This plastic tastes like cheese

Casein is the main protein in cows milk. It precipitates rapidly when vinegar is added to warm milk. The precipitated casein and other milk solids can be filtered, shaped and dried into a hard plastic like material frequently referred to as milk plastic.

Equipment & ingredients

Both full cream milk and low fat milk contain casein and give a good yield of product. The casein obtained from full cream milk gives a more pliable, less crumbly material, presumably due to the higher fat content. However, the solid also exudes more oily liquid during the drying stage.

We used the following recipe:
1. Heat 300ml milk to 65 C.
2. Add 20ml of white vinegar and stir.
3. Allow the precipitated casein to cool a little before filtering through a sieve.
4. Press the filtered solid between several layers of paper towel.
5. Mould the solid into the desired shape and dry in a low temperature oven or on a warm shelf near to a radiator.

Making milk plastic

Heat the milk to 65 C

Add vinegar and stir to precipitate out the casein and other milk solids.

Precipitating the casein

A closer look:

Lower the pH to precipitate the casein micelles

Cottage cheese anyone?

How do we get the cheese out?

Filter it through a sieve

Separating the milk solids from the whey

Press out the excess fluid before drying

Press between layers of paper towel

What shape will you make?

Peace dove or bunny ears?

You can add food dye to colour the milk during the heating stage

This season it’s orange

The drying stage takes several days or even weeks, until a very hard plastic-like material is produced.

Drying time, days to weeks

Hey presto! Milk plastic.

Milk plastic or dried cheese?

It’s too hard to eat, but it may still go mouldy under the right conditions.

Movies for the precipitation, filtration and shaping are on You Tube.

More information on making milk plastic and cottage cheese.

Finally, on the chemical reaction in the baking powder submarine last time

The baking powder used in the toy submarine last time contained two active ingredients, disodium diphosphate and sodium bicarbonate. These two chemicals react to produce carbon dioxide which was released as bubbles that caused the submarine to rise.

Na₂H₂P₂O₇ + 2NaHCO₃ → Na₄P₂O₇ + 2H₂O + 2CO₂

The slow release of carbon dioxide is what allows the submarine to rise and dive

This reaction tends to release carbon dioxide at a slower rate than the sodium bicarbonate reacting with an acid like vinegar.

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Voyage to the bottom of the water trough

Baking powder submarines are great fun.

All you need is a submarine, baking powder and water

The baking powder provides the chemical reactants needed to produce the gas which raises the sub.

What are the key chemical reactants listed here?

At the surface the sub takes on water and sinks to the bottom of the water trough. A chemical reaction takes place producing gas inside the sub which causes it to rise back to the surface. The gas is emitted from the sub and the cycle repeats, diving and rising.

Diving and rising

A video clip from which the above animation was constructed can be viewed here.

Questions

1. What is the chemical reaction that takes place when the baking powder dissolves in water?

2. What is the gas produced that raises the sub?

Answers next time.

 

TLC under UV

Thin layer chromatography of three analgesics and caffeine under U.V. light was carried out in order to show the separation taking place. It is not a recommended technique in the laboratory.

U.V. Hazard
Due to the nature of the uv hazard polycarbonate safety spectacles (which absorb short wavelength U.V. light) and rubber gloves were worn throughout.

Five samples were run on a single TLC plate.

The samples were (left to right on the plate):
IBU = Ibuprofen
CAF = caffeine
u? = a commercial ‘pain relief’ medicine, used as an unknown
PAR = paracetamol
ASP = aspirin

Five samples prior to elution

The samples were dissolved in ethanol for spotting onto the plate. The TLC plate was run in an open beaker under short wavelength u.v. light using ethyl ethanoate as the eluting solvent.

Separation of the samples

The movement of the dark purple spots (samples) during the running of the plate can be observed in the animation. The original movie can be viewed here.

Rf values can be measured

It is easy to see which are the two active ingredients in the commercial pain relief medicine by comparison of the spots with the standard reference materials running on either side.

This kind of thin layer chromatography experiment is often carried out at school.

Making two sulfates

At GCSE level Chemistry making salts is divided into three categories:

1. by reacting an acid with an excess of an insoluble base
2. by titration
3. by making a precipitate

Making the salt copper (II) sulfate from copper (II) oxide falls into the first category, since copper (II) oxide is a basic oxide which is insoluble in water, but it reacts with sulfuric acid in a typical acid/base reaction.

acid + base -> salt + water

Copper (II) oxide is insoluble in water (shown on the left), but reacts with sulfuric acid producing a solution of copper (II) sulfate (blue on the right)

Making ammonium sulfate from aqueous ammonia requires a titration technique, since both the acid (sulfuric acid in this case) and the base (ammonia) are soluble in water. We call bases that are soluble in water alkalis, so we can re-write the general equation here as:

acid + alkali -> salt + water

Making copper (II) sulfate

Add copper (II) oxide to sulfuric acid whilst heating gently and stirring with a glass rod

Do not boil!

The copper (II) oxide reacts with the acid producing copper (II) sulfate as a blue solution

Stop adding copper (II) oxide when no more reacts and the black solid collects at the bottom of the beaker. At this stage all of the acid has reacted.

All of the acid has been used up and the excess of insoluble base collects at the bottom of the beaker

The excess copper (II) oxide is removed by filtration.

Filter off the insoluble black solid

Transfer the filtrate to an evaporating basin.

The product copper (II) sulfate is dissolved in water as a blue solution

Water is removed by evaporation over a Bunsen burner. Stop heating when crystals start to form at the edge of the solution.

Solid forming at the edge of the solution

Place the evaporating basin on a windowsill or in a drying cabinet to allow crystals to form.

Making ammonium sulfate

In an approximation to carrying out a titration, small volumes of ammonia can be added to sulfuric acid in a beaker using a dropping pipette. After each addition of ammonia a sample of the reacting mixture is spotted onto a small piece of Universal Indicator paper using a glass rod. Stop adding ammonia when the Universal Indicator turns green or green/blue.

In the picture below the beaker in the centre shows the reaction mixture, the small squares of Universal Indicator paper on the tiles have been spotted after adding aliquots of ammonia (from right to left).

The small squares of Universal Indicator paper show the progress of the reaction from right to left, (the tile on the right used up before the tile on the left).

Stop adding ammonia when the indicator turns green. The last four squares of paper on the bottom left (of the tile of the left) were not used.

Reduce the volume of the reaction mixture by heating over a Bunsen burner.

Allow crystals to form in a warm, dry place

The third method of making a salt, by precipitation, is not illustrated above.

Which sulfate is formed as a white precipitate and used as a test for the presence of the sulfate ion in GCSE Chemistry?

Which method to use?

Making salts is a topic studied at GCSE Chemistry. Copper (II) sulfate and ammonium sulfate are two salts commonly made at school. Typically, these two salts are made by reacting a base with sulfuric acid.

Here is a GCSE Chemistry problem to solve:

If the bases available are copper (II) oxide (a black powdered solid) and aqueous ammonia (a 1M solution of ammonia in water), describe the method you would use to make each of the salts?

Examples of the salts crystallising in evaporating basins made from experiments performed at school are shown in the gif animations shown below.

Which salt is being crystallised here?

Colourless needles – which product is shown above?

Answers next time.

 

Using TLC to monitor a reaction

Thin layer chromatography (TLC), is often used to monitor the progress of chemical reactions. Paper chromatography and silica TLC may be used.

First a solvent system which leads to a good separation of the starting materials and products for the reaction must be found, along with a technique for visualising the compounds on the TLC ‘plate’.

In the examples below, aspirin and salicylic acid are shown on silica TLC plates run in a solvent system consisting of a mixture of petroleum ether (boiling point range 60 to 80^oC) and acetone.

Two mixtures of these solvents were compared; pet.ether : acetone 85:15 and pet.ether : acetone 7:3. The ratios indicate the relative volumes of each solvent in the mixture.

Both TLC plates were labelled at the top in pencil to indicate the samples spotted. S = salicylic acid and A = aspirin. The aspirin and salicylic acid were dissolved in a mixture of ethanol and dichloromethane 50:50 for ‘spotting’ on the plates, although the aspirin did not dissolve completely.

Pet. ether : acetone, 85:15

Pet. ether : acetone, 7:3

The spots were visualised under short wavelength u.v. light (around 254 nm). The TLC plates can be seen to ‘glow’ with a characteristic yellow/green colour due to a u.v. fluorescent indicator that is incorporated into the silica. Organic molecules such as salicylic acid and aspirin block the u.v. light from reaching the fluorescent indicator and appear as dark purple spots.

One must wear gloves and goggles which absorb the harmful u.v. light when carrying out such experiments.

A hydrolysis experiment using aspirin

Approximately 0.1g of aspirin was dissolved in 1cm³ ethanol in a small glass vial. Then 1cm³ of 2M sodium hydroxide was added to the vial and a sample spotted onto a TLC plate and labelled t=0 (for time = zero at the start of the experiment). The mixture was left to react overnight at room temperature and then acidified (to pH 1) using 6M HCl. Another sample of the reaction mixture was then spotted onto the TLC plate and labelled t=18h (for reaction time = 18 hours). Reference samples of aspirin and salicylic acid were also spotted. All samples were spotted using small glass capillary tubes drawn out to a fine point in a Bunsen burner flame.

Pet. ether / acetone, 7:3 Asp = aspirin, t=0 is from the start of the reaction and t=18h is after 18 hours. SA = salicylic acid.

A few points to note:

Samples t=18h and SA are overloaded in size, meaning too much of these two samples were spotted on the plate.

The sample at t=0 was underloaded and nothing much could be seen apart from the material remaining on the baseline.

The sample at t=18h does show a small spot with approximately the same R_f value as the large salicylic acid reference spot. But then again another small spot at this R_f value is also shown by the aspirin sample spot, although this is much weaker.

Does the TLC plate show evidence of a base hydrolysis reaction having occurred?

We think it does, but the evidence is not conclusive and more tests would need to be carried out. What the TLC plate does show is that no starting material was detected in the reaction mixture after 18 hours indicating that some reaction had occurred.

Flexible carbon electrodes

Carbon fibres can be used as flexible electrodes in high school electrolysis experiments. Here is an example showing the electrolysis of aqueous silver nitrate solution.

Cathode in the middle, anode around the outside

Silver grows out from the centre in all directions.

Go micro and have fun with flexible carbon electrodes.

 

Pencil lead reduction

Iron (III) oxide can be reduced to iron by using a pencil lead.

Experimental set-up

In the experiment a pencil lead was supported on a white tile and connected at a 12 volt power pack.

A mixture of iron (III) oxide and powdered carbon was placed on top of the pencil lead. A bar magnet was passed over the top of the powdered mixture and none of the solid was attracted to it.

Chemicals and equipment

Placing the reaction mixture on a ceramic tile

Nothing picked up by a bar magnet

The power pack was turned on and the pencil lead glowed brightly. After about a minute and a half of heating the pencil lead broke.

Thermal reduction

On cooling, when a magnet was passed over the powdered reaction mixture several specks of solid containing iron were attracted to it.

Attractive product – we’ve made some iron!

Specks of solid stick to the magnet

A movie of the experiment can be viewed here.

Answers to the questions at the end of last months blog:

1. Chlorine gas was produced at the anode. The chlorine gas diffused across the tile reacting with the blue litmus paper, the potassium iodide solution and the potassium bromide solution.
2. Copper metal was produced at the cathode.
3. Chlorine gas dissolved in water on the surface of the paper. The products of the reaction of chlorine with water were hydrochloric acid which turned the litmus pink and hydrogen hypochlorite which bleached it white.
4. The potassium iodide turned orange/brown due to the formation of iodine. Chlorine displaced the iodide ions from solution producing chloride ions and iodine.
5. The potassium bromide turned yellow/orange due to the formation of bromine. Chlorine displaced bromide ions from solution producing chloride ions and bromine.
6. The blue colour of the copper (II) chloride solution faded as copper metal was deposited at the cathode. The blue colour was due to the copper (II) ions which were removed from solution as they reacted at the cathode.

 

Microscale electrolysis of copper (II) chloride

Using pencil leads as electrodes it is possible to carry out electrolysis experiments at school on a very small scale or micro-scale. The advantages of microscale are reduced equipment needs and environmental burden, speed, simplicity and safety.

The picture below shows the equipment used in a microscale electrolysis experiment on copper (II) chloride solution. This is modelled on the same electrolysis experiment described by Bob Worley on Microchemuk.

Microscale electrolysis on a plastic sheet placed on a white tile

A movie showing how the experiment was set up and carried out can be seen here on You Tube.  A close-up of the experiment just before the power pack was turned on (6v d.c.) is shown below, complete with labels.

Funny face labels

Here’s a gif animation of the electrolysis experiment:

Face the changes

The results of the experiment:

Face change, what happened?

Questions

1. What was the gas produced at the anode?

2. What was the solid produced at the cathode?

3. Why did the blue litmus paper turn pink and then white?

4. Why did the potassium iodide turn orange / brown?

5. Why did the potassium bromide turn yellow / orange?

6. Why did the blue colour of the copper (II) chloride fade?

Answers next time.

Go microscale!

A microscale mag sulfate experiment

The materials for making magnesium sulfate in a microscale reaction

Making salts on a micro-scale in a high school chemistry laboratory has the advantages of being quick, simple and easy to prepare. For example, here is an experiment we recently carried out making magnesium sulfate in an IGCSE Chemistry class.

A two decimal place electronic balance, tweezers and a 0.2ml (200 microlitre) automatic pipette were used in addition to the materials shown above.

Method

1. Cut out a 10cm x 10cm plastic sheet (from an A4 plastic wallet) and record its mass on an electronic balance.

2. Place 0.2ml of 1M sulfuric acid on the centre of the sheet.

3. Take a 4 cm length of magnesium ribbon and bend it into an oval shape which can be placed around the globule of acid.

4. Place the magnesium around the acid.

Ring fence hubble bubble

5. When the bubbling has stopped, pick up the magnesium with a pair of tweezers and rinse off with a small volume of distilled water.

Rinse, rinse

6. Place the plastic sheet in a suitable place, such as a sunny windowsill or a low temperature drying cabinet, to allow the water to evaporate and the salt to crystallise.

Magnesium sulfate – has all the water gone or not?

7. Weigh the plastic sheet + the solid using an electronic balance and determine the mass of product formed.

8. If time allows, break up the crystals with a spatula and dry further using warm air from a hair dryer until a constant mass is obtained.

Calculations can be performed to determine the percentage yield of magnesium sulfate (presumably the heptahydrate) made.

The plastic sheets can be washed and re-used.

Thanks to Bob Worley for inspiring us to try out this and other high school chemistry lab experiments on a microscale. Here is a link to Bob’s Microchemuk website which is a goldmine of resources and ideas.