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.

Two dark solids

Two dark solids, one giant ionic, the other covalent molecular.

Iodine

Iodine is a molecular solid. It consists of I₂ molecules. The molecules are non-polar and are held together by weak intermolecular attractions in the iodine crystals.

Potassium manganate (VII)

Potassium manganate (VII) is an ionic solid, made up of K⁺ ions and MnO₄^– ions in a giant crystal lattice structure.

The differences in bonding and structure between the two materials means they have quite different properties.

Solubility

Water dissolves many polar molecules and ionic solids.

Here water is being added to iodine and potassium manganate (VII).

Which tube contains the iodine and which the potassium manganate (VII) ?

Hexane is a non-polar organic liquid which dissolves non-polar molecules, but not ionic solids.

Here hexane is being added to iodine and potassium manganate (VII).

Which tube contains iodine and which potassium manganate (VII) ?

Behaviour in an electric field

What would happen if solid iodine and solid potassium manganate (VII) were placed on some filter paper moistened with tap water and subjected to 10 volts?

In potassium manganate (VII) the K⁺ ions are colourless, whilst the MnO₄^– ions are deep purple in colour.

Which way do the purple ions move and why?

Iodine dissolves in potassium iodide solution

The I₂ molecules in iodine do not dissolve very well in pure water, but they do dissolve in a solution of potassium iodide.  I₃^– ions are formed when iodine dissolves in potassium iodide.

Behaviour in an electric field

What would happen if solid iodine and solid potassium manganate (VII) were placed on some filter paper moistened with potassium iodide solution and subjected to 10 volts?

Why does the brown colour move to the left at the top of the slide?

This one maybe harder to explain fully. So let’s close with a little rhyme:

“Two dark solids, sometimes purple, sometimes brown. One ionic, one molecular, with behaviour that can make you frown.”

 

Spooky

Sometimes things don’t go the way the way they’re planned. When this happens in chemistry the first thing is to see if you can repeat the observation. Some examples follow, maybe you can explain them.

Here’s an experiment we did recently using the juice extracted from red cabbage as an acid/base indicator. We were surprised by the colour produced in strong alkali on the far left of the picture.

Yellow at pH 12 to 14?

We checked with a second extract and got the same result

Red cabbage juice in 1M NaOH and 2M HCl

In another experiment we tried to grow one big crystal of sodium chloride by suspending a ‘seed’ crystal in a saturated solution of brine. The seed crystal fell off and then something magical happened around the string.

Boxes, little boxes..

Finally, when we tried to demonstrate diffusion by dissolving a small lump of potassium manganate (VII) in a beaker of water, things didn’t go as smoothly as planned:

Jittery time lapse animation (overnight)

Science, the fun part is in figuring out what happened.