TOPIC: Solubility


Lime Water/Carbon Dioxide Reactions

The addition of carbon dioxide gas to lime water, Ca(OH)2(aq), first results in the precipitation of CaCO3. Additional CO2, however, redissolves the precipitate. The reaction can be reversed to give the CaCO3 precipitate once again either by boiling the solution or bubbling air through it.


# dry ice
Saturated Ca(OH)
2 solution
Tall, narrow container (like a 5 L graduated cylinder, unmarked if available)
Carbon dioxide generator (500 mL flask with stopper and glass tube connected to 3 ft of rubber hose ending in 1 ft of glass tubing.
Bunsen burner (or a stirring hot plate, if there is plenty of time)
two demo-size (a.k.a. "Texas-size") test tubes
8 ft of rubber hose to connect to compressed air line


Add equal parts of lime water and distilled water to the tall container. Add water and about an ounce of dry ice to the gas generator. Bubble CO2 through the lime water solution until the precipitate is well established (20-30 seconds), and then remove the tube and discuss the precipitate. Continue to bubble CO2 through the solution until the precipitate just disappears. Fill the test tube about 1/5 of the way full with the solution (or if a hot plate is being used, fill the beaker about 1/3 full and start it stirring on the hot plate). Begin bubbling air through the remainder of the solution, using the hose and glass tube from the generator. Discuss the significance of the re-dissolution and how the reaction can be reversed. Boil the solution for a couple of minutes with the Bunsen burner. The precipitate with reappear almost at once, but it takes about 10-20 minutes for the solution through which air is being bubbled to turn cloudy.


When carbon dioxide dissolves in water, it forms carbonic acid. Lime water neutralizes the carbonic acid and carbonate ion is formed. Calcium carbonate is insoluble

CO2(g) + H2O(l) H2CO3(aq)

H2CO3(aq) + 2OH-(aq) CO32-(aq) + 4H2O(l)

Ca2+(aq) + CO32- CaCO3(s)

and precipitates. As still more CO2 is bubbled into solution, all the OH- is used up, and the solution becomes acidic. Carbonate ion is converted into bicarbonate, which has a very

CaCO3(s) + H2CO3(aq) Ca2+(aq) + 2HCO3-(aq)

soluble calcium salt. In Nature, calcium is a very common ion in igneous rocks. As these rocks weather and the calcium goes into solution, CO2 in the atmosphere precipitates some of it as CaCO3. Found in massive sedimentary deposits, this material is called limestone. Cracks in limestone deposits allow water (containing more CO2) to percolate through, dissolving part of the material and forming limestone caves. Re-precipitation of the CaCO3 results in the formation of stalactites and stalagmites. Tap water that is "hard" because it comes from wells and has been in contact with limestone for a long time contains relatively large amounts of calcium and bicarbonate ions. When this water is heated in coffee pots or water heaters and allowed to cool, limy deposits called "boiler scale" build up.


None. Although Ca(OH)2 is a strong base, it is so insoluble that a saturated solution of it does not even give the bitter taste typical of basic substances.


B. Z. Shakhashiri, "Chemical Demonstrations", Volume 1, p. 329.


Directions for this demonstration normally call for undiluted lime water. However, the use of this solution causes the re-dissolution of the calcium carbonate to be inconveniently delayed.