Balls of fire

Cotton wool balls are very easy to nitrate on a micro-scale and can provide a stimulating demonstration to support teaching chemistry at Advanced Level at school. Adding a mixture of concentrated nitric acid and concentrated sulfuric acid to cotton wool produces nitrocellulose and when dried this burns spectacularly. A bright yellow flame characteristic of the presence of sodium is evident.

Do not do this!

Teacher demonstration A method is described below for the nitration of cotton wool on a microscale which considerably reduces the risks involved and minimises the impact on the environment in terms of the quantities of chemicals used and amount of waste generated.

Microscale nitration of cotton wool

After allowing the nitrating mixture of concentrated nitric acid and concentrated sulfuric acid to react with the cotton wool for 20 to 30 minutes whilst sitting on ice, the reaction is terminated by washing the cotton wool with water and then sodium bicarbonate solution. The latter step neutralises any acids present. We decided to see what effect it would have on the product and its burning characteristics by varying the base used in the neutralisation step.

Accordingly, potassium carbonate, lithium carbonate, limewater (calcium hydroxide) and a barium carbonate slurry were each used in place of the sodium bicarbonate in a series of experiments. The results of burning the nitrocelluloses produced are shown in the animations below, with the nitrocellulose produced using a sodium bicarbonate wash shown first for comparison.

Nitrocellulose – sodium

Nitrocellulose – sodium (from 1000fps movie)

Nitrocellulose – lithium

Nitrocellulose – lithium (from 1000fps movie)

Nitrocellulose – calcium

Nitrocellulose – calcium (from 1000fps movie)

Nitrocellulose – barium

Nitrocellulose – barium (from 1000fps movie)

Nitrocellulose – potassium

Nitrocellulose – potassium (from 1000fps movie)

Two questions for chemistry students (answers next time)

1. What are the characteristic flame test colours produced by a) sodium, b) potassium, c) lithium and d) calcium?

2. Why does barium produce a bright white flame in the above images whilst its characteristic flame test colour is green?

Here are some links to the original movie files uploaded to You Tube from which the animations were constructed.

Burning nitrocellulose – sodium, filmed at 120fps

Burning nitrocellulose – sodium, filmed at 1000fps

Burning nitrocellulose – potassium, filmed at 120fps

Burning nitrocellulose – potassium filmed at 1000fps

Burning nitrocellulose – lithium, filmed at 120fps

Burning nitrocellulose – lithium, filmed at 1000fps

Burning nitrocellulose – calcium, filmed at 120fps

Burning nitrocellulose – calcium, filmed at 1000fps

Burning nitrocellulose – barium, filmed at 120fps

Burning nitrocellulose – barium, filmed at 1000fps

Finally, thanks to Tony Pluck for his help and encouragement in developing this work.

Food Colour Frenzy

We have had great fun with food colours and microscope slides recently reproducing an experiment with Dancing Droplets described by Nate Cira, Adrien Benusiglio and Manu Prakash. They posted a movie on Vimeo showing how to do it “Dancing Droplets: How to easily recreate the phenomena at home“.

We took up their invitation to have a go and found that red food colour was the best at pushing droplets of other food colours around.

The Power of Red!

Here’s our movie posted on You Tube ‘Red Peril’ Food Dye, from which the above gif animation was constructed.

More details about the phenomena discovered by Nate Cira can be found on Scientific American “2 Common Liquids Spontaneously Form Dancing Droplets“.

More efforts from us (sorry about the wobbly camera work):

Our first success

Twins

The Big Red

Can you see a heart?

You might decide to have a go for yourselves. We found using an automatic micropipette (200 microlitres) speeds up the procedure considerably and increases the likelihood of success.

More fun with luminol

We recently carried out a reaction with luminol dissolved in sucrose solution.

Luminol reaction in a measuring cylinder

The idea was that by dissolving the luminol reaction mixture in a concentrated sucrose solution we could carry it out in the bottom of a measuring cylinder and ‘float’ a solution containing copper (II) ions and iron (II) ions above it. The experiment worked quite well and a short movie clip can be viewed here.

After a few seconds

The details of how we carried out the experiment are shown below

How we did it

Take care, by the time its all over it can get fairly messy as the images show below.

Lights on, lights off

Almost over

Another fine mess…

Here’s an image of our first experiment done on a smaller scale (approximately one tenth) using 0.025g luminol and in a 20ml measuring cylinder.

On a smaller scale

Thanks to Chris Kruger and Jonathan Barton in developing this experiment.

We hope you have fun with luminol too!

Answers to the “Bicarb rockets” questions last time:
1. sodium hydrogen carbonate + ethanoic acid -> sodium ethanoate + water + carbon dioxide
NaHCO₃ + CH₃COOH -> CH₃COONa + H₂O + CO₂

2. f = ma
The carbon dioxide produced by the chemical reaction builds up inside the bottle and when the bung finally comes out the gas pressure forces the liquid out of the bottle quickly. It is the force of the liquid moving downwards which pushes the bottle upwards.

Bicarb rockets

Bicarb rockets are fun to make and launch.

Lift off!

We use 150 ml vinegar and 3.5g sodium bicarbonate in a 600 ml plastic bottle. It’s easy to make a bicarb rocket if you have a rubber bung which tightly fits into the top of the bottle. We use a tripod to launch from as shown below and can easily reach a height of a three storey building using these materials.

The rocket building materials

The rocket, a 600ml water bottle

Vinegar (dilute ethanoic acid)

Sodium bicarbonate
(sodium hydrogen carbonate)

The bicarb is wrapped in paper towel to produce a small pellet

Do this outside!

Countdown

5, 4, 3, 2,…

Here are some .gif animations of various bicarb rocket flights. Most of them were recorded at 240 fps (frames per second) using a Casio EX-FH100 camera.

A whole flight in frame

A little closer

Closer still

Too close – do you want vinegar on that!

Hit the roof

How high can you go?

Viewed from above (shot at 120 fps)

Questions (answers next time)

1. What are the word and balanced symbol equations for this chemical reaction?

2. What is it that actually propels the rocket upwards?

“Propane Rockets”

SAFETY FIRST: This demonstration must be done with care and adequate ear protection must be worn by all present, along with safety spectacles. The use of commercial ear protectors are advised.

A propane rocket

One demonstration that can be used to show how the ratio of reacting gas volumes relates directly to mole quantities in a balanced chemical equation is with a “propane rocket”.

In the experiments described below we used bottled gas, or LPG, which contains a mixture of propane and butane, in a 600 ml plastic bottle. Oxygen was also required. When ready to launch the bottle rocket was suspended from a fishing line using two paper clips and two rubber bands.

Fuelled and ready to launch

The balanced equations for the complete combustion of propane and butane are shown below.

Equations for the complete combustion of propane and butane

Thus, it can be seen from the first equation that 5 volumes of oxygen are required to completely burn 1 volume of propane.

The gases are introduced into the bottle rocket by displacing water as shown. It is recommended that a fresh plastic bottle is used each time.

Oxygen from a gas syringe

Followed by LPG

When the gases are used in the proportion of 5 volumes of oxygen to 1 volume of bottled gas, a very loud explosion propels the bottle at speed across the laboratory.

Complete combustion.. BANG!

Reduce the proportion of oxygen and incomplete combustion occurs. By trial and error a slower flying rocket can be obtained.

With 3 to 4 volumes of oxygen : 1 volume of propane a slower rocket results

However, if the proportion of oxygen is reduced too far the rocket will fail to launch.

Too much fuel, not enough oxygen – disaster!

All of the above .gif animations were produced from high speed movies shot at 120 fps (frames per second). Here is another animation produced from a movie shot at 1000fps

High speed movie 1000fps

Frame by frame

Thus, the effect of varying the molar ratio of the two reacting gases, using reacting gas volumes, can be shown.

Match head rockets

Match head rockets are fun to make.

For safety we use no more than four match heads.

We have lift off

There are many videos on You Tube describing how to make these rockets.

Here’s how we do it:

About 15 x 10 cm

But trial and error is involved

Roll smoothly

and tightly

This is the rocket body

Making way for the fuel

No more than four

Don’t worry if the solid crumbles off

Experiment with a mixture of crumbles and whole match heads

Don’t let the stick fall out whilst loading

Use a second stick as a ram rod

This tail is probably a little too short

The pin forms part of the launch pad

Wrap tightly to form an exhaust tube

You can twist the foil too

Make sure the pin moves in and out smoothly

This is going to roll up

Like a party toy or tube of toothpaste

Squash the end as tightly as you can

Pliers do work

But so does biting down with your teeth

All systems go

You can vary the angle easily by pushing the end of the pin into Plasticene

The launch – Wear eye protection and stand well back

This one seemed to have too many whole match heads which shot out of the rocket quite spectacularly

Here are three more animations of another successful launch.

Students were not allowed to hold the Bunsen burner as shown.

Teachers do so at their own risk, but they must wear gloves and a full face mask if they do.

Crumbled up match head solid fuel seems to burn more smoothly

They can sometimes fly several metres

MUST WEAR GLOVES (unlike here)

The squashed and rolled up aluminium foil unwinds on launch and is puffed out

Launch pad failure

Launch pad failures are quite common and seem to occur particularly when the nose is not flattened tightly enough during manufacture or when a hole is made in the foil.

Answers to two chemical reactions last time

Equations for the reactions are:

iron + sulfur   –>  iron sulfide

Fe  +  S   –>   FeS

and

copper (II) oxide  +  carbon   –>   copper  +  carbon dioxide

2 CuO   +   C    –>   2 Cu   +   CO₂

Both reactions give out a lot of heat and this can be seen by the intense glow which continues even when the external heat source is taken away.  Reactions which give out heat like this are called exothermic reactions.

Here are some pictures of the products.  First the iron sulfide:

We had to smash the test-tube to get the product out of this one

Unfortunately, it was still attracted to a magnet, typically due to unreacted iron filings

The fused lump of product probably consists of iron sulfide plus unreacted iron filings and some sulfur

Here is the product of the thermal reduction of copper (II) oxide with carbon:

Brown coloured copper metal is produced

 

 

Two chemical reactions

Two chemical reactions often done at school are the reaction of iron filings with sulphur and the reduction of copper (II) oxide with carbon. Answers to the questions posed below will be given in the next blog.

1. The reaction of iron filings with sulphur

Iron filings and sulphur

The reaction is hot!

How can you tell that the reaction produces a lot of heat?

What is the name of the product?  Will it be magnetic like the iron filings?

The product of the reaction

2. The reduction of copper (II) oxide with carbon

A mixture of copper (II) oxide and carbon

Mini volcano!

How can you tell that the reaction produces a lot of heat?

What are the products of this reaction? What will the contents of the test-tube look like when it has cooled down?

What is the new brown coloured material?

Two exothermic reactions often carried out in school.

Can you write equations for these two reactions?

Fun with luminol

Luminol is a chemical frequently used to produce blue light by chemiluminescence.

One classic demonstration uses luminol in a reaction carried out as it spirals down a length of tubing such as the one shown here:

Chemiluminescence, courtesy of Nigel Evans, Bloxham School.

Another demonstration uses luminol in a spray to detect fake ‘blood spatter’ such as is often seen in TV forensic science dramas. In our school chemistry club we used a recipe published by Anne Marie Helmenstine on About.com This recipe requires two solutions which are mixed together in equal proportions shortly before use. Solution 1 – Luminol stock solution (2 g luminol + 15 g potassium hydroxide + 250 mL water). Solution 2 – 3% hydrogen peroxide in water (our local pharmacy supplied a 6% hydrogen peroxide solution which needed to be diluted by half with water). However, because this mixture uses a fairly concentrated solution of potassium hydroxide, which is corrosive, it must be used in a fume cupboard. Under such conditions we were unable to see much of a glow when sprayed onto test paper ‘painted’ with an iron (III) chloride solution.

We had much more fun by pouring the luminol mixture into a Petri dish in a darkened room. We then added 0.1M iron (III) chloride and 0.1M copper (II) sulfate using dropping pipettes. In this way more students were able to investigate this exciting reaction for themselves:

Luminol reaction in a Petri dish

In another similar experiment we tried electrolysing the reaction mixture using graphite electrodes and 9v from a power pack:

Electrolysing the luminol reaction

You too might have some fun if you try this. Happy Halloween from our Chemistry Club.

Strawberry blonde

How do you make a strawberry blonde?
Why, hydrogen peroxide of course!

We have been doing some experiments with strawberries recently. First we wanted to look at the effect of freezing on strawberries.

Strawberry frostbite

So we did an experiment where we observed what happened when a frozen strawberry (on the left) and a chilled strawberry (on the right) were allowed to warm up to room temperature side by side.

On your marks!

Here’s an animation of the warming up process:

Slushy

The ice crystals in the frozen strawberry have destroyed much of the internal structure of the strawberry such that it collapses into a mush, oozing out juice on thawing.

Mushy

Our second experiment looked at what happened when a strawberry was placed in hydrogen peroxide bought from a local pharmacy. Take care to read the safety precautions on the label when using hydrogen peroxide.

How do you make a strawberry blonde?

Hydrogen peroxide can be used to take the colour out of hair or a strawberry.

In we go… just sit back and relax.

Hey presto, a white strawberry.

And the colour is gone

The bleached strawberry still smelled like a strawberry. Do not eat!

Bleached strawberry

Our third experiment was a combination of the first two experiments. We decided to look at the effect of temperature on the bleaching process. Accordingly, we dipped a frozen strawberry and a room temperature strawberry into hydrogen peroxide side by side.

Which one will bleach quicker?

We measured the temperature difference between the two beakers.

Warm and cold

And we watched what happened.

Which one loses colour faster?

After about thirty minutes the cold one had warmed up a little.

Time to come out

And the warmer strawberry had lost more colour than the frozen one.

Warmer bleaches faster

Increasing the temperature, increased the rate of reaction. In this case the bleaching of a strawberry. Although the loss in colour was not even across the surface of the strawberry in this experiment.

Close up

Last time I asked which metals were unlikely to react with 2M hydrochloric acid? The best answer is to pick metals such as gold and platinum at the bottom of the reactivity series. This would be an expensive experiment to carry out though.

Unreactive metals

One way to make a salt is to react a metal with an acid.

acid + metal -> salt + hydrogen

However, not all metals react with acids and those that do often react at different rates.

GCSE level students learn a reactivity series of metals. This lists a number of common metals in order of decreasing reactivity in chemical reactions.

One such list (in order of decreasing reactivity) would be:

Potassium
Sodium
Lithium
Calcium
Magnesium
Aluminium
Zinc
Iron
Tin
Lead
Copper
Silver
Gold

Potassium is the most reactive metal in this list and gold is the least reactive.

At school students frequently carry out reactions on a test-tube scale to illustrate the reactivity series of metals. One such reaction is the reaction of metals with hydrochloric acid.

In the pictures below magnesium, zinc, iron* and copper are shown reacting with 2M hydrochloric acid.

Four metals used at school to illustrate the reactivity series of metals

What is the order of reactivity of the four metals with 2M hydrochloric acid?

After about a minute magnesium is seen reacting the most vigourously of the four, followed by zinc which is the second most reactive.

After a few tens of seconds

*The ‘iron’ nail is almost certainly not pure iron, but probably a form of steel.

Under these conditions the magnesium completely reacts in a couple of minutes. After about an hour the zinc is seen reacting steadily and bubbles of hydrogen gas can be seen on the surface of the nail.

After about an hour

Thus, the order of reactivity as illustrated by this experiment is:
Magnesium
Zinc
Iron
Copper

Copper is sometimes described as being unreactive in this experiment, but the word unreactive must not be taken to mean ‘does not react at all’. Copper metal does react with 2M hydrochloric acid, only it reacts rather slowly. This is illustrated by the animation below (apologies for the wobbly nature of the timelapse photography here).

Iron and copper with 2M HCl for 1 week

In this experiment, at the end of one week the rather thin piece of copper foil had virtually all gone, whilst the much bigger iron nail, or its remnants, were still seen bubbling away.

Copper foil after one week in 75ml 2M HCl

Iron nail after one week in 75ml 2M HCl

Thus, copper can be described as being relatively unreactive in this reaction with 2M hydrochloric acid, but it does react slowly. Can you suggest any metals which would not react with 2M hydrochloric acid?

Last time quiz crystal 1 was quartz and quiz crystal 2 was iron pyrite or Fool’s Gold. The answer to the third quiz question was neutralisation.