teacher’s notes
student’s notes
The chemistry of
limestone: teacher’s
notes
Level
This activity is most appropriate for students aged 14-
16 to illustrate chemical reactions and useful materials
made from rocks.
I
Topic
This activity illustrates some of the simple chemical
reactions of limestone (calcium carbonate, CaCO
3
) and
lime (calcium oxide, CaO).
Description
The activity is suitable as a class practical or as a
demonstration. The students (or, less preferably, the
teacher) heat limestone (mainly calcium carbonate) to
form lime (calcium oxide) and note the differences
between the reactions of limestone and of lime with
water, acids and carbon dioxide.
Context
Students should know about differentiation of
materials, eg rocks, on the basis of physical properties,
and the activity assumes that some work has already
been done on the physical examination of rocks.
Students should know that carbon dioxide is a gas and
have simple ideas about reversible and irreversible
changes. They should be aware of simple properties of
acids, alkalis and indicators.
The activity concentrates almost exclusively on
chemistry, but there are also important potential links
with fossils and evolution.
Teaching points
The chemical and physical properties of limestone,
especially when reasonably pure, make it highly sought
after for hundreds of everyday uses. With salt and coal,
it formed the main feedstock for the chemical industry
until about 1914. It is still important today as shown by
the wide range of uses in Tables 1 - 4 (Appendix). The
chemistry is relatively straightforward and can be used
to illustrate many types of simple reactions and
properties. It is then possible to relate these to
industrial and domestic applications (see Limestone in
everyday life).
Timing
It should be possible to carry out the activity either as a
class practical or as a demonstration within a teaching
period of about one hour.
Apparatus
Each student (or group) will need:
eye protection
Bunsen burner, tripod and gauze
heatproof mat
tongs
3 test-tubes
test-tube rack
dropping pipette
drinking straw
Chemicals
Each student (or group) will need:
a few small lumps of limestone (each about
1cm
3
) (limestone includes chalk). Marble chips
will do if no local source of limestone is
available
deionised / distilled water
Universal Indicator solution and colour chart
Safety notes
Wear eye protection.
Take care when heating as the lumps will become
very hot.
Calcium oxide (lime), the material formed when
the lumps are heated, is corrosive. It causes
burns and is irritating to eyes, skin and the
respiratory system. The reaction of calcium
oxide with water is vigorous and exothermic.
It is the responsibility of the teacher to carry out
an appropriate risk assessment.
The activity
For fuller details of the experiment, see the student's
material.
Students take about half a dozen small (about 1cm
3
)
lumps of limestone. They examine the stone and
describe briefly its colour, texture and any other
notable features such as fossils. The colour of a piece
of limestone may be misleading. For example coarse
brown limestones may be wrongly described as
sandstones. (Limestone comes in almost every
imaginable hue – from white, through yellows, reds,
oranges, blues, purples, olives to browns and black.)
These colour variations are almost all due to iron
content. Some of the darker colour may be due to
carbon or possibly manganese. (If no local source of
limestone is available, marble chips, available from the
prep room, will do.)
Students heat a couple of lumps on a tripod and gauze
with a roaring Bunsen flame for 15 minutes. If possible
darken the room briefly to allow students to note what
happens when the flame is trained directly on the
lumps. It may be possible to see the lumps glowing –
this is the origin of the term ‘limelight’.
After allowing the lumps to cool, students compare the
heated lumps with unheated ones.
Lumps that have been heated:
may appear whiter than the unheated ones
should crumble more easily than the unheated
ones
will react exothermically when a few drops of
water are added
will show an alkaline pH
Blowing through a straw into the clear solution formed
by reacting the heated lumps with water will turn the
solution cloudy.
Note. If it is necessary to spread the practical work over
two teaching periods, teachers should be aware that,
in the intervening period, the lime (calcium oxide)
produced by heating the limestone may combine with
carbon dioxide from the air to re-form calcium
carbonate, thus reducing its reactivity very significantly.
It would be worth making some fresh lime just before
the second lesson.
The chemistry of the reactions is as follows:
Heating the limestone (calcium carbonate) drives off
carbon dioxide gas leaving behind lime, the base
calcium oxide.
CaCO
3
(s) → CaO(s) + CO
2
(g)
The lime is white and will have a more crumbly texture
than the original limestone.
Calcium carbonate does not react with water.
Adding water to the lime produces slaked lime (calcium
hydroxide) in an exothermic reaction.
CaO(s) + H
2
O(l) → Ca(OH)
2
(s)
Some of the calcium hydroxide dissolves in the water
producing an alkaline solution called limewater.
Ca(OH)
2
(s) + (aq) → Ca(OH)
2
(aq)
On blowing into this solution through a straw, the
calcium hydroxide solution reacts with the carbon
dioxide in exhaled breath to form a cloudy precipitate
of calcium carbonate (this is the basis of the limewater
test for carbon dioxide). In effect, we have regenerated
the original limestone.
Ca(OH)
2
(aq) + CO
2
(g) → CaCO
3
(s) + H
2
O(l)
Continuing to blow through the straw for some time
will result in the calcium carbonate precipitate re-
dissolving as soluble calcium hydrogencarbonate.
CaCO
3
(s) + CO
2
(g) + H
2
O(l) → Ca(HCO
3
)
2
(aq)
Appendix: limestone
data for Great Britain
and Northern Ireland
Table 1 The uses of limestone in Great Britain and
Northern Ireland(1999)
Notes on Table 1
(a) mainly foundation and fill
(b) ie architectural, walling, dimension stone
(c) mainly iron- and steel-making flux
(d) powders + ‘whitings’ used in animal feeds, polymers
(plastics, rubber) paint, paper, pharmaceuticals
(e) estimated
(f) in addition about 1.8 Mt of dolomite were used for
industrial purposes (especially furnace linings and
production of magnesium compounds, notably magnesia)
* construction aggregates total = 76 326
n.e.s. not elsewhere specified
Sources: British Geological Survey, Minerals Year Book;
Office for National Statistics; National Stone Centre
Table 2 Production of limestone in Great Britain
and Northern Ireland (1999) by country
Notes on Table 2
(a) almost all for aggregates – figure includes hard chalk
Table 3 Production of limestone in Great Britain
and Northern Ireland (1999) by producing area
Notes on Table 3
(a) N.B. recent data in some cases published for counties
which were reorganised in 1990s
N.B. all figures in Tables 1 – 3 (except N.Ireland) are for
limestone excluding chalk
Table 4 Production and uses of chalk in England
(1999)
Use
Quantity /
kt
Construction
Roadstone coated
9175*
Roadstone uncoated
22 481*
Railway ballast
99*
Concrete
15 309*
Other
29 262* (a)
Cement
9831
Building stone
301 (b)
Asphalt filler / mine dust
216
Building lime
460
Industrial
Agricultural /
horticultural
795
Iron & steel
3239 (c)
Specialist fillers
875 (d)
Soda ash
1000 (e)
Sugar refining
250 (e)
Glass
203
Other lime n.e.s.
139
Other uses n.e.s.
666
Total
94 547 (f)
Country
Quantity /
kt
England
72 820
Wales
17 220
Scotland
1507
N.Ireland
4219 (a)
Total
98 766
Producing area (a)
Quantity /
kt
Derbyshire (inc. Peak National Park)
19 240
Somerset
11 550
N.Yorkshire
7528
Clwyd
7269
Mid Glamorgan
5076
Lancashire
5072
Avon
4948
Durham
4401
Cumbria
4389
Leicestershire
3419
Region
Quantity /
kt
England
9667
Of which: South East Region
Yorkshire/Humber Region
4144
3268
Of which: cement
construction
Misc. uses (inc. fillers)
6345
1021
1701