Apart from carbon dioxide what is the other product of thermal decomposition of calcium carbonate?

Calcium carbonate is the principal mineral component of limestone. Its chemical and physical properties lie behind the modern-day uses of limestone as well as the unique limestone landscapes of the countryside.

Calcium carbonate – mineral forms

The principal mineral component of limestone is a crystalline form of calcium carbonate known as calcite. Although calcite crystals belong to the trigonal crystal system, shown below, a wide variety of crystal shapes are found.

Single calcite crystals display an optical property called birefringence (double refraction). This strong birefringence causes objects viewed through a clear piece of calcite to appear doubled.

Another mineral form of calcium carbonate is called aragonite. Its crystal lattice differs from that of calcite, resulting in a different crystal shape – an orthorhombic system with needle-shaped crystals.

Solubility

Calcium carbonate has a very low solubility in pure water (15 mg/L at 25°C), but in rainwater saturated with carbon dioxide, its solubility increases due to the formation of more soluble calcium bicarbonate. Calcium carbonate is unusual in that its solubility increases as the temperature of the water decreases.

The increased solubility of calcium carbonate in rainwater saturated with carbon dioxide is the driving force behind the erosion of limestone rocks, leading to the formation over long periods of time of caverns, caves, stalagmites and stalactites. Rainwater is weakly acidic, and when it meets with limestone, some of the calcium carbonate reacts to form a solution of calcium bicarbonate.

Thermal decomposition

When heated above 840°C, calcium carbonate decomposes, releasing carbon dioxide gas and leaving behind calcium oxide – a white solid.

Calcium oxide is known as lime and is one of the top 10 chemicals produced annually by thermal decomposition of limestone.

The thermal decomposition of calcium carbonate to lime is one of the oldest chemical reactions known. For several thousand years, lime has been used in mortar (a paste of lime, sand and water) to cement stones to one another in buildings, walls and roads. The setting of mortar involves several chemical reactions.

First, the lime is ‘slaked’ by the water to produce calcium hydroxide (slaked lime

Over time, this reacts with carbon dioxide in the air to form crystals of calcium carbonate, which lock the sand grains together to form a hard rock-like material.

Reaction with acids

Like all metal carbonates, calcium carbonate reacts with acidic solutions to produce carbon dioxide gas. It is this reaction that is responsible for limestone fizzing when dilute hydrochloric acid is placed on its surface.

Limestone, which consists mostly of calcium carbonate, has been used in agriculture for centuries. It is spread on fields to neutralise acidic compounds in the soil and to supply calcium, which is an essential plant nutrient. Today, depending on the soil requirements, options available to the farmer are:

• lime – CaO
• slaked lime – Ca(OH)2
• crushed pure calcitic limestone – CaCO3
• dolomitic limestone – CaMg(CO3)2

In medicine, antacids containing small amounts of calcium carbonate are used in the treatment of ‘acid stomach’. The chemical reaction occurring involves the neutralisation of excess acid with calcium carbonate. Brands such as Quick-Eze and TUMS have calcium carbonate as the ‘active’ ingredient.

In trying to understand the world around us, scientists often look for patterns of behaviour that allow general rules or principles to be formulated. The watchful scientist, however, needs to have an open mind and know that there are always exceptions to the general rule, for example, where the solubility of calcium carbonate decreases with increasing temperature.

Calcium carbonate is strongly heated until it undergoes thermal decomposition to form calcium oxide and carbon dioxide. The calcium oxide (unslaked lime) is dissolved in water to form calcium hydroxide (limewater). Bubbling carbon dioxide through this forms a milky suspension of calcium carbonate

This experiment can be carried out conveniently in groups of two or three and takes about 40–45 minutes.

Equipment 

Apparatus

• Eye protection
• Tripod
• Gauze
• Bunsen burner
• Tongs
• Boiling tubes, x2 (note 1)
• Drinking straw (note 2)
• Dropping pipette
• Filter funnel, small
• Filter paper

Apparatus notes

1. Use large (150 x 25 mm) test tubes (boiling tubes).
2. Freshly purchased drinking straws should be used and each student issued with their own straw.

Chemicals

• Calcium carbonate
• Universal Indicator solution (HIGHLY FLAMMABLE)

Health, safety and technical notes

• Read our standard health and safety guidance.
• Wear eye protection.
• Calcium carbonate, CaCO3(s) – see CLEAPSS Hazcard HC019b. The calcium carbonate used should be in the form of pea sized lumps of chalk. Blackboard chalk should not be used as it is likely to be mostly calcium sulfate.
• Universal indicator solution (HIGHLY FLAMMABLE) – see CLEAPSS Hazcard HC032 and CLEAPSS Recipe Book RB047.

Procedure

1. You need to prepare a tabulated results sheet before you start your experiments. An example table is provided below in the teaching notes.
2. Set a lump of chalk (calcium carbonate) on a gauze. If your gauze has a coated central circle, use the edge where there is no coating.
3. Heat the chalk very strongly for 5–10 minutes. Write down what you observe.
4. Let the chalk cool and use tongs to move it into a boiling tube. Add 2–3 drops of water with a dropping pipette. Write down your observations.
5. Add about 10 cm3 more water to the solid. What happens now?
6. Filter half the mixture into the other boiling tube and, using a straw, gently blow a stream of bubbles through the filtrate. What do you see?
7. Test the remaining half of the mixture with Universal Indicator solution. Write down what you observe.

Teaching notes

Keep an eye on less mature students who might be tempted to suck rather than blow through the filtrate.

The results expected are as follows:

This set of experiments involves a variety of important reactions and types of reactions, with several references to industrial processes. The roasting of limestone and the hydration of the quicklime formed has relevance in the manufacture of plaster and cement, and in the laboratory limewater is a common reagent for the testing of carbon dioxide. Students could be asked to carry out web research on these applications.

Some question and answers for the class after the experiment

1. Why does the chalk crumble slightly on strong heating?
   Carbon dioxide/a gas is evolved; this forces its way out of the solid and breaks down its structure.
2. What type of reaction is taking place during the heating process? Write an equation for the reaction.
   Thermal decomposition; CaCO3(s) → CaO(s) + CO2(g)
3. Why is steam evolved when drops of water are added? Write an equation for the reaction occurring.
   The reaction is highly exothermic and the small amount of water added is partly converted to steam in the process: CaO(s) + H2O(l) → Ca(OH)2(s)
4. Why does the limewater turn cloudy? Write an equation for the reaction which is occurring.
   Insoluble calcium carbonate is being precipitated: Ca(OH)2(aq) + CO2(g) → CaCO3(s) + H2O(l)
5. What does the colour change occurring when limewater is added tell you about the pH of the solution? Explain why the pH would be expected to have this value.
   The pH is about 11 - 14; soluble metal hydroxides are alkaline and therefore give high pH values.