Cyano Type Blue Printing
Cyanotype is a photographic printing process that produces a cyan-blue print. Engineers used the process well
into the 20th century as a simple and low-cost process to produce copies of drawings, referred to as blueprints.
Invented by Sir John Herschel in 1841, this simple process produces a continuous tone
image of Prussian Blue using a sensitizing solution of ferric ammonium citrate and
potassium ferricyanide. These iron salts, when exposed to natural or artificial ultraviolet
light, are reduced to their ferrous state, producing a high contrast blue image when
oxidised. The process was eminently suited to its traditional role in reproducing technical
drawings, its most common use in engineering and architecture until the advent of modern
photocopiers. However, it was a versatile process, and was used throughout the 19th
century from Anna Atkins’ photograms of plants and seaweed for her books on botany
(1843–55) to Henri LeSecq’s still life studies of the 1850s. Photographers at the end of the century used
cyanotype paper for proofing negatives.
Procedure:
1. Immerse pieces of bond paper one by one in the sensitizing solution and keep them immersed for 4-5 minutes.
2. Remove the wet pieces of paper and place them between sheets of filter paper. This should be done as quickly
as possible and in a partially closed locker. Dry it for 10-15 minutes.
3. After the paper has dried, place an opaque object on top of the sensitizing paper, compress it between sheets of
glass and expose to sunlight for 4-6 minutes.
4. Make 3-4 exposures, varying the time of exposure to optimize the best condition.
5. After the exposure, dip the paper into 0.1M ferric cyanide. It is important that the paper is immersed all at once,
otherwise lines will appear on the blue field of the paper.
6. Remove the paper and dip it in 0.3M potassium dichromate solution for one minute. Afterwards, wash the paper
first with 0.1M HCl and then tap water and dry.
7. Paste all the images in your lab notebook with their exposure times.
Synthesis of Dibenzalacetone
Aldol condensation is an important route of organic synthesis because it provides an efficient way to
form carbon-carbon bond. In this condensation, an enol or enolate ion reacts with a c arbonyl
compound to from a β-hydroxy ketone or β-hydroxy aldehyde, which is followed by dehydration. The
reaction is used to manufacture solvents such as isophorone, used in printing inks, lacquers,
adhesives and many other products. It is also used in the manufacture of α, β- unsaturated ketones
known as chalcones. This condensation is generally used to create plasticizers which ae used to
convert rigid plastic polyvinyl chloride into a soft, flexible elastic material. In today’s experiment, you
will utilize this reaction to prepare dibenzalacetone. It is a common ingredient in sunscreens since it
absorbs UV light. Dibenzalacetone was first prepared in 1881 by Claisen and Claparede. It is used as
ligand for making organometallic complexes which are used as catalysts in coupling reactions. For
example, Pd-DBA.
O
C
O
H
NaOH
O
+
H 3C
+ 2H2O
CH3
Procedure:
1. Take 0.8 mL of benzaldehyde in a small conical flask.
2. Add 0.3 mL of acetone and 5 mL of ethanol to the above. Add 2 mL of NaOH solution to the
mixture of benzaldehyde, acetone and ethanol mixture.
3. Mix and swirl the reaction mixture occasionally for 15 minutes.
4. A yellow flocculent precipitate should form. Filter the product using a Buchner funnel on a
vacuum filtration set up.
5. Wash the obtained solid with 10-15 mL of water, followed by a washing with 8-10 mL of chilled
ethanol.
Recrystallization
1. Dissolve the crude product in 2 mL of ethyl acetate in a beaker and heat on water bath till you
get a clear solution.
2. Allow the solution to cool slowly to room temperature and then cool it in an ice bath. Pure
dibenzalacetone crystallizes.
3. Filter the pure product and allow to dry. Weigh the product and report the percentage yield.
4. Determine the melting point of the product.
Synthesis of Aspirin
Aspirin (acetylsalicylic acid) is a synthetic organic derived from salicylic acid. Salicylic acid is a natural product
found in the bark of the willow tree and was used by the ancient Greeks and Native Americans, among others,
to counter fever and pain. However, salicylic acid is bitter and irritates the stomach. In a Bayer laboratory in
Wuppertal, Germany, young scientist Dr. Felix Hoffmann was the first to succeed in synthesizing a chemically
pure and stable form of acetylsalicylic acid (ASA), which becomes the active ingredient in Aspirin. Aspirin is the
most frequently sold pain reliever in the world, has been the subject of a Nobel prize (1982), and has been
termed the ‘wonder drug’ of the century. It is singlehandedly responsible for the foundation and success of
Bayer Pharmaceuticals (2019 revenue: 49 billion US dollars).Acetylation of Salicylic Acid
O
C
OH +
OH
Salicylic Acid
O
O
H3C
O
O
C
catalyst
CH3
Acetic Anhydride
H 3C
O
C
O
OH +
H3C
OH
O
Procedure:
1. Take a petri dish or watch glass and weigh x g of salicylic acid (Mol. Wt. 138.12 g/mol).
2. Transfer the weighed salicylic acid to a dry 150 mL conical flask.
3. Add 2.7 equivalents of acetic anhydride (Mol. Wt. 102.08 g/mol; Density 1.08 g/mL) using a measuring
cylinder to the salicylic acid. Now add 5-6 drops of concentrated sulfuric acid and stir until all salicylic
acid is dissolved.
4. Leave the reaction mixture undisturbed for 15-20 minutes.
5. Add 50 mL of water to the flask and swirl for two minutes and filter.
6. Collect the solid from the filter paper.
Recrystallization
1. Dissolve the crude product in 7 mL of ethanol in a beaker and add 15 mL of distilled water. Heat on
water bath till you get a clear solution.
2. Allow the solution to cool in an ice bath without disturbing. Pure acetylsalicylic acid crystallizes.
3. Filter the pure product and dry it by placing in between sheets of filter paper. Report the percentage
yield.
4. Dissolve a f ew crystals of the dry compound in 0.5 mL methanol and a dd 2 d rops of FeCl3 solution.
Note the color change. Repeat the above test with similar amount of salicylic acid and note the color
change.
5. Determine the melting point of acetylsalicylic acid.
Estimation of Iodine in iodized common salt using iodometry
Iodine is an essential element for life and one of the heaviest elements required by living
organisms. However, around one third of the world’s population
lives in areas of iodine deficiency. The practice of adding iodine to
salt is a safe, easy and effective way of overcoming iodine
deficiency in our diet. Globally two chemical forms of iodine are
used for iodization; Iodates (IO3-) and Iodides (I-). The iodides
degrade more readily in presence of impurities, exposure to
sunlight, moisture and exposure to heat, whereas the iodates
remain stable under extremes of weather and handling. USA uses
potassium iodide (77 mg/Kg) while Germany and India use
potassium iodate (25-20 mg/Kg) for iodine fortification. Today you will use iodometry to
estimate the amount of iodine in a salt sample.
Procedure:
Rapid test to determine the nature of iodizing reagent
Take a pinch of common salt on a watch glass and divide into two parts:
Test for iodate solution
Moisten the first part with 2-3 drops of the given solution (mixture of A, D and E). If iodate is present
the salt will turn blue/grey and the colour will be retained for several minutes before turning brown.
Test for iodide solution
Moisten the second part with 2-3 drops of the given solution (mixture of A, B and C). If iodide is present
the salt will turn blue and remain blue for several minutes before fading.
Determination of iodate content (if test A is positive)
1. Weigh X g of the given salt sample in a watch glass or petri dish.
2. Transfer the weighed salt into a 250 mL conical flask and dissolve it in 50 mL water.
3. Add 1 mL 2N H2SO4 then add 5 mL of 3 % KI solution using a measuring cylinder. The solution will
turn yellow. Wrap the mouth of conical flask with a piece of filter paper and keep it in cupboard for
10 minutes.
4. Take 60-80 mL solution of (approx.) 0.005 M Na2S2O3 in a plastic beaker and use for titration. Rinse
burette and fill it in and adjust zero level.
5. Remove flask from cupboard and titrate with Na2S2O3 solution until the solution turns pale yellow.
Now add approx. 10 drops of starch indicator. The solution will turn dark purple. Continue titrating
until the solution becomes colourless.
6. Record the volume of titrant (Na2S2O3 solution) used and calculate the amount of iodine present in part
per million (ppm). Repeat three times to get average value.
Standardization of Sodium Thiosulfate
1. Pipette out 10 mL copper sulfate of concentration 0.005 M in a conical flask and add 5 mL of 3% KI
solution. The solution will turn yellow in colour.
2. Titrate with Na2S2O3 solution until the solution turns pale yellow. Add 7-8 drops starch indicator
solution at this stage.
3. Continue the titration until the purple colour fades, then add 5-6 drops of KSCN solution and titrate
again. The end point gives a colourless solution.
CAUTION: To ensure that you have obtained the true end point, stir the flask for 20 seconds and
then wait for 20 seconds to make sure that the purple colour does not reappear.
4. Repeat the titration to get concordant readings.
5. Calculate the molarity of the given sodium thiosulfate solution.
Conductometric Titration of HCl vs NaOH
Several compounds wholly or partially dissociate into ions when dissolved in water. The
conductance of such electrolytes depends on the concentration of the ions, temperature of the
medium and nature of the ion. The nature of ion becomes a variable with respect to the charge
and mobility of the particular ion. The electrolyte solutions obey Ohm’s law. Today a
conductometry cell will be used to perform a titration of HCl vs NaOH.
Procedure:
CAUTION: THE ELECTRODES ARE FRAGILE INSTRUMENTS. HANDLE THEM
GENTLY AND CAREFULLY.
DO NOT ADJUST THE KNOBS ON THE INSTRUMENT. THE INSTRUMENT IS
CALIBRATED.
1. Take NaOH (0.1 N) solution in a 50 mL burette and adjust zero reading.
2. Pipette out 25 mL of the given HCI solution in a 100/150 mL beaker. Add 25 mL of
water to this.
3. Now add the NaOH solution from the burette in (2 drops) 0.2 mL increments and record
the conductance after mixing the solution.
4. Continue the titration till you reach the initial conductance.
5. Repeat the experiment twice.
6. Plot the graph of volume of NaOH vs conductance and determine the equivalance point
of the titration.
7. Calculate the normality of HCl solution.
Determine the pI of a glycine using pH metry
Glycine (means ‘sweet tasting’ in Greek) is the simplest amino acid with a single hydrogen atom
as its side chain. In aqueous solution, glycine itself is amphoteric: at low pH the molecule can be
protonated and at high pH it loses a proton. Today you
will find the isoelectric point for glycine, which is the pH
at which a molecule or surface carries no net electrical
charge. At this pH, the molecule will not show any
motion in an electric field. For a molecule to have a
sharp isoelectric point, it should be amphoteric. Proteins
and amino acids are common molecules that meet this
requirement. For an amino acid with only one amine and
one carboxyl group, the pI can be calculated from the pKa values of this molecule. You will
perform a potentiometric titration for this study.
Procedure:
CAUTION: ELECTRODES ARE FRAGILE INSTRUMENTS. HANDLE
THEM GENTLY AND CAREFULLY.
DO NOT ADJUST THE KNOBS ON THE INSTRUMENT. THE
INSTRUMENT IS CALIBRATED.
1. Take 50 mL of NaOH (0.1 M) solution in a 50 mL burette and adjust zero reading.
2. Pipette out 25 mL of the given solution of amino acid in a 250 mL of beaker and add 25 mL of
distilled water to the amino acid solution using a pipette.
3. Insert the cleaned pH electrode into the beaker solution and record the initial pH of the solution.
4. Add NaOH in 0.5 mL increments from the burette. Stir the solution and mix it well.
5. Record the corresponding pH values until the pH starts increasing drastically. At this time, add
0.1 mL increments of NaOH till the pH stabilizes around 8.
6. After reaching pH 8, continue adding 0.5 mL increments of NaOH till you reach pH 11.
7. Plot the graph of pH vs volume of NaOH solution.
8. The two almost horizontal parts of the graph give the values of pKa1 and pKa2 for glycine. Use
mid-points of these regions to get the values.
9. The average of these values (pKa1 and pKa2) gives the pI of glycine.
To Determination the Energy of Activation
In today’s experiment you will study the reaction between potassium permanganate and dilute oxalic acid at
different temperatures. The permanganate ion MnO4- reduces to MnO2 changing the colour from bright
purple/pink to yellow/brown. You will find the rate constant for this reaction at five different temperatures and
then determine the activation energy for the reaction.
PRECAUTION: DO NOT TOUCH THE WATER IN THE WATER BATH
Procedure:
1. Using burettes, place 20 mL oxalic acid (0.25 M) in a conical flask and 10 mL KMnO4 (approximately
0.01 M) in a test tube. The exact concentration of KMnO4 should be noted from the blackboard.
2. Place both the vessels in a water bath to equilibrate for at least 5 minutes.
3. Mix the reactants in the conical flask and start the stop watch.
4. Swirl the reaction mixture regularly without taking it out of the water bath.
5. Record the time it takes for the mixture to turn yellow/brown.
6. Repeat the procedure with another sample at this temperature.
7. Repeat steps 1 to 6 for three other temperatures.
8. For the reading at 0 degree Celsius, time taken is 2215 seconds. Use this information as the fifth
temperature reading in your experiment.
9. Determine the activation energy by plotting ln k Vs 1/T. (Temperature in degree kelvin)
Data collection should consist of a table similar to:
[KMnO4] =
[Oxalic Acid] =
Observation and Calculations:
S.
Temp Temp 1/T Time
No. ( ͦ C)
(K)
for
trial
1 (s)
1
2
3
4
5
6
7
8
9
10
Time Average Rate=
k
ln
for
Time
[KMnO4]/time =Rate/[KMnO4][Oxalic] (k)
trial
(s)
2 (s)
Show sample calculation for determining the rate: Rate = [KMnO4]/time.
Show sample calculation for determining k: k =Rate/[KMnO4][Oxalic].
Plot of ln k Vs 1/T and determine Ea in kJ/mol from the slope.
Preparation of [Ni (NH3)6] 2+ and its Analysis by Complexometric Titration
(A) Preparation of Hexamminenickel (Il) chloride
1. Take 10 mL solution of nickel chloride hexahydrate (contains 3g of NiC12) in a 250 mL
beaker.
2. Take 10 mL solution of aqueous ammonia in a measuring cylinder.
3. Add the ammonia solution drop wise to the solution of nickel chloride with constant stirring
till the colour of the solution has changed from pale green to intense violet.
4. Allow the solution to stand at room temperature for 5 minutes, cover with watch glass. Then
cool it in an ice bath for about 15 minutes.
5. Filter the solution and wash the crystals with 3-5 mL ammonia solution.
6. Dry the crystals using filter paper.
7. Report the weight of the dried complex.
(B) Estimation of nickel(ll) by EDTA
1. Take 80 mL of 0.01M EDTA solution in a 250/500 ml plastic beaker and fill it in a clean burette
up to the mark.
2. Weigh accurately 0.23 g of [Ni(NH3)6]Cl2 complex and transfer this to a 100 mL volumetric flask.
Now add 50 mL of 1 N H2SO4 to dissolve it and makeup the solution to the mark with distilled
water.
3. Pipette out 10 mL of the complex solution in a 250 mL conical flask and dilute it with 15 mL of
distilled water.
4. Add 2-3 drops of murexide indicator and 5 mL NH4Cl solution (0.5M) to the conical flask. Now
add ammonia solution (7-10 drops) to maintain a pH 7 (light green colour of the solution).
5. Titrate it with EDTA solution till the endpoint is near, add 3ml of ammonia solution and continue
the titration till the endpoint (bluish violet colour appears).
6. Repeat the titration and get concordant values.
7. Calculate the amount of Ni present in the complex.
8. Put the remaining Ni complex, [Ni(NH3)6]Cl2, in a paper sachet, write your Roll Number on it and
submit the sample in the box labelled ‘Product’.
(C) Estimation of nickel by spectrophotometry
1.
2.
3.
Several solutions (of NiCl2) of known concentration and one solution of unknown concentration
will be provided to you.
Measure absorbance of all the solutions at 395 nm using a UV-Visible spectrophotometer.
Plot absorbance versus mg/mL of nickel. Determine the concentration of nickel present in the
unknown solution in g/L.
Extraction and Identification of DNA
DNA or Deoxyribonucleic acid contains all genetic information necessary for growth, functioning and
reproduction of almost all living organisms. DNA molecules consist of two biopolymer strands coiled around
each other to form a double helix. The chemical and molecular structure of the DNA is illustrated below.
Chromosomal DNA, exists in the well known X shape and is bound by proteins into a supercoil. DNA was first
isolated by Friedrich Miescher in 1869. Its molecular structure was first identified by James Watson and
Francis Crick at the Cavendish Laboratory within the University of Cambridge in 1953. In 1960, Nirenberg and
Har Gobind Khorana decoded DNA. Today you will extract DNA from peas and then identify it using UV-Visible
spectroscopy and a chemical test.
Procedure:
1. Take 5 mL of the (pea/onion) extract in a boiling tube and add 1.0 mL of the SDS solution and gently
swirl. Let the mixture stand for 10 minutes in ice.
2. Add 4 drops of papain extract to the mixture and stir gently.
3. Now hold the boiling tube at an angle and pour very slowly 15 mL of ice cold ethanol down the wall of
the test tube so that it forms a layer above the extract layer.
4. Allow the boiling tube to stand straight for a few minutes.
5. Some stringy white substance comes in the alcohol layer. This is
DNA.
6. Use a hooked glass rod and place it such that its end is just below
the alcohol layer. Now try to spool the DNA out of the tube.
Identification of DNA
Diphenylamine Test
1. In a test tube, add a small amount of crude DNA and 2 mL of 4% sodium chloride solution. Add 2 mL of
diphenyl amine reagent and mix.
2. Place the test tube in boiling water bath for one hour and record changes.
UV-Vis Absorption
1. Dissolve DNA in 2-3 mL of TE buffer solution and determine the ratio of absorption at 260 & 280 nm.
Amount of Calcium in Milk
Calcium is one of the minerals that the body needs daily. Milk and milk products have the best reserves
of calcium.
Interestingly, calcium
seems to come
in fifth place wherever it goes: it is the fifth most abundant element by mass in the Earth's crust (after
oxygen, silicon, aluminum and iron); the fifth most abundant dissolved ion in seawater (after sodium,
chloride, magnesium and sulfate); and the fifth most abundant element in the human body (after
oxygen, carbon, hydrogen and nitrogen). It is, however, the most abundant metallic element in the
human body, 99 percent of which can be found in our bones and teeth. Today you will find out the
amount of calcium in milk by doing a complexometric titration with ethylenediaminetetraacetate (EDTAsodium salt).
Procedure
STEP A: Preparation of Mg-EDTA Indicator Solution
1. Take 50 mL of Mg-EDTA solution in a 250 mL glass beaker. (This Mg-EDTA solution contains
0.74 g of EDTA and 0.49 g of MgSO4 in 100 mL of water.)
2. Add 2-3 drops of phenolphthalein indicator to the Mg-EDTA solution. Fill a burette with 0.1 M
NaOH solution.
3. Keep a white tile below the beaker containing Mg-EDTA solution.
4. Add the NaOH solution dropwise from the burette to the Mg-EDTA solution till it becomes light
pink. Note the volume of NaOH required for this process and discard this solution in chemical
waste bucket.
5. Take again 50 mL of Mg-EDTA solution in a 250 mL beaker and add the same volume of 0.1 M
NaOH solution as noted in step 4.
6. Add distilled water such that the total volume; Mg-EDTA + NaOH + Water = 95 mL
7. Add 2 mL of pH 10 buffer solution and 8 drops of Eriochrome black-T indicator to the above
solution. At this stage there are two possibilities.
7.1 If the solution is red in color, add 0.01M EDTA solution dropwise until the solution turns blue.
7.2 If the solution is blue in color add 0.01M MgSO4 solution dropwise until the solution turns red
then add 0.01M EDTA solution dropwise until the solution turns blue again.
8. This solution will be used as the Mg-EDTA indicator solution in part B of the experiment.
STEP B: Estimation of Calcium
1. Remove NaOH from the burette. Rinse the burette with 0.01 M EDTA and fill it with 0.01 M EDTA.
2. Pipette out 25 mL of the milk solution into a 250 mL conical flask.
3. Add 2 mL of pH 10 buffer, 10 mL of Mg-EDTA indicator solution (Prepared in step A) and 3 drops
of Eriochrome black-T indicator.
4. Titrate it with the standard 0.01M EDTA (exact concentration on blackboard) solution until the color
change is red to blue.
5. Repeat the titration to get concordant value.
6. Calculate the amount of calcium in the given sample in percentage or grams/liter.