium and magnesium at the high pH required
for titration with EDTA. These interferences can be removed by
precipitation with potassium metastannate, yielding a clear filtrate in
which calcium and magnesium may be determined. But this filtration is
time consuming and because the precipitate occupied some volume in the
flask, the amount (quantity) of the sample must be corrected. Another
means of removing phosphate ions is anion exchange, but application of
this technique to milk requires prior removal of protein by precipitation
or ashing. Elution of the ion exchange column is also time consuming,
and yields a very dilute sample for titration, leading to indistinct end
points.
Kamal has shown (1) that addition of excess Na 2 EDTA to a
neutral or slightly acidic solution containing calcium, magnesium, and
phosphate ions causes chelation of calcium and magnesium, preventing
42
their precipitation with phosphate when the solution is made alkaline. By
back-titrating the excess EDTA at pH 10, the EDTA complexed with the
sum of calcium and magnesium can be calculated. If the titration is
performed at pH 12 or higher, magnesium is released from its complex
with EDTA and precipitates phosphate and / or hydroxide. The
back-titration in this instance determines the EDTA complexed with
calcium alone. These data, together with a blank determination, allow
calculation of both calcium and magnesium in the sample. Kamal applied
this method to milk and found no interference from the proteins present.
Apparatus and Reagents
1. Magnesium turnings 0.2g
2. Calcium carbonate 2.5g
3. HCl
4. EDTA 15g
5. ammonium chloride 67.5g
6. ammonium hydroxide 570ml
7. KOH or NaOH 0.5N
8. Calmagite
9. Hydroxy Naphthol blue
10. Pipet
11. 125ml flask
12. Buret
Experimental
Reagents
Standard Magnesium Solution. Dissolve 0.2 g ( accurately weighed )
of reagent grade magnesium turnings in the minimum amount of dilute
HCl ( about 3 ml of 6 N ) and dilute to 1-l. This solution contains about
0.2 mg Mg per ml.
Standard Calcium Solution. Dissolve 2.5 g ( accurately weighed ) of
previously dried reagent grade calcium carbonate in dilute HCl, and
dilute to 1-l. This solution contains about 1 mg Ca per ml.
EDTA Solution. Dissolve 15 g of reagent grade disodium EDTA in
water and dilute to 1 l. The solution will be standardized against the
standard Mg and Ca solutions. The titer of this solution is about 1.6 mg
43
Ca per ml and just under 1 mg Mg per ml.
Ammonia Buffer, pH 10. Dissolve 67.5 g of ammonium chloride in
200 ml distilled water, add 570 ml of reagent grade concentrated
ammonium hydroxide, and dilute to 1-l.
Base for High pH Titration. 0.5 N KOH or NaOH.
Magnesium Indicator. Calmagite; either a dry preparation or a
0.05% solution in water of the pure material.
Calcium Indicator. Hydroxy Naphthol Blue
Procedure
Calcium.
Pipet 5 ml of milk and 10 ml of EDTA into a 125-ml erlenmeyer
flask, and mix thoroughly. Add 15 ml of 0.5 N KOH to the sample to
adjust the pH to about 13, and add 200-300 mg calcium indicator. The
color of the solution will change from white to light blue. The excess
EDTA is back-titrated with standard calcium solution until the color
changes from blue to violet. Dilution with water should be avoided as
much as possible. The volume of EDTA reacted with calcium is
determined from the difference between this titration and a similar
titration where water is substituted for the milk. The amount of calcium in
the sample is calculated by multiplying the volume of EDTA complexed
with calcium by the calcium titer of the EDTA, which can be determined
from the blank titration.
Magnesium.
Pipet 5 ml of milk and 10 ml of EDTA into a 125-ml erlenmeyer
flask and mix thoroughly. Add 10 ml of ammonia buffer and 200 mg of
magnesium indicator ( or an appropriate volume of indicator solution ).
The color of the solution will change from white to light blue. To
minimize loss of ammonia, do not add the buffer and indicator until just
before titrating. Titrate the sample with standard magnesium solution
until the color changes from blue to red-violet, avoiding dilution with
water. From the difference between this titration and a blank titration, run
with water instead of milk, determine the volume of EDTA complexed
44
with the sum of calcium and magnesium. Subtract the volume of EDTA
which was found to be complexed with calcium in the previous
experiment, and multiply the result by the magnesium titer of the EDTA
to obtain the amount of magnesium in the sample.
Reference
Kamal, T. H., J. Agr. Food Chem., 8, 156(1960).
Result and Discussion
45
Experiment 11
Colorimetric Determination of Iron in Vitamin
Supplement Tablets
The basic concepts of instrumental analysis are introduced through
modification of a standard procedure for the colorimetric determination
of iron. By utilizing standard brand name products ( e.g. Geritol and
One-A-Day Plus Iron ) student interest is kept high throughout the
experiment. The use of spectrophotometers ( e.g. Bausch and Lomb
Spectronic 20’s ) appeared to be more suitable for analytical chemistry
laboratories.
Principle
In spectrometric methods, the sample solution absorbs
electromagnetic radiation from an appropriate source, and the amount
absorbed is related to the concentration of the analyte in the solution. In
other words, spectrometry is based on the absorption of photons by the
analyte. The color of an object we see is due to the wavelengths
transmitted or reflected. The other wavelengths are absorbed.
All biochemicals absorb energy from at least one region of the
spectrum of electromagnetic radiation. The energies at which absorption
occurs depend on the available electronic, vibrational and rotational
energy levels of the molecule. When absorption is from the UV /visible
region of the spectrum (200-700nm), transitions occur between electronic
energy levels. Molecules absorb energy only when the incident photon
has precisely equal to the difference in energy between two allowed states,
the photon promoting the transition of an electron from the higher energy
state. Before another photon can be absorbed, the excited state must lose
this energy and revert to the ground state.
The measurement of the absorption of radiation by chemical species
is known as spectrophotometry. Spectrophotometers are constructed so
that the sample to be studied can be irradiated with light ( or other
radiation ) of know wavelength and intensity. The wavelength can be
46
varied continuously by the operator (or automatically) and the amount of
radiation absorbed or transmitted by the sample determined for each
wavelength used. In this way, it is possible to learn which wavelength of
radiation is absorbed by the sample and how effective the species in the
sample are in absorbing a particular wavelength. From this information,
an absorption spectrum for a species can be obtained and used to identify
the species in unknown samples. In many cases, the amount of a
substance present in a sample can be determined by spectrophotometry.
Apparatus
1. Colorimeters (Bausch and Lomb Spectronic 20 or equivalent )
for recording absorption at 508 nm.
2. Colorimeter tubes ( 24 for a class of 20 students ) and racks.
3. 100 ml-volumetric flasks --- 2.
4. One-and two-ml transfer pipets ( one of each per class ).
5. Five-and ten-ml transfer pipets ( one of each per pair of
students ).
6. Fifty-milliliter burets ( three per class ).
7. PHydrion paper, range 1-5.5 ( one roll ).
Reagents
1) Hydroquinone solution, 1% in water ( freshly prepared and
stored in an amber bottle ).
2) Sodium citrate solution, 25g/l in water.
3) o-Phenanthroline solution, 0.25% in water containing 10%
alcohol ( freshly prepared and stored in an amber bottle ).
4) Standard iron solution, 0.04 mg Fe/ml. Dissolve 0.281 g of
reagent grade FeSO4 ( NH4 )2 SO4 6H2 O in 50 ml of
water containing 1 ml of conc. H 2 SO4 . Transfer to a 1-l
volumetric flask and dilute to the mark with water.
Procedures
The spectrophotometer should be turned on at least 20 minutes
before any measurements are taken.
47
Calibration Curve
Pipet a 10-ml aliquot of the standard iron solution into a beaker and
test the pH using the pHydrion paper. Add sodium citrate solution
dropwise until a pH of about 3.5 is obtained as determined by the
pHydrion paper color (about 30 drops (1.5 ml dispensed from a buret ) of
sodium citrate should suffice). Now pipet a second 10-ml aliquot of the
iron solution into a 100-ml volumetric flask. Add the same number of
drops of sodium citrate followed by 2 ml hydroquinone solution and 3 ml
of o-phenanthroline (dispensed from burets). Dilute to the mark with
distilled water and mix well. This solution contains 4 ppm (4 mg/l) of
iron. Prepare three other standard solutions (2.0, 0.8, 0.4 ppm) by
pipetting 5, 2 and 1 ml, respectively, of the standard iron solution into
100-ml volumetric flasks, adding the appropriate amount of sodium
citrate solution to each, and proceeding as described above.
Allow the standard solutions to stand 10 min and measure the
absorbance of each solution at 508 nm against a blank containing all of
the reagents except the iron solution.
The absorbance values are obtained for each colorimeter, and those
values prominently displayed. Each student should then construct a
calibration curve by plotting absorbance versus ppm Fe.
Vitamin Tablet Iron Content
Fe content Fe content Std. Deviation
( Label ) a ( Relative )
Vitamin Brand ( This Expt )
( mg ) ( mg ) (%)
Geritol 50 53.1 2.84
One-A-Day Plus Iron 18 19.1 4.48
Unicap M 10 10.2 0.56
One-A-Day ( Regular ) 0 <1
a
Average value for four determinations.
Sample preparation and Determination
One tablet of the brand vitamin to be analyzed to be analyzed is
placed in a 100-ml beaker and heated to a slow boil with 25 ml 6 N HCl
for 15 min. The mixture is then diluted slightly with water and filtered
48
while hot through #40 paper directly into a 100-ml volumetric flask. After
washing the residue with hot water the filtrate is allowed to cool and
diluted to the mark. A 5-ml aliquot (10 ml if the package label indicates
the tablet contains less than 15 mg Fe ) is then pipetted into a 100-ml
volumetric flask and diluted to volume.
A 10-ml aliquot of this new solution is pipetted into a beaker and
the pH adjustment performed as described for the calibration curve. A
second aliquot is then pipetted into a 100-ml volumetric flask, the
required amount of sodium citrate added (about 3.5 or 7.0 ml, depending
on the aliquot taken), 2 ml hydroquinone and 3 ml o-phenanthroline
added, and the solution diluted to the mark. After 10 min, the absorbance
is determined.
The concentration of iron is determined by interpolation of the
sample absorbance value on the standard calibration curve. A simple
calculation involving dilutions used for the particular sample allows the
amount of iron in the original tablet to be readily determined.
Result and Discussion
49
Experiment 12
Liquid Chromatography ---Ion-Exchange
Principle
The principle of ion-exchange chromatography is that charged
molecules adsorb to ion exchangers reversibly so that molecules can be
bound or eluted by changing the ionic environment. Separation on ion
exchangers is usually accomplished in two stages: first, the substances to
be separated are bound to the exchanger, using conditions that give stable
and tight binding; then the column is eluted with buffers of different pH,
ionic strength, or composition and the components of the buffer compete
with the bound material for the binding sites.
In this experiment, Iron ( III ), Copper, and Nickel may be separated
by passing a solution in 8M hydrochloric acid through the anion
exchanger and rinsing with 8M acid; Nickel passes through and the other
ions are held on the resin. Then 3M acid is passed through to elute
Copper, and finally distilled water to elute Iron.
Properties of Ion Exchangers
An ion exchanger is usually a three-dimensional network or matrix
that contains covalently linked charged groups. If a group is negatively
charged, it will exchange positive ions and is a cation exchanger. A
2
typical group used in cation exchangers is the sulfonic group, SO 3 . If
an H + is bound to the group, the exchanger is said to be in the acid
form: it can for exchange one H + for one Na or two H + for one
+
Ca 2+ . The sulfonic acid group is called a strongly acidic cation
exchanger. Other commonly used groups are phenolic hydroxyl and
carboxyl, both weakly acidic cation exchangers. If the charged group is
positive, for example, a quaternary amino group, it is a strongly basic
anion exchanger. The most common weakly basic anion exchangers are
aromatic or amino groups.
The matrix can be made of various materials. Commonly used
materials are dextran, cellulose, agarose and polystyrene or polyphenolic
50
resins. Usually, polystyrene (Dowex resin) is used for isolation and
separation of small molecules and cellulose ( DEAE, CM ) for
macromolecules.
The total capacity of an ion exchanger measures its ability to take
up exchangeable ions and is usually expressed as milliequivalents of
exchangeable groups per milligram of dry weight. This number is
supplied by the manufacturer and is important because, if the capacity is
exceeded ions will pass through the column without binding.
Material:
(1) Amberlite CG-400; 100-200 mesh
(2) Chromatography column (1.6 cm 15 cm)
(3) 8 M HCl, 3 M HCl, 1 M HCl
(4) 0.1 M Fe3+ , Cu2+ , Ni 2+ in 9 M HCl solution.
(5) Glass wool
(6) Na2SO4
Important note: Throughout this entire chromatographic analysis, never
let the fluid level in the column drop below that of the resin. If this
happens, channels will develop in the resin bed and the separation
efficiency will be vastly decreased. If the fluid level should drop below
the resin, the column must be emptied and repacked.
Preparing the stock solution
A solution containing 1.5 mg/ml of each ion is prepared in 9M HCl by the
instructor. Ions used in this study are Ni(II), Cu(II), and Fe(III), dissolved
in the medium as their nitrate salts. Each group will need 0.5 ml of this
solution.
Preparing the Column
Place approximately 5g of dry resin in a 100 ml beaker. Add 20 ml
distilled water and swirl the contents, Add to this 5 ml 8M HCl and pour
the slurry into the column slowly but all at once. Rinse the beaker with 2
or 3 ml 8M HCl and pour this immediately onto the column. Tap the
column gently to settle the resin, which should form a column about 13
51
cm high. During all these operations the valve at the bottom of the
column should be open. Insert a small glass wool plug at the top of the
column and allow the solution in the column to flow though the resin to
within 1 cm above the plug. (Ensure that there are no bubbles trapped in
the column)
Now activate the column by adding 15 ml more 8M HCl and draining it
to within 1 cm of the plug. Discard this eluant. The resin at this point may
have changed from yellow to brown.
Separating the Ions
Add 0.5 ml of the solution containing the chloro complexes to the column.
Drain the liquid to within 1 cm of the plug. Add 1 ml 8M HCl to rinse
any sample off the walls of the column and again drain too within 1 cm of
the plug. Add 15 ml of 8M HCl and begin to collect eluant fractions of 4
ml each in numbered test tubes at a flow rate of 4 ml per min. (about 1
drop/sec).
When the 8M HCl is within 1 cm of the plug, add 30 ml 3M HCl and
continue collecting fractions of 4 ml each at the same flow rate. When
this solution is within 1 cm of the plug, continue the elution with 30 ml
distilled water. When this is gone, you should have collected 20 fractions,
each with a volume of 4 ml.
Reclaiming the Resin
Before proceeding with the analysis of the fractions, reclaim the resin by
running 10 ml 1M Na2SO4 through the column, followed by 30 ml
distilled water. Discard this eluant. Remove the resin from the column
and store it in an appropriately labeled bottle for future use.
Result and Discussion
Outside Reading:
Analyzing the Eluant fractions, Spot tests
Copper
52
Transfer 6 to 10 drops of the eluant to a clean tube, Add 15M ammonia
until the solution is just barely basic to litmus. Adjust the pH of the
solution to slightly acidic to litmus by adding 2 drops of 6M acetic acid.
Add several drops of 10% hydroxylamine hydrochloride solution
followed by several drops of cuproine reagent (a saturated solution of 2,
2’-biquinoline in isoamyl alcohol). The appearance of a rose to purple
color in the top layer of liquid in the tube confirms the presence of
copper(II). Record the tube numbers that give a positive test for copper.
Iron
Transfer 6 to 10 drops of the eluant to a clean tube. Add 5 drops of 1M
KSCN. The formation of the deep red FeSCN+2 complex confirms the
presence of iron, Since this test is extremely sensitive, a pale pink color
indicates a possible trace contamination rather than the true presence of
iron. Record the tube numbers which give a positive test for iron.
Nickel
Transfer 6 to 10 drops of the eluant to a clean test tube. Neutralize the
solution by adding drops of 15M NH3 until the solution is just basic to
litmus. Add several drops of 1% dimethylglyoxime reagent and mix the
contents. The formation of a brick red precipitate confirms the presence
of nickel. The precipitate may take as long as 5 min to form. If no
precipitate appears after 5 min, add 1 drop of ammonia and wait another 5
min. Record the tube numbers which give a positive test for nickel.
Spectrophotometric analysis of elution profiles
Copper
Add to 1 ml aliquots of the eluant the following solutions in the indicated
order:
(1) 2 ml 10% hydroxylamine hydrochloride solution;
(2) 2 ml 10% tartaric acid solution;
(3) adding 6M NH3 dropwise adjust the pH to 5-6 using narrow-range
pH test paper,
(4) add 5 ml cuproine reagent.
Shake each solution for 1-2min. Set aside and allow the layers to separate.
Carefully remove the top layer and place it in a cuvet to make the
spectrophotometric measurement.
Record the absorbance of each solution at 545 nm using isoamyl alcohol
53
as the blank.
Iron
Add to 2 ml aliquots of the eluant the following in the order given:
(1) 1ml 10% hydroxylamine hydrochloride solution;
(2) 1ml 0.5% phenanthroline solution,
(3) using 2M sodium acetate adjust the pH to 3-4.
Allow the solutions to sit for approximately 1 hr.
Dilute the solutions by the addition of 25 ml distilled water. Record the
absorbance of each solution at 505 nm using distilled water as a blank.
Nickel
Add to 1 ml aliquots in 125 ml flasks the following:
(1) 0.5 ml saturated bromine water,
(2) 2 ml 15M ammonia,
(3) 20 ml 95% ethanol,
(4) 10 ml 1% dimethylglyoxime solution,
(5) 7 ml distilled water.
Record the absorbance of each solution at 450 nm using distilled water as
a blank. (Because this test mixture is unstable, the absorbance of these
solutions must be measured within a half hour of mixing.)
54
55