I organic air pollutants I 1 Volatile Organic Compounds (vocs)

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I.2.5.1. Acidity

The acidity of a water sample is by definition, its capacity to react with a strong base until a certain value of pH. The acidity is expressed as the concentration in "milli-equivalent by gram" of hydrogen ions or like the equivalent amount of calcium carbonate that is required to neutralize the acidity.

The determination of the acidity intends "to quantify the acid concentration in a water sample or in a liquid residue". This data is important because the acid substances present in the water are responsible for the corrosivity increase and they interfere with the reaction capacity of many substances by interfering some chemical processes to the interior of the water systems.

Thus, the quantification of acid substances is useful and necessary, as much as it allows its later neutralization and in general, the adjustment of the water for a certain aim or application.

The acidity in the water can be associate to the weak acid presence such as carbon dioxide, to the strong acid presence like the sulfuric, hydrochloric and nitric acids and to the presence of strong salts from weak bases, such as those of ammonium, Fe3+, Al3+, etc. Although the acidity from CO2 has little importance from the point of view of the drinking degree, from the industrial point of view it is very important due to the corrosive power of present acid substances in the water.

The most frequent form to measure the acidity, it is by titration with a strong base, 0.020 mol L-1 Na(OH), generally using as indicator the methyl orange, (turn between 3.1 and 4.4) or the bromocresol green, (color change between 3.0 and 4.6), for the first point and the phenolphthalein, (change over 8.0 - 10.0 for the second point).

Because the pH of the water saturated with CO2 and one atmosphere of pressure is of the order of 4.5, any water sample whose pH is lower than 4, contains some additional acid substances different from carbon dioxide. The concentration of that additional acid substance is known and measured like “mineral acidity”, by titration with sodium hydroxide from the original value of pH to pH 4.5.

Nevertheless, in a natural water system whose pH is over the range 4 - 6, it can be assumed that the acidity must be almost exclusively due to the concentration of CO2. This acidity, known as “carbonic acidity”, is titrated up to pH 8.2 with NaOH by means phenolphthalein.

The addition of these two measurements, expressed in terms of milligrams of calcium carbonate L-1 of solution, is known as “total acidity”. In general, it is not usual for natural water to present mineral and carbonic acidity simultaneously; however, it is frequent in industrial residual waters. The carbonic acidity is by far, the most frequent form of acidity in water.

I.2.5.2. Alkalinity

The alkalinity of a water sample is the capacity to react or to neutralize hydrogen ions, (H+), up to a pH value of 4.5. The alkalinity is caused mainly by bicarbonates, carbonates and hydroxides in solution and in smaller degree by borates, phosphates and silicates. In a strict sense the main species causing alkalinity and their association with the source, are the following:

Hydroxides, natural, residual and industrial waters

Bicarbonates and carbonates, natural and residual

CO2, underground (deep) and residual water

Silicates, SiO32- and HSiO3, underground waters

Borates, BO33-, HBO32-, H2BO3 underground and agricultural residual waters

Phosphates, PO43-, HPO42-, H2PO4 agricultural and industrial waters

In spite of the reported in most of natural water systems, alkalinity is associated to the carbonate system, which means the sum of carbonates and bicarbonates. Due to that, the alkalinity usually is taken as an indicative of the concentration of these substances, (Figure I.2.3.), however this does not mean for all the cases, that the alkalinity is exclusively only due to bicarbonates and carbonates.

The relatively old underground waters running through sandy layers, constitutes a good exception, where the alkalinity is also related to dissolved silicates.

The alkalinity in most of the natural water systems has its origin in the system carbonate, because carbon dioxide and the bicarbonates come from the metabolism of the live organisms, aerobic or anaerobic, which can survive in any water, with organic matter and minimum conditions of survival. Since this possibility is frequent in most of the water that surround us, the "system carbonate" is present in all of them:

M. O. + O2 / Fe3+, NO3, etc. ® CO2 + H2O

CO2 + H2O ® H2CO3 ≈ 100 mg of CO2 / liters of water with CN-

H2CO3 ® H+ + HCO3

HCO3 ® H+ + CO3-2

One of the main consequences from the acid-basic characteristics of carbonate system is the slight "buffering capacity" of the waters. Thus, the concentration of the system carbonate in the water determines its buffering capacity, and the proportion of these ions CO2, HCO3- and CO3 2- determines the pH.


Figure I.2.3. System Carbonate in waters

The alkalinity in the water expresses the equivalent hydroxyl ion concentration, in mg L-1 or the equivalent amount of CaCO3, in mg L-1. The alkalinity, understood as the alkaline-earth metal concentration, it important in the determination of the quality of the water for irrigation and is in addition, an important factor in the interpretation and the control of the processes of residual water purification.

The alkalinity is titrated with 0.02 N HCl by using indicator phenolphthalein, when the original pH of the sample is over 8.3, or methyl orange in the opposite case. The former case is called P Alkalinity, (to phenolphthalein) and in the second of M Alkalinity (to methyl orange). Figure I.2.4 facilitates the selection of an indicator for the measurements of alkalinity or acidity.

Since the alkalinity is a direct function "of the system carbonate" in the sample, the alkalinity values obtained "in situ" usually differ from the ones obtained in the laboratory on transported samples, because these can absorb or release CO2 before measurements in the laboratory.

These differences are relevant especially in underground water samples because the prevailing pressures to these depths determine the balances of dissolution of gases.

To obtain a more exact value of the alkalinity, it is important to measure the alkalinity "in situ" or in the mouth of the well for deep waters. To obtain a better ionic balance in the joint analytical results of these samples, it could be preferable to let the sample balance to the atmospheric pressure before measurements.

Figure I.2.4. Common indicators and change intervals

I.2.5.3. Salinity

It is the concentration of the soluble mineral salts in the water, mainly of metals like sodium, magnesium and calcium. The mineral salts are good conductors; organic and colloidal matters present low conductivity. Consequently, in the case of residual waters, this measurement does not give an immediate idea of the total contents in the water sample.

The concentration is usually expressed in parts by million (ppm) and according to the present salt concentration there are different water types: Fresh water, less than 1 000 ppm; Slightly salty water, from 1 000 ppm to 3 000 ppm; Moderately salty water, from 3 000 ppm to 10 000 ppm; Highly salty water, from 10 000 ppm to 35 000 ppm (ocean waters are about 35 000 ppm, 3.5%).

Other usual way is to be expressed as psu (practical salinity units); the ocean waters are about 35 psu. The relative amounts of solved salts in sea water have approximately the following figures (in %): 55.3 Cl (I); 30.8 Na (I); 3.7 Mg (II); 2.6 SO42-; 1.2 Ca (II) and 1.1 K (I).

For seawater, the salinity is defined conventionally as the weight in grams of dried solid compounds until constant weight to 480º C, obtained from 1 kg of water of sea. It is assumed that the organic matter has oxidized, the bromine and the iodine have been replaced by their equivalent in chlorine and carbonates turned oxides.

In fresh waters the most abundant mineral salts are the carbonates, the sulfates and the chlorides. Cations of greater importance are: Ca2+, 64%; Mg2+, 17%, Na+, 16%; and K+, 3%. The calcium plays a fundamental role in fresh waters since it determines two different water types: a) hard waters, and b) soft waters, the last type is when the concentration is inferior to 9 mg by litter. Many mollusks, crustaceans and other invertebrates, have calcium necessity to form their shells and therefore can be a lim.

The salinity measurement is very important in irrigation waters. Generally, the content of salts is reduced and does not interfere in the good development of the plants, but, in certain places and circumstances, water can contain high concentrations of salts. For this reason, it is highly recommendable to make periodic water analysis in irrigated land operations especially if settled irrigations by dripping are present, where the possibilities of an accumulation of salts in the ground are greater.

The progressive increase of the concentration of soluble salts, due to the continued irrigation, brings an increase of the osmotic pressure of the dissolution in the ground. The greater the concentration of salts is, the greater the osmotic pressure that the roots of the plants will have to surpass before being able to absorb water.

When saline waters are used for irrigation some recommendations are required. The saline water preferably will be used in light and permeable grounds and for tolerant cultures to salinity. And, it is necessary to facilitate the removing of excess of salts by means of a suitable system of drainage. Frequent rain over permeable grounds, makes unnecessary the above-mentioned recommendations, since the rain would wash the grounds. The phreatic level is deep enough to avoid problems by excess of salinity from ground water. The irrigation by aspersion is not recommendable for waters with the conductivity greater than 2 mmhos/cm, since these can damage the installation and produce burns in the leaves of the plants.

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