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 80oC) 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 Rf value as the large salicylic acid reference spot. But then again another small spot at this Rf 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

Bendy carbon fibres

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?


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.


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 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.


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.”




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.



Silane is toxic and all reactions involving magnesium silicide to produce silane should be carried out in a fume cupboard.

In the August 2016 blog we looked at the reaction of magnesium with oxygen and chlorine.

Magnesium is so reactive it will take the oxygen away from silicon in sand, (silicon dioxide, SiO2) .

Magnesium reacts with sand to produce magnesium oxide and silicon.

2Mg + SiO2  -> 2MgO + Si

If an excess of magnesium is used, magnesium silicide (Mg2Si) is also formed.

2Mg + Si  ->  Mg2Si

Silane is produced when magnesium silicide is reacted with dilute acids, such as hydrochloric and sulfuric acids, to produce silane.

For example,  4HCl + Mg2Si  ->  SiH4 + 2MgCl2

Silane bursts into flames as soon as it is produced.

For example, SiH4 + 2O2  ->  SiO2  + 2H2O

Experiment 1 – heating magnesium and silicon dioxide using a Fresnel lens


The Fresnel lens here was taken from an old overhead projector


Quite a violent reaction, over in a few seconds


Boiling tube cracked by the ferocity of the reaction

Experiment 2 – heating magnesium and silicon dioxide using a Fresnel lens


Powdered magnesium and silicon dioxide


Heat using the focussed rays of the sun

Experiment 3 – heating magnesium and silicon dioxide in a crucible using a Fresnel lens


Adjust the lens to focus the suns rays onto the mixture


An exothermic reaction, over in a flash


The product mixture contains magnesium silicide which must be left to cool for several minutes

Experiment 4 – adding dilute sulfuric acid to the product in a boiling tube


Experiment 5 – adding dilute hydrochloric acid to the product mixture


The product mixture containing magnesium silicide


Small, sputtering explosions


Caught in the act

Experiment 6 – slow motion animation of adding dilute hydrochloric acid to the product mixture


Slow motion

Links to movies on You Tube:

Experiment 1
Magnesium and silicon (IV) oxide reaction 1 – boiling tube / Fresnel lens
16 seconds

Experiment 2
Magnesium and silicon (IV) oxide reaction 2 – boiling tube / Fresnel lens
39 seconds

Experiment 3
Magnesium and silicon (IV) oxide reaction 3 -crucible / Fresnel lens
24 seconds

Experiment 4
Magnesium silicide and sulfuric acid – boiling tube
14 seconds

Experiment 5
Magnesium silicide and sulfuric acid 3 – watch glass in the dark
40 seconds

Experiment 6
Magnesium silicide and hydrochloric acid at 120fps, no sound (watch glass)
56 seconds

Don’t look into the flame!

At the end of the last post which was about the reactions of the alkali metals with chlorine, I asked what you would see if burning magnesium were lowered into a gas jar containing chlorine gas. The answer is you would see a bright white flame.

The flame produced by magnesium in such reactions is so bright you must not look at it directly, because of the risk of damaging your eyes. Always wear eye protection and follow the recommended safety advice in your laboratory.

Sometimes its easier, although less exciting to look at a movie of such reactions. Movie here, (51 seconds).


Magnesium burns in oxygen with a bright, white flame

What is the product of the reaction between magnesium and oxygen? Can you write a balanced symbol equation for the reaction?

Magnesium also reacts violently with chlorine, producing the same bright, white flame. In the movie here (54 seconds), burning magnesium ribbon is lowered into a gas jar containing a mixture of chlorine gas and air. Listen for the crackles at about 18 seconds into the movie, indicative of the very energetic nature of the reaction.


Magnesium reacting with chlorine and oxygen from the air

What is the product of the reaction between magnesium and chlorine? Can you write a balanced symbol equation for the reaction?

Another movie of the reaction between magnesium and chlorine from the RSC Science Skool can be seen here,(1 minute 28 seconds).

Magnesium is so reactive, it will also take the oxygen away from silicon in silicon (IV) oxide. Carry out some research and find out the products of this reaction and an equation for it.

Answers to all of the questions in this post next time.


Making lithium chloride, sodium chloride and potassium chloride by direct combination

The alkali metals lithium, sodium and potassium all react with chlorine gas to produce their respective chlorides.

2Li + Cl2 –> 2LiCl

2Na + Cl2 –> 2NaCl

2K + Cl2 –> 2KCl


Experimental set-up

In the animation below samples of lithium, sodium and potassium metals are shown being heated in a blue Bunsen burner flame until they catch fire. The flaming metals are then dropped into gas jars containing chlorine gas. Each metal reacts showing its characteristic flame colour and a product seen as a white smoke (solid particles of the chloride).


Flaming metals and chlorine

Here are three pictures taken from a movie of the reactions which can be viewed on You Tube here (3 minutes).


Lithium reacting with chlorine to make lithium chloride


Sodium reacting with chlorine to make sodium chloride


Potassium reacting with chlorine to make potassium chloride

We carried out the experiments in broken porcelain crucibles because when we used combustion spoons the iron metal, in the steel from which the combustion spoons were made, also reacted with the chlorine gas (movie here). Dense red smoke, presumably of iron (III) chloride, was produced in addition to the alkali metal chlorides.

We had to use broken crucibles because whole crucibles would not fit inside our gas jars.


Combustion spoon: lithium and chlorine


Combustion spoon: sodium and chlorine


Combustion spoon: potassium and chlorine


Products from experiments carried out in porcelain crucibles were all white crystalline solids

As for comparing the reactivity of the three metals, potassium reacted the fastest, with the reaction over in a couple of seconds. Both lithium and sodium reacted for several seconds, with longer lasting flames.

However, the reaction between potassium and chlorine shown in the animation above did not go to completion because there was unreacted chlorine gas at the end of the experiment, (visible in the picture). There was also unreacted potassium and this reacted violently with water when the gas jar was was washed out, after allowing it to cool.

An animation of the reaction between potassium and chlorine in another experiment is shown below (movie here). This one did not react so violently, but was also over in a couple of seconds.


Close-up: potassium reacting with chlorine

Thus, burning potassium reacts quicker than lithium and sodium with chlorine gas, but its almost as if the potassium chloride produced smothers the reacting metal.

What colour flame would you expect to see if burning magnesium ribbon was dropped into a gas jar full of chlorine gas?