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Hardness of Water






Waters that do not readily form a lather with soap are termed hard. Such Hardness of Water may be caused by the presence in considerable quantity of salts of the alkali metals, as, for example, in the case of sea water and brine. More usually, however, the term is reserved for such hardness as is due to the presence of very much smaller quantities of salts of magnesium or calcium.

The majority of vegetable and animal oils and fats consist essentially of an organic salt, comprising glycerine, C3H5(OH)3 or CH2OH.CHOH.CH2OH, as base, combined with an organic acid such as stearic acid, C17H35.COOH, oleic acid, C17H33.COOH, or palmitic acid, C15H31.COOH. Thus, glyceryl stearate or stearin, the essential constituent of mutton fat, is represented by the formula

(C17H35.COO)3C3H6,
or


Olive oil is largely glyceryl oleate, and palm oil, glyceryl palmitate. When warmed with solutions of caustic alkalies, these fats are decomposed, yielding soaps, hence the term saponification. Thus, for example, with sodium hydroxide, glyceryl stearate yields free glycerine and sodium stearate, which latter is a sodium soap. Thus

(C17H35.COO)3C3H5 + 3NaOH = C3H5(OH)3 + 3C17H35.COONa
Or
Glyceryl stearate. = Glycerine + Sodium stearate or soap.

The sodium soap is soluble in water and a very small quantity suffices to produce a lather if the water is pure. If, however, it contains dissolved salts of calcium or magnesium the lather is destroyed by these, yielding the familiar insoluble curd, so characteristic of the action of hard water on soap. This curd is really the insoluble soap of the alkaline earth metal formed by double decomposition as shown in the two following equations, in which it is assumed the hardness is due to the presence of calcium carbonate and magnesium sulphate respectively .

Ca(HCO3)2 + 2C17H35.COONa = 2NaHCO3 + (C17H35.COO)2Ca;
and
MgSO4 + 2C17H35.COONa = Na2SO4 + (C17H35.COO)2Mg.

Not until the whole of the alkaline earth metal has been "fixed" as insoluble curd is the sodium soap free to yield a lather. Consequently the higher the percentage of calcium or magnesium, the larger the amount of soap required and the greater the hardness of the water. The amount of soap required to produce a lather is thus a measure of the hardness of the water and, as indicated below, is used in quantitatively determining the same.

Two kinds of hardness are ordinarily recognised, namely temporary and permanent.


Temporary hardness

Temporary hardness is that caused by the presence of the bi- carbonates of calcium or magnesium. Whilst the normal carbonates of these metals, namely, CaCO3 and MgCO3 respectively, are practically insoluble in pure water, they readily dissolve in the presence of carbon dioxide - a normal constituent of natural waters, owing to the presence of this gas in the atmosphere - the soluble acid- or bi-carbonates, CaH2(CO3)2 and MgH2(CO3)2, being produced. The mere process of boiling temporarily hard water suffices to soften it, for the expulsion of the dissolved gases effects the decomposition of the bicarbonates with consequent precipitation of the insoluble normal carbonates.

Permanent hardness

Permanent hardness is caused by soluble salts of calcium and magnesium other than the bicarbonates. The more common of these are the sulphates, chlorides, and nitrates, particularly the first named. Such waters cannot be softened by merely boiling.

The total hardness of a given sample of water may be due in part to the presence of bicarbonate and in part to the presence of other soluble salts. When boiled, the normal carbonate is precipitated and, on account of the decrease in solubility of calcium sulphate with rise in temperature above 38° C., there is always a tendency for this substance to separate to some extent with the carbonate. This causes the deposit to form a coherent film on the containing vessel, whereas the pure carbonate gives a more or less powdery suspension. The boiled water is now softer than before, such hardness as it now possesses is termed permanent, whilst its temporary hardness is the difference between the total and permanent hardness, namely, that lost by boiling.

Degree of Hardness

To render comparison easy it is usual to record the hardness in terms of the calcium oxide, CaO, or calcium carbonate, CaCO3, that would produce the same amount of hardness if added to pure water.

In British water reports, according to a decision of the Local Government Board, the number of grains of CaCO3 required per gallon (70,000 grains) of pure water to render it as hard as a given sample, constitutes the degree of hardness or the degree Clark of the sample. Thus a water of 5 degrees of hardness would be obtained by dissolving 5 grains of CaCO3 per gallon of distilled water. According to the metric system, it is usual to express the hardness in terms of CaCO3 per 100,000 of water. This figure is readily obtained from the degree Clark by dividing the latter by 0.7. Thus

1 degree Clark = 1.43 parts of CaCO3 per 100,000.

In Germany it is customary to express the hardness in terms of 10 mg. of CaO per litre. Hence

1 degree German = 1.79 parts CaCO3 per 100,000.

Determination of Hardness

Hardness, whether temporary or permanent, is conveniently estimated by means of Clark's soap test, which consists in adding from a burette small quantities of standard soap solution (vide infra) to 50 c.c. of water which have been carefully measured out with a pipette into a 250-c.c. bottle. After each addition of soap solution the bottle is vigorously shaken, and the titration is complete when the lather remains unbroken for five minutes after laying the bottle on its side at rest.

Should the water he so hard that 8 c.c. upwards of soap solution are required, a smaller volume than 50 c.c. should be taken and diluted to this amount with freshly boiled distilled water. Reference to the table indicates the amount of hardness corresponding to each titration. The total hardness is given by this method. If a second sample of water be boiled and after settling or filtering titrated in a similar manner, the permanent hardness is obtained. Subtraction gives the temporary hardness.

Degrees of hardness (Clark) corresponding to amounts of soap solution used

Soap Solution, c.c.CaCO3 per 100,000.Soap Solution, c.c.CaCO3 per 100,000.Soap Solution, c.c.CaCO3 per 100,000.
0.70.003.03.255.66.86
0.80.163.23.515.87.14
0.90.323.43.776.07.43
1.00.483.64.036.27.71
1.20.793.84.296.48.00
1.41.114.04.576.68.29
1.61.434.24.866.88.57
1.81.694.45.147.08.86
2.01.954.65.437.29.14
2.22.214.85.717.49.43
2.42.475.06.007.69.71
2.62.735.26.297.810.00
2.82.995.46.578.010.30

Clark's Standard Soap Solution

This can be prepared in several different ways. Commonly dry Castile soap is dissolved in 80 per cent, alcohol in such proportions as will yield a solution well above the desired final concentration; 100 grams per litre is a convenient ratio. After allowing this solution to stand at rest for several days for the deposition of undissolved matter, a quantity of the clear liquid is withdrawn (usually 75-100 c.c. per litre of final solution), and so diluted with 80 per cent, alcohol as to produce a solution which on titration with a known weight of calcium chloride solution under the standard conditions will give results in accordance with Clark's table. The calcium chloride solution is best prepared by dissolving 0.2 grams of Iceland spar in dilute hydrochloric acid; excess of acid is removed by evaporation on a water bath and the solution then diluted to 1 litre with distilled water. A mixture of 25 c.c. of this solution, mixed with 25 c.c. of water, should require 7.8 c.c. of Clark's standard soap solution for the production of a permanent lather.

The standard solution of soap can also conveniently be prepared by neutralising an alcohol solution of oleic acid with a solution of potassium hydroxide in the same solvent; the neutral solution of potassium oleate is then suitably diluted. A solution of potassium soap for dilution can also be obtained by the interaction of lead plaster (" lead soap ") and potassium carbonate.

Although the soap method is still widely applied to the determination of hardness, it is inferior in accuracy and general trustworthiness to more recent methods, which also possess the additional advantage of allowing a direct determination of the temporary hardness. In this case, however, the conception of temporary hardness is narrowed so as to include merely the bicarbonates, the whole of the calcium sulphate being included in the permanent hardness. In the simplest of these methods a measured volume of the water is carefully titrated with decinormal hydrochloric-acid solution, using methyl orange as indicator; alizarin is a still better indicator for the purpose, but titration must then be in boiling solution. The titration depends on the decomposition of the bicarbonate of calcium and magnesium with formation of carbon dioxide and the corresponding chlorides.

Permanent hardness can also be estimated by the alkalimetric method of Wartha and Pfeifer. A measured volume (200 c.c.) of the water is boiled with 50 c.c. of a mixture of decinormal solutions of sodium carbonate and hydroxide in equal amounts; after restoring to the original volume and allowing the solution to settle, the residual alkali is determined by titration with standard acid. As the bicarbonates do not cause any consumption of alkali, there is a direct proportionality between the quantity of alkali which disappears and the total amount of sulphates and chlorides of calcium and magnesium. Sodium carbonate alone does not efficiently precipitate magnesium salts from solution, but precipitation as the hydroxide is complete if excess of sodium hydroxide is present; it is for this reason that a mixture of sodium carbonate and hydroxide is applied.

The last method can also be extended to the measurement of total hardness by first neutralising the bicarbonates as described above for the determination of temporary hardness, and subsequently treating with the mixed alkali solution.

Another satisfactory process for the determination of total hardness, based on a somewhat similar principle, is due to Blacher. The water is first titrated with decinormal hydrochloric acid until it is neutral to methyl orange, as in the method described above for temporary hardness. After the removal of the carbon dioxide by a current of air, the methyl orange is bleached by the addition of a drop of bromine water; a little phenolphthalein and a few drops of alcoholic potassium hydroxide are added, the liquid is just decolorised with decinormal hydrochloric acid and is then titrated with an alcoholic decinormal solution of potassium palmitate until a decided red colour is produced. The quantity of the potassium palmitate solution required is proportional to the total hardness.

A water, the total magnesium and calcium salts of which could be represented by an equivalent of approximately 7 parts (or less) of calcium carbonate per 100,000, would generally be considered soft, whilst it would be described as very hard if the quantity exceeded 45 parts per 100,000.
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