Hard water -- formation of chalk on boiling

In this experiment it is nicely demonstrated what happens when so-called hard water (which contains a lot of calcium ions and bicarbonate ions) is boiled for some time. Everyone knows the formation of chalk, which covers heating spirals, glass walls and many other things and every few months needs to be removed.

The mechanism behind the formation of chalk from hard water is the decomposition of bicarbonate ions when the water is boiled, and the subsequent precipitation of carbonate ions with calcium and magnesium ions, which also are present in the water.

Required chemicals:

• sodium bicarbonate

• calcium chloride

• phenolphtalein pH indicator solution

Required equipment:

• test tubes

Safety:

• Besides the heating of the test tubes, no real risk is involved in this experiment.

Disposal:

• No toxic chemicals are used in this experiment, the waste can be flushed down the drain.

Rise of pH when boiling a solution of sodium bicarbonate

This experiment is very simple. Dissolve some sodium bicarbonate in cold distilled water and add a few drops of phenolphtalein indicator. The solution of sodium bicarbonate is very slightly basic. The phenolphtalein indicator solution has a faint pink color.

When this solution is boiled for a while, then the color of the solution becomes much more intense. This means that the pH of the solution is rising considerable, well beyond the trajectory where phenolphtalein changes from colorless to red/purple. The solution remains deep red/purple after cooling down again.

This experiment demonstrates that the pH is going up when a solution of bicarbonate is heated. The chemical explanation is given at the end of this web page.

Formation of chalk from hard water

In this experiment some very hard water is made and then this is boiled. Making the very hard water can be done as follows:

• Take a few ml of distilled water and put this in a wide test tube.
• Add a few mm³ of solid NaHCO₃. Swirl the test tube, until all of the solid has dissolved. Only take a small amount, just a few mm³ of the powder, not more.
• Dissolve some calcium chloride in another test tube with distilled water and add this solution to the solution of NaHCO₃.

The solution is clear and colorless, as shown in the picture below.

This colorless solution is heated and the liquid is kept boiling for a few minutes. After the boiling, the liquid looks as follows:

Many small solid particles are formed, which tend to stick to the glass and the liquid becomes turbid. Carefully letting settle the solid at the bottle and rinsing along the glass with the liquid above the precipitate results in a small, but clearly visible concentration of white solid at the bottom. Also some solid remains stuck on the glass walls.

Discussion of results

Hard water contains a high concentration of the following ion:

• Ca²⁺ (and possibly Mg²⁺)
• HCO₃^–

In the first experiment, a solution of NaHCO₃ is prepared. This solution is very slightly alkaline, almost neutral. In such a solution the following equilibrium exists (which is largely to the left):

 (1)   HCO₃^– + H₂O ↔ H₂CO₃ + OH^–

There also is another set of equilibria in this solution, which also are largely to the left:

 (2)   HCO₃^– + H₂O ↔ CO₃^2– + H₃O⁺

 (3)   HCO₃^– + OH^– ↔ CO₃^2– + H₂O

This latter two equilibria are less pronounced than the previous one, and this makes a solution of a bicarbonate slightly alkaline.

The compound H₂CO₃ is a weak acid, but more important is that it is in equilibrium as follows:

 (4)   H₂CO₃ ↔ CO₂ + H₂O

At room temperature, the CO₂ remains in solution. When the solution is heated, then the CO₂ becomes less soluble and escapes from the solution as a gas and equilibrium (4) now becomes a reaction which is driven much more to the right:

 (4')   H₂CO₃ → CO₂↑ + H₂O

Because H₂CO₃ is taken away from the system, equilibrium (1) also is driven to the right, resulting in the following:

 (1')   HCO₃^– + H₂O → CO₂↑ + H₂O + OH^–

With the rising concentration of the hydroxide ion, equilibrium (3) is driven more to the right as well, because CO₃^2– is only a fairly weak base and even low concentrations of OH^– drive this to the right. Free H₂CO₃ becomes less abundant and the reaction in which CO₂ is formed comes to a halt. At a certain point, when there hardly is any bicarbonate left, which is the case after a few minutes of boiling, the situation has changed drastically. The main species in solution now is carbonate. Approximately half of the carbon is driven out as CO₂ and the remaining half is present as carbonate, which is in equilibrium with a small amount of hydroxide and bicarbonate. The pH of the solution has risen considerably, because now the main species is CO₃^2– and although this only is a fairly weak base, it is a much stronger base than bicarbonate.

The first experiment demonstrates what is explained above.

Presence of calcium ions

When calcium ions are present in solution, then the calcium ions are precipitated as calcium carbonate. Calcium bicarbonate is soluble in water fairly well, but calcium carbonate has very low solubility. So, the carbonate, which results from heating the solution with the bicarbonate, causes precipitation of calcium (and magnesium) ions.

On heating, the net reaction is

    Ca²⁺ + 2HCO₃^– → CaCO₃↓ + CO₂↑ + H₂O

This effect is demonstrated in the second experiment.

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