The chemistry of limestone: teacher’s notes
This activity is most appropriate for students aged 14-16 to illustrate chemical reactions and useful materials made from rocks.
In Northern Ireland it also matches KS3 requirements for the above topic.
English National Curriculum reference 4.3.3k
CCEA (Northern Ireland) reference 3.4.1
ACCAC (Wales) reference 220.127.116.11
AQA modular 3468 reference 15.1
AQA linear 3462 reference 11.5
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.
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 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).
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.
Each student (or group) will need:
Each student (or group) will need:
For fuller details of the experiment, see the student's material.
Students take about half a dozen small (about 1cm3) 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:
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.
CaCO3(s) → CaO(s) + CO2(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) + H2O(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) + CO2(g) → CaCO3(s) + H2O(l)
Continuing to blow through the straw for some time will result in the calcium carbonate precipitate re-dissolving as soluble calcium hydrogencarbonate.
CaCO3(s) + CO2(g) + H2O(l) → Ca(HCO3)2(aq)
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
(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
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)