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Potable water treatment
Potable water treatment

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3.4 Chemical characteristics of natural waters

Since water is such a good solvent, it is not surprising to find many different chemical substances present in it. Water, on reaching a river, will contain inorganic and organic compounds which were dissolved as rainwater percolated through the soil and rocks. In addition, some gases will dissolve in rainwater during its passage through the air.

The substances present in water may be conveniently grouped into:

  1. those from dissolved gases such as oxygen and carbon dioxide;

  2. those involved in the bicarbonate–carbonate equilibrium derived from carbon dioxide;

  3. other dissolved substances.

All gases will dissolve to a greater or lesser extent in water. As we have seen, oxygen has a low solubility in water, as do nitrogen, argon and some of the other gases present in the atmosphere.

Carbon dioxide, by contrast, is very soluble in water: 1 m3 of water at 20°C will hold 878 g of pure carbon dioxide. However, carbon dioxide is special for another reason. When carbon dioxide dissolves, it reacts with the water to form bicarbonate and carbonate ions.

The chemical equation describing the process is called the bicarbonate–carbonate equilibrium.

Because all the reactions are reversible, the whole system reaches equilibrium, so that natural waters will contain various proportions of carbon dioxide, bicarbonate and carbonate.

How does the acidity (hydrogen ion concentration) of the water affect the equilibrium in Equation (1)?

For a reversible reaction, a change in the concentration of one of the chemical species in the reaction will produce a corresponding shift in the concentrations of other species in order to 'compensate' for the change. So in Equation (1), if the concentration of hydrogen ions (H+) increases, the reactions move towards the left to compensate. A new equilibrium is reached with higher concentrations of CO2 and bicarbonate, and a lower concentration of carbonate. Conversely, a decrease in hydrogen ions shifts the reactions in Equation (1) to the right. Figure 13 illustrates the relationship between pH and the concentrations of CO2, HCO3 and CO32−. The definition of pH states that low values (lower than a numerical value of 7) of pH correspond to 'acid' (high H+) conditions, whereas high values (greater than 7) of pH correspond to 'alkaline' (low H+) conditions. Figure 13 shows that in high-pH water, most of the carbon dioxide ends up as bicarbonate and carbonate, whereas in low-pH water, the carbon dioxide stays in solution without reacting further.

Figure 13
Figure 13 The bicarbonate–carbonate equilibrium

The bicarbonate–carbonate equilibrium is important in the process of photosynthesis, in which aquatic plants take up the inorganic carbon in carbon dioxide in the presence of sunlight for synthesising new cell material. All plants can use dissolved CO2 for this purpose, but none apparently can use carbonate directly. Blue-green algae can also use bicarbonate for photosynthesis. Thus, low-pH waters, with available carbon dioxide, are more favourable for photosynthesis.

The supply of carbon dioxide to the aquatic environment comes from the atmosphere through diffusion and as the product of aerobic and anaerobic metabolism, and is consumed during photosynthesis.

Apart from substances derived from the atmosphere, there are usually other substances dissolved in natural waters. Salinity is a general term which means the concentration of ionic constituents dissolved in water. These include the carbonates, sulphates and chlorides of sodium, calcium, potassium and magnesium. It may also mean specifically the sodium chloride content which comes from either sewage effluent or sea water intrusion. High chloride contents can also arise in watercourses receiving run-off from salted roads in winter.

The pH of natural water usually varies from approximately 6.0 to 8.0 depending on the types of rocks and substrate surrounding the watercourse, although in some drainage areas this can be as low as 4.0. There has been growing interest and concern regarding the acidity of rainwater in Europe and North America. Acidic precipitation (acid rain) can reach lakes and streams either directly or indirectly after interaction with the vegetation and soils. The magnitude of the effects depends on the buffering capacity of the water.

You might be familiar with hard water, from seeing scale deposited in kettles. As well as scale formation, both temporary and permanent hardness make lathering with ordinary soap difficult. The result is the formation of scum that floats on the surface of washing water. On the benefit side, the dissolved solids or minerals often give hard water a pleasant taste, and they are of nutritional importance to plants and micro-organisms, and may have various medicinal functions for humans.

Hardness in water is mainly due to the presence of ions of the metals calcium (Ca2+), magnesium (Mg2+), and iron (Fe2+). Rivers and lakes fed by water that has run from chalky areas and limestone (CaCO3) contain an abundance of calcium. Calcium and magnesium account for at least 70% of the total cations in water.

If calcium, magnesium and iron are present in water as bicarbonate salts, e.g. Ca(HCO3)2, and the water is boiled or heated above 70°C, carbonate salts of the metals are precipitated. Such water is said to possess temporary or carbonate hardness because the carbonate salts (e.g. calcium carbonate) are largely insoluble, and are thus removed from the water and deposited as scale:

If the scale deposits on heating elements, it shortens their life and makes them less efficient.

When calcium, magnesium and iron are present as chloride or sulphate salts (e.g. CaCl2), the hardness is called permanent or non-carbonate hardness. Although this type of hardness also contributes to scaling, in this case the precipitate is due to the decreased solubility of these metal salts at higher temperatures and not to the formation of new insoluble compounds.

The extent of hard water in the UK tends to follow a north to south-east gradient; the softest water being in Scotland, north England and Wales, and the hardest in East Anglia and south-east England. Also, groundwaters are more likely than surface waters to be hard. Mortality from cardiovascular (CV) disease (heart disease and stroke) tends to follow the same pattern, a higher rate in the north and north-west than in the south and south-east. Several statistical surveys have shown an inverse relationship between CV disease and water hardness. After adjustment for socioeconomic and climatic factors, this relationship is somewhat weakened but remains statistically significant. It can be shown that towns with very soft water (CaCO3 concentration 25 g m−3) have a CV mortality 10–15% higher than in areas with medium-hard water (170 g m−3 CaCO3), while any further increase in hardness above these figures does not additionally lower CV mortality. CV disease may be said to be associated with soft water districts; this association may be influenced by either water hardness itself or by some factor closely associated with it. As a consequence, softening of water for domestic use is rarely carried out except for very hard sources. In homes with their own water-softening system, a tap is usually installed allowing hard water to be drawn off for drinking purposes.