Small-scale science
Some more microscale gas experiments
Bob Worley
ABSTRACT Microscale techniques can assist chemists to carry out some experiments which on a
large scale would be quite hazardous. Hydrogen–oxygen explosions, reducing metal oxides with
hydrogen and working with toxic gases can all be carried out very quickly and safely once the
techniques have been assimilated and practised.
In 1984, a teacher was successfully prosecuted
by the Health and Safety Executive. The
case centred around chemically prepared
hydrogen that was dried by bubbling it through
concentrated sulfuric(VI) acid and then passed
over hot copper(II) oxide. It was traditional
to burn the excess hydrogen at the outlet, as
shown in the diagram in Figure 1 from an old
textbook (Brown, 1954). The procedure was
often carried out quantitatively to find the mass
of copper in a sample of copper(II) oxide and
the hydrogen acted as an inert atmosphere to
avoid any re-oxidation of the copper. However,
if the demonstrator tried to ignite the excess
hydrogen before all the air was flushed out then
the apparatus exploded. In this case, the bottles
of concentrated sulfuric(VI) acid exploded,
showering the watching pupils (not wearing eye
protection) with acid.
Teachers could also use coal gas for this
experiment, which in many ways was safer. With
natural gas (methane), more heat was required and
it was difficult to explain why the metal oxide was
reduced even when various tricks were used (such
as passing the gas through ethanol).
I remember being told at the school where
I taught that the procedure with hydrogen
was ‘banned by law’. This was not true, as I
subsequently found out when I joined CLEAPSS
(see Websites).
I then saw Bruce Mattson of Creighton
University carry out the reaction in microscale and
was amazed at its simplicity. With the microscale
approach, there is very little dead space for there
to be an explosive atmosphere. A syringe full of
hydrogen can be made by the Mattson method
(see page 43), or by other methods shown at the
end of this account.
Reduction of metal oxides with hydrogen
The procedure for reducing copper(II) oxide with
hydrogen is shown in Box 1.
The same approach can be used with lead(II)
oxide [toxic], iron(III) oxide and nickel(II) oxide
[toxic]. Cobalt (II) oxide [toxic] can also be used;
it is prepared by heating a small amount of cobalt
carbonate in a test tube – wear eye protection,
disposable gloves and use a fume cupboard.
Cobalt compounds are known to be sensitisers.
They are not covered in commonly adopted model
risk assessments such as Hazcards, so a special
risk assessment will be necessary for most schools
(check with CLEAPSS or SSERC – see Websites).
The reactions are not noticeably exothermic
and the flame should be kept on. The products of
Figure 1 Traditional arrangement for reduction of a metal oxide; reproduced from Brown, 1954
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BOX 1 Reduction of copper(II) oxide with hydrogen
1. Wear eye protection. Do not light any spirit
burners or Bunsen burners while the hydrogen
[extremely flammable] is being placed in the
syringes.
2. Copper(II) oxide [harmful] during storage
adsorbs water and is naturally ‘damp’. Heat a
small quantity (say 2 g) of copper(II) oxide in a
borosilicate test tube first and allow it to cool
before using it.
3. Prepare a 60 ml plastic syringe full of hydrogen
[extremely flammable] and fit the syringe cap.
4. Using a microspatula, place a small amount of
dried copper(II) oxide in the Pasteur pipette.
5. Set up the appartus as shown in Figure 2,
clamping around the end of the Pasteur
pipette. The syringe can be left loose; it does
not fall off.
6. Light the spirit burner. A Bunsen burner flame
is too hot and would cause the Pasteur
pipette to bend as the glass softens. A tea
light could be used but is perhaps too weak,
and carbon would be deposited over the
pipette, preventing the reaction from being
observed. Experience has shown that small
spirit burners produce the best results.
7. After about 2 minutes, extinguish the flame, hold
the syringe in one hand and push the barrel to
force hydrogen over the hot copper(II) oxide.
8. The reaction is exothermic, causing the solid
to glow. Water can be seen condensing at the
exit of the pipette (Figure 3).
9. Let the apparatus cool before disconnecting
the pipette.
Figure 3 The exothermic reaction between
copper(II) oxide and hydrogen. The spirit burner
flame is extinguished. There are water droplets on
the right-hand side of the tube.
Figure 2 Arrangement for reduction of a metal oxide on a small scale
the iron, nickel and cobalt oxide reductions are all
magnetic (Figure 4).
Another variation of performing this reaction
is shown in the CLEAPSS guide L195, Safer
chemicals, safer reactions. A diagram of the
arrangement is shown in Figure 5. Unfortunately,
this method does not show the exothermic nature
of the reaction.
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Hydrogen–oxygen explosions
One of the first accident reports I received after
joining CLEAPSS concerned a teacher who
exploded a large oxygen/hydrogen balloon.
With a radius of about 15 cm (which means the
volume was about14 dm3) and under pressure,
the explosion was so loud that both he and the
pupils were deafened for over 24 hours and
all the windows were blown out. I am always
Worley
Some more microscale gas experiments
(a)
(b)
(c)
(d)
Figure 4 Reduction of four other metals by hydrogen: (a) lead(II) oxide reduced to lead; (b) iron(III) oxide
reduced to magnetic iron; (c) nickel(II) oxide reduced to magnetic nickel; (d) cobalt(II) oxide reduced to
magnetic cobalt
(a)
mineral wool
copper(lI) oxide
spirit burner
vial
2 mol dm Ñ3
hydrochloric acid
zinc
(b)
indebted to Alan Goodwin for showing me the
microscale method to carry the procedure out
safely (Box 2). All I did was to use nickel rather
than expensive platinum electrodes. The nickel
anode will oxidise slightly but not enough to spoil
the experiment. It gives a really good ‘crack’ with
only about 10 cm3 of gas.
I have exploded larger volumes of hydrogen
and oxygen in soap bubbles. These were made by
combining hydrogen and oxygen from syringes
in a 2 : 1 ratio by volume. With volumes from 10
to 60 cm3, I would recommend that earplugs be
worn by the demonstrator and that the audience
is at least 10 m away, with their hands over their
ears. The pressure wave is such that I have found
Bunsen burner flames close by (to light the splint)
are extinguished.
Microscale hydrogen–oxygen explosions for
pupils to carry out
Microscale allows pupils themselves to carry out
hydrogen–oxygen explosions, with only 4 cm3 of
gas (Box 3).
Bruce Mattson has developed a piezoelectric
sparker that allows pupils to set these off as
minirockets (see Websites).
Figure 5 (a) Arrangement for an alternative
CLEAPSS-approved method of reducing copper
oxide; (b) the resulting copper mirror
Reactions of toxic gases in a Petri dish
The Royal Society of Chemistry book on
microscale chemistry (Skinner, 1998) described
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BOX 2 Microscale method for demonstrating hydrogen–oxygen explosions
Apparatus
A 100 ml wide-necked glass bottle is fitted with a
rubber bung into which two holes are drilled with
a very fine drill bit. A third, wider, hole is bored
though the bung, which is then fitted with 6 mm
medium-wall borosilicate glass tubing.
Two copper wires are fed through the two small
holes in the bung and nickel-foil electrodes are
soldered onto the wire.
The glass tubing is bent into shape using a nonluminous Bunsen burner flame to melt the glass;
be careful not to heat the bung.
The nickel anode does oxidise during the process
if the electrolysis cell is left on for long periods.
The solution becomes green and green nickel(II)
hydroxide precipitates out.
All solutions can be poured down the foul-water
drain.
The electrolyte is 0.2 mol dm−3 sodium sulfate(VI)
solution rather than sulfuric(VI) acid. More
advanced students will appreciate that water
molecules are being oxidised and reduced at
the electrodes to form oxygen and hydrogen
molecules.
(a)
Method
1. Wear eye protection.
2. Pour a 0.2 mol dm−3 sodium sulfate(VI) solution
into a 100 ml bottle so that it reaches the very
top of the bottle.
3. Have a beaker or tray handy (or work over
a sink) to collect the overspill as the bung
containing the electrodes is inserted into the
neck of the bottle (Figure 6a). The overflow
rises up the delivery tube and empties into the
beaker or tray. There must be no dead space
above the level of the liquid.
4. Connect the copper wires to a dc low-voltage
electrical supply (6–8 V is a suitable setting)
and pass current until no more solution is
pushed out of the bottle into the beaker or
tray. The tube is then full of gas.
5. Quickly place a crucible or a small weighing
boat filled with soap solution under the tube so
that the end of the tube dips into the solution,
and support this on a laboratory jack or
wooden blocks so that gases from the bottle
will bubble through. Switch on the current to
collect bubbles of gas (Figure 6b).
6. Switch off the low-voltage supply, lift the bottle
so that the tube is clear of the crucible and move
the crucible closer to an ignited Bunsen burner.
7. Light a splint with the Bunsen burner flame and
apply it to the bubbles on the top of the crucible.
(b)
Figure 6 (a) Diagram and (b) photograph of the apparatus for the microscale preparation and explosion
of hydrogen
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Some more microscale gas experiments
BOX 3 Microscale hydrogen–oxygen explosions that pupils can carry out themselves
1. Wear eye protection.
2. Prepare two labelled 60 ml syringes sealed
with syringe caps, one containing oxygen
[oxidising] and the other containing hydrogen
[extremely flammable].
3. Cut off the bulb from a nominal 3 ml plastic
pipette with scissors.
4. With another pipette, add water in 1 ml aliquots,
marking the level it comes to in the cut-off bulb
each time with a felt-tip pen. These pipettes
usually hold about 4 ml (Figure 7a).
5. Fill the bulb with water. Turn it upside down so
that the surface tension effect holds the water
in place. Hold the bulb in a clamp which is just
tight enough to grip the bulb.
(a)
(b)
6. To fill the pipette with gas, remove the syringe
cap from the syringe and attach the silicone
tubing. Just push the plunger slightly (holding
the stem of the plunger and not pushing the
end) and insert the tube up the inverted bulb.
Insert the silicone tube into the pipette bulb.
Add the gas and note the approximate volume
in the bulb (Figure 7b). Do this with oxygen first
to the 1 ml mark, then make it up to the 4 ml
mark with hydrogen using the same method.
7. Bring the bulb close to a gentle flame from a
Bunsen burner. Squeeze the bulb to blow out
the water and a little of the gas. The mixture
then ignites with a bang if the gas ratios are
correct. Figure 7c is a frame from a movie of
the explosion.
(c)
Figure 7 (a) The plastic pipette bulb temporarily filled with water; (b) filling the pipette bulb with gas;
(c) the hydrogen–oxygen explosion when the gas mixture is squeezed out of the pipette bulb near a
Bunsen burner flame
some elegant experiments with toxic gases in
Petri dishes. Now that plastic Petri dishes are
available in most schools, these make an excellent
piece of kit to use. These experiments with toxic
gases can be carried out in the open laboratory
as most of the gas is consumed in the reactions
(Boxes 4 and 5).
The reaction to produce the gas is carried out
in a blister pack well. Drugs are often supplied in
blister packs, one side being plastic and the other
aluminium foil. Pressing the plastic side ejects the
pill through the aluminium side. After dispensing,
the indented plastic sheet can be cut up and used
for tiny reaction vessels in this and other similar
experiments.
Other gases can be examined in a similar way,
as shown in Table 1.
Other methods of obtaining hydrogen
Syringes filled with hydrogen [extremely
flammable] can be prepared from hydrogen gas
cylinders and canisters (Figure 12).
Alternatively, syringes can be filled from
a zinc/acid chemical generator. However, it is
imperative that all air is first displaced from the
bottle. I ensure this by using a 250 ml plastic
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BOX 4 The chemistry of ammonia
Procedure
1. Wear eye protection. The base of a
plastic 90 mm Petri dish is used.
2. Place the reagents around the base
of the Petri dish as shown in Figure 8.
Pieces of indicator paper are also
placed in the Petri dish and these
undergo colour changes as the pH
changes (Figure 9a).
3. Add about 0.3 ml of freshly made
2 mol dm−3 ammonia solution to a blister
pack well in the centre and replace the
top of the Petri dish.
Comments
After about a minute, the indicator papers
will have changed colour and precipitates
will be forming in the various solutions
(Figure 9b).
After 10 minutes, the acid solution is
gradually being neutralised and the
precipitates are fully formed (Figure 9c).
After even more time, the
tetraaminecopper(II) complex,
[Cu(NH3)4(H2O)2]2+, begins to form.
0.3 ml of freshly made 2 mol dm−3 ammonia
solution has the capacity to produce about
21 cm3 of ammonia gas [toxic] The reaction
(a)
(b)
2 drops of 0.1 mol dm−3
hydrochloric acid plus
1 drop of universal
indicator
2 drops of 0.1 mol dm−3
lead(II) nitrate solution
[toxic]
2 drops of
0.1 mol dm−3 zinc
sulfate(VI) solution
5 small grains of iron(II)
sulfate(VI) crystals plus
3 drops of water
blister pack well
containing 0.3 ml of
2 mol dm−3 ammonia
solution
2 drops of 0.1 mol dm−3
copper(II) sulfate(VI)
solution
3 drops of
0.1 mol dm−3 iron(III)
chloride solution
damp pH papers
with ranges 1–14,
8–10 and 10–12
Figure 8 Placement of reagents and indicator papers in
the Petri dish
can be speeded up by adding anhydrous calcium
chloride [irritant] to the blister pack well. The exothermic
reaction with water increases the temperature and
releases the gas more rapidly.
The solutions can be washed down the sink.
(c)
Figure 9 Reactions of ammonia gas with droplets of various solutions: (a) before the addition of the
ammonia solution to the central blister pack well; (b) after about 1 minute; (c) after 10 minutes
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Some more microscale gas experiments
BOX 5 The chemistry of sulfur dioxide
Procedure
1. Wear eye protection. The gas is toxic
and those with asthma should take
particular care. The base of a plastic
90 mm Petri dish is used.
2. Place the reagents around the base of
the Petri dish as shown in Figure 10.
Pieces of indicator paper are also
placed in the Petri dish and these
undergo colour changes as the pH
changes (Figure 11a).
3. Add about 0.04 g of sodium
metabisulfite to a blister pack well in
the centre, add three or four drops of
1 mol dm−3 hydrochloric acid [irritant],
and replace the top of the Petri dish.
Comments
After about a minute, the indicator papers
will have changed colour, the iodine will
have decolourised and the potassium
manganate(VII) will have reduced to
manganese(IV) oxide (Figure 11b).
After 10 minutes, the potassium
manganate(VII) will have reduced further
to the manganese(II) state, and the
dichromate(VI) to the chromium(III) state
(Figure 11c). There is no precipitate in
the plain barium chloride solution (barium
sulfate(IV) is soluble in water) but the solution
that also contained hydrogen peroxide does
(a)
(b)
3 drops of 0.002 mol dm−3
potassium dichromate(VI)
solution
2 drops of 0.1 mol dm−3
3 drops of 0.005 mol dm−3
barium chloride solution
acidic
potassium
[harmful] plus 1 drop of
manganate(VII) solution
20-volume hydrogen
peroxide solution
[irritant]
blister pack well
3 drops of
containing 0.04 g of
2 drops of
0.05 mol dm−3
sodium metabisulfite
−3
0.1 mol dm
iodine solution
[harmful] to which
barium chloride
3 drops of 1 mol dm−3
solution [harmful]
hydrochloric acid [irritant]
are added
damp pH papers
with ranges 1–14,
1–4 and 4–6
Figure 10 Placement of reagents and indicator papers
in the Petri dish
form a precipitate of barium sulfate(VI): the hydrogen
peroxide oxidises sulfuric(IV) acid to sulfuric(VI) acid.
About 10 cm3 of sulfur dioxide gas [toxic] is produced
from the reaction in the blister pack well.
At the end of the experiment, place the Petri dish in a
fume cupboard and remove the top. The solutions can
be washed down the fume cupboard sink.
(c)
Figure 11 Reactions of sulfur dioxide gas with droplets of various solutions: (a) before the addition of
the sodium metabisulfite and hydrochloric acid to the central blister pack well; (b) after about 1 minute;
(c) after 10 minutes
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Table 1 Other gases whose reactions can be investigated in a Petri dish
Gas
Carbon dioxide
Preparation of 10 cm3 of gas
Add about 0.2 ml of 2 mol dm−3
hydrochloric acid [irritant] to
marble chips
Reactions
Calcium hydroxide
solution; 0.01 mol dm−3
sodium carbonate
solution with universal
indicator; 0.1 mol dm−3
barium chloride solution
[harmful]; 0.1 mol dm−3
sodium hydroxide
solution [irritant]
Comments
When carbon dioxide
dissolves into a mixture of
barium chloride and sodium
hydroxide, a precipitate of
barium hydroxide [corrosive]
forms. Barium chloride
solution on its own forms no
precipitate.
Hydrogen sulfide
Add 0.5 ml of 2 mol dm−3
[extremely flammable hydrochloric acid [irritant] to
and toxic]
iron(II) sulfide
Metal salt solutions
produce precipitates;
acidified potassium
manganate(VII) is
reduced
The odour is off-putting, to
say the least. The gas level
for a class experiment is well
below the WEL level* for
hydrogen sulfide. Only open
the Petri dish in the fume
cupboard.
Chlorine [toxic]
Potassium iodide and
potassium bromide
solutions; moist blue
litmus is bleached
Although only small amounts
of chorine are formed, the
Petri dish should only be
opened in a fume cupboard.
Add about 0.4 ml of 2 mol dm−3
hydrochloric acid [irritant]
to fresh bleaching powder
[oxidising and corrosive]
*Workplace Exposure Limits (WELs) are specified by the Health and Safety Executive: They define the extent to which a person
may be safely exposed to a hazardous substance (typically a gas or solvent vapour) without endangering his or her health.
Figure 12 Obtaining hydrogen
from a gas cylinder
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(a)
Some more microscale gas experiments
(b)
Figure 13 (a) Apparatus to obtain hydrogen from a zinc/acid generator; (b) checking the liquid levels to
ensure that the hydrogen in the cylinder is at atmospheric pressure
bottle as the generator and then collecting over
250 cm3 of displaced gas over water first before
attaching the syringe. The equipment is set up
as shown in Figure 13a. The syringe barrel will
not move on its own. The pressure of hydrogen
rises above atmospheric pressure and the level of
acid rises in the thistle funnel as shown. Once the
liquid rises into the funnel, withdraw the syringe
barrel a little at a time, which brings the level of
acid down in the tube; always keep the level of
liquid in view (Figure 13b). Keep repeating this
until the syringe is full and then attach a syringe
cap to the syringe. Gases, including hydrogen,
will keep for several days in the syringe without
diffusing away.
References
Websites
Brown, G. I. (1954) Essentials of certificate chemistry.
London: Longmans, Green & Co.
Skinner, J. ed. (1998) Microscale chemistry: experiments in
miniature. Cambridge: Royal Society of Chemistry.
CLEAPSS (provides safety advice in the UK excluding
Scotland): www.cleapss.org.uk
Microscale Gas Chemistry: Experiments with oxygen
– Experiment 4. Hydrogen–oxygen rockets: mattson.
creighton.edu/O2/index.html
SSERC (provides safety advice in Scotland): www.sserc.
org.uk
Bob Worley is the lead chemistry adviser at CLEAPSS and the guest editor for the small-scale science
theme in this issue. Email: [email protected]
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