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CES Technical Report 133,   August 2013

8.1 WATER TEMPERATURE:  The water temperature parameter is important for its effects on the chemistry, and biological reactions in the organisms in water. A rise in temperature of the water leads to the speeding up of the chemical reactions in water, reduces the solubility of gases and amplifies the tastes and odours. At higher temperatures with less dissolved gases, the water becomes tasteless. At elevated temperatures metabolic activity of organism’s increases, requiring more oxygen but at the same time the solubility of oxygen decreases. The disease resistance in the fishes also decreases with rise in temperature. The temperature of drinking water has an influence on its taste (Trivedy and Goel., 1984)

Procedure

Clean the pH-TDS-EC probe with distilled water and fill the container with water sample. Then check the water temperature of the sample by dipping the pH-TDS-EC probe into sample and express in degrees Celsius (ºC).

8.2 CONDUCTIVITY: Conductivity is a numerical expression of the ability of an aqueous solution to carry an electrical current. This ability depends on the presence of ions, their total concentration, mobility, valence, relative concentration and on the temperature of measurement (APHA, 1985). Solution of most inorganic acids, bases and salts are relatively good conductors. Conversely molecules of organic compounds that do not dissociate in aqueous solution conduct a current very poorly, if at all. Conductivity in water is affected by the presence of inorganic dissolved solids such as chloride, nitrate, sulfate, and phosphate anions (ions that carry a negative charge) or sodium, magnesium, calcium, iron, and aluminum cations (ions that carry a positive charge). Organic compounds like oil, phenol, alcohol, and sugar do not conduct electrical current very well and therefore have a low conductivity when in water. Conductivity is also affected by temperature: the warmer the water, the higher the conductivity. For this reason, conductivity is reported as conductivity at 25° C.

Normally the physical measurement of conductivity is done through determination of resistance, measured in ohms or megohms. The resistance of a conductor is inversely proportional to its length. The magnitude of the resistance measured in an aqueous solution therefore depends on the characteristics of the conductivity cell used; it is not meaningful without this knowledge.

Using this method, a low-voltage alternating current is applied across the electrodes. The resistance of the electrodes is measured and is converted to conductivity according to the following,

K = L/AR

Where K is the conductivity, L is the distance between the electrodes (cm), A the surface area of the electrode (cm²), and R the resistance [ohm = Siemens (S)^-1] Siemens is the S.I unit of the electrical conductivity (Roger Reeve, 2002).

The cell is calibrated by using solutions of known conductivity. Conductivity is highly temperature dependent and so care has to be taken that calibration solutions and the unknown sample are at the same temperature. A standard temperature of 25° C is often used. The relationship between conductivity and total salt content is not simple. All ions having the same charge have approximately the same conductivity, but unfortunately most environmental waters contain ions with different charges in varying concentrations.

The substance dissociates positive and negative ions and imparts conductivity in a solution. Those with poor solubility in water are called weak electrolytes and those with high solubility, the strong electrolytes. Thus higher the concentration of electrolytes in water the more is its electrical conductivity. i.e. lesser the resistance. The conductance of distilled water ranges from 1 to 5 μmho (μS).
The amount of current that can flow through a solution is proportional to the concentration of dissolved ionic species in the solution. The rough guide to the conversion between μS/cm to mg/l for natural waters is the conductivity in μS/cm is about 110-115 % of the dissolved solids in mg/l.

Conductivity in streams and rivers is affected primarily by the geology of the area through which the water flows (which is a natural process). Streams that run through areas with granite bedrock tend to have lower conductivity because granite is composed of more inert materials that do not ionize (dissolve into ionic components) when washed into the water. On the other hand, streams that run through areas with clay soils tend to have higher conductivity because of the presence of materials that ionize when washed into the water. Ground water inflows can have the same effects depending on the bedrock they flow through. (Chapman, 1996)

Procedure
Clean the pH-TDS-EC probe with distilled water and fill the container with water sample. Then the electrical conductivity of the sample is determined by dipping the conductivity probe into sample and express in μS/cm.

8.3 SOLIDS: The general term solids refer to matter is suspended (insoluble solids) or dissolved (soluble solids) in water. Solids can affect the water quality in several ways. Drinking water with high dissolved solids may not taste good and may have a laxative effect. Boiler water with high dissolved solids requires pre-treatment to prevent scale formation. Water high in suspended solids may harm aquatic life by causing abrasion damage, clogging of fish gills, harming spawning beds, and reducing photosynthesis by blocking sunlight penetration, among other consequences.

8.4 Total Dissolved Solids (Non-Filterable Solids): The substances that would pass through the filter paper but will remain as residue when the water evaporates which includes dissolved minerals and salts, humic, tannin and pyrogens. A constant level of minerals in the water is necessary for aquatic life. Changes in the amounts of dissolved solids can be harmful because the density of total dissolved solids determines the flow of water in and out of an organism’s cells. Many of these dissolved solids contain chemicals, such as nitrogen, phosphorous, and sulfur, which are the building blocks of molecules for life. Concentration of total dissolved solids that are too high or too low may limit the growth and may lead to the death of many aquatic organisms. High concentrations of total dissolved solids may reduce water clarity, which contributes to a decrease in photosynthesis and lead to an increase in water temperature. Many aquatic organisms cannot survive in high temperatures (Ramteke and Moghe, 1988).

The conductivity in water is affected by the presence of inorganic dissolved solids such as chloride, nitrate, sulfate, and phosphate anions or sodium, magnesium, calcium, iron, and aluminum cations. These particles contribute the dissolved solids of the water.

Procedure

Clean the pH-TDS-EC probe with distilled water and fill the container with water sample. Then the total dissolved solids is determined by dipping the pH-TDS-EC probe into sample and express in parts per million (ppm).

8.5 pH: pH is one of the most important parameter in water chemistry and it indicates the acid/ alkaline or neutral status. pH of 7 is considered to be neutral. When the pH is less than 7, it is acidic; a pH greater than 7 is basic. A pH value between 7.0 and 8.0 are optimal for supporting diverse aquatic ecosystem. A pH range between 6.5 and 8.5 is generally suitable. Acidity or alkalinity are the acid and base neutralizing capacities of a water and usually expressed as milligrams CaCO3/L. pH is given by  –log[H+], it is the intensity factor of acidity (Sorenson, 1909). At a given temperature the intensity of the acidic or basic character of a solution is indicated by pH or H+ activity.

pH = -log[H]+  As H+ increases, pH decreases. Since pH is on a log scale based on 10, the pH changes by 1 for every power of 10 changes in [H] (APHA, 1985). Pure water is very slightly ionized and at equilibrium the ion product is

At a given temperature, pH indicates the intensity of the acidic or basic character of a solution and is controlled by the dissolved chemical compounds and biochemical processes in the solution. In unpolluted waters, pH is principally controlled by the balance between the CO2, carbonates and bicarbonate ions as well as the other natural compounds such as humic and fulvic acids (Chapman, 1996). The level of acidity of the water is important to the plant and animal life there. Most animals are adapted to living in neutral conditions. Changes in pH endanger the lives the organisms in the water.

The level of acidity can be changed by human’s actions. Acid rain, a result of air pollution and matter emitted from tailpipes and smokestacks affect the pH. When these things combine with water in the atmosphere, they form sulfuric and nitric acids, then fall to the earth as acid rain, snow, hail, and fog. This precipitation mixes with water already on the earth, in creeks, rivers and wetlands. Other pollutants carried by runoff from the land, also change the acidity of the water (Chapman, 1996).

Procedure

Clean the pH-TDS-EC probe with distilled water and fill the container with water sample. Then pH of the sample is determined by dipping the pH probe into sample.

8.6 SALINITY: Salinity is defined as the total solids in water after all carbonates have been converted to oxides, all bromide and iodide have been replaced by chloride, and all organic matter has been oxidized. It is numerically smaller than the total dissolved solids (APHA, 1985).

Procedure

Clean the pH-TDS-EC probe with distilled water and fill the container with water sample. Then determine the salinity of the sample by dipping the pH-TDS-EC probe into sample and express in parts per million (ppm).

8.7 HARDNESS: The hardness of natural water depends on the presence of calcium and magnesium salts. The total content of this hardness is total hardness or general hardness. The total hardness is further classified in to carbonate and non-carbonate hardness. Carbonate hardness determined by the concentration of the calcium and magnesium hydro carbonates and non-carbonate hardness is determined by calcium and magnesium salts of strong acids. Hydro carbonates are transformed during the boiling of water into carbonates, which usually precipitates, so the carbonates hardness is also called as temporary hardness. On the other hand the hardness remaining in the water after boiling is called constant hardness or permanent hardness (Chapman, 1996).

In total hardness the calcium hardness is usually prevalent (up to 70%) although in some cases magnesium hardness can reach up to 50-60%. Water hardness is the ability of water to precipitate soap. It is roughly measured by the amount of soap required for adequate lathering and also the rate of scale formation in the hot water boilers.

Hardness has some similarities with alkalinity. Like alkalinity, hardness is also not attributed to a single constituent, so some convention must be adopted to express hardness quantitatively as a concentration. As with alkalinity hardness is usually expressed as an equivalent concentration of CaCO3. Hardness is the property of cations (Ca2+ and Mg2+) and alkalinity is the property of anions (HCO3- AND CO32-) (Weiner, 2000)

Hardness is sometimes useful as an indicator proportionate the total dissolved solids, because Ca2+. Mg2+ and HCO3- often represent the larger portions of TDS. There is no report relating hardness with the human health except some inverse effect with the cardiovascular disorders. The high level of the hardness may inhibit the growth of fish, but low level of hardness increases the sensitivity of the fish towards the toxic metals, so the higher hardness is beneficial by reducing the toxicity to fish. The soft water is readily dissolvable than the hard water, so it brings metal ions from the plumbing network. Lime stones, which are percolated by rainwater, are the natural source and artificial sources such as mine discharge, industrial and domestic effluents also contribute hardness.

EDTA titrimetry for hardness

Ethylene diamine tetra acetic acid (EDTA) forms a chelated complex when it is added to a solution containing metal cations. The dye Erichrome Black-T forms a complex with cations in the solution at pH of 10.0 ± 1.0 and turns to wine red colour. When the EDTA is added as a titrant it gets complexed with the existing cations. When all the magnesium and calcium gets complexed, the indicator comes back to its original blue colour.

METAL + INDICATOR          METAL-INDICATOR COMPLEX (Wine red colour)

METAL - INDICATOR COMPLEX + EDTA               METAL-EDTA COMPLEX    +   INDICATOR (Blue colour)

Reagents

1. Standard EDTA indicator: (0. 01 M) Weigh 3.723 g of EDTA and dissolve in 1000 ml of distilled water and standardise against standard calcium solution.
2. EBT Indicator: The powdery form of the Erichrome Black-T can be used for this purpose.

Procedure

Select the sample volume requires less than 15 ml of EDTA titrant in a conical flask and complete titrant within 5 minutes after addition of the buffer. Take 25 ml of the sample (dilute it with distilled water if required) and add 1-2 ml of the buffer solution and 1-2 g of EBT indicator. Then titrate against the EDTA, with continuous stirring; until the last reddish tinge disappears. At the end point the solution gives blue colour.

Calculation

                                                             A*B*1000
Hardness (EDTA) as mg CaCO₃/L =    * Dilution factor (If diluted)
                                                           ml of sample

Where A= ml titrant required and B= mg CaCO₃ equivalent to 1 ml EDTA titrant

Degree of hardness	Mg.CaCO3/L	Effects
Soft	< 75	May increase toxicity of dissolved metals. No scale deposition
Moderately hard	75-120	Not objectionable for most purposes. Requires more soap Slightly scale formation
Hard	120-200	Considerable scale formation and staining
Very hard	>200	Requires softening for household and commercial use.

8.8 CALCIUM HARDNESS: The presence of calcium (fifth most abundant) in water results from passage through or over deposits of limestone, dolomite, gypsum and such other calcium bearing rocks. Calcium is present in all waters as Ca2+ and is readily dissolved from rocks rich in calcium minerals, particularly as carbonates and sulphates. The salts of calcium, together with those of magnesium are responsible for hardness of  water (Chapman, 1996). Calcium is an important element for all organisms and is incorporated in to the shells of many aquatic invertebrates, as well as the bones of vertebrates. Calcium concentrations in natural waters are typically <15 mg/l. For water associated with carbonate rich rocks concentrations may reach 30-100 mg/l. Salt waters have concentrations of several hundred milligrams per litre or more.

Ethylene diamine tetra acetic acid (EDTA) Titrimetric Method

The simplicity and rapidity of the EDTA titration procedure make it the method of choice for general use. When EDTA is added to water containing both calcium and magnesium, it combines first with the calcium, because during the increase of pH magnesium is largely precipitated as the hydroxide and the murexide indicator combines only with calcium.

Reagents

1. Sodium hydroxide: (NaOH) 1Normality.
2. Indicator: Murexide  (Ammonium purpurate)
3. Standard EDTA indicator: (0. 01 M) Weigh 3.723 g of EDTA and dissolve in 1000 ml of distilled water and standardise against standard calcium solution.

Procedure

Sample preparation: Start titration as soon as the addition of the alkali and indicator, because high pH is required for this reaction .Use 50 ml sample or a smaller portion diluted to 50 ml, so that the calcium content is about 5-10 mg. Analyse hard waters with alkalinity higher than 300 mg CaCO₃/l by taking the smaller portion and diluting to 50ml, or by neutralizing the alkalinity with acid, boiling 1 min and cooling before beginning the titration.

Titration: Add 2 ml NaOH solution per volume sufficient to produce the pH 12-13. Stir well and add a pinch of indicator. Add EDTA titrant slowly with continuous stirring to the proper end point [pink to purple]. Run the blank along with the sample.

Calculation

                                                          A*B*1000
Calcium hardness as CaCO₃ =     * Dilution factor (If diluted)
                                                    ml of sample taken

Where A = ml titrant for sample and B = mg Caco₃ equivalent to 1.0 ml EDTA titrant.

8.9 MAGNESIUM HARDNESS: Magnesium arises principally from the weathering of rocks containing Ferro magnesium minerals and from the carbonate minerals, ranking eighth in abundance among the elements. It is found in all natural waters and its source lies in rocks, generally present in lower concentration than calcium. It is also an important element contributing to hardness and an essential element for photosynthetic pigment chlorophyll. Its concentration greater than 125 mg/L can influence cathartic and diuretic actions. Magnesium hardness can be calculated from the estimated total hardness and calcium hardness.

Calculation

Magnesium = (A - B) as mg/L

where, A = Total hardness mg\L (as CaCO₃) and  B = Calcium hardness mg\L (as CaCO₃)

8.10 SODIUM AND POTASSIUM:  Sodium ranks sixth among the elements in order of abundance and is present in most natural waters. The high solubility of sodium confirms its presence in the all the natural water. Sodium available in Na⁺ (ionic form) in plant and animal. The increased concentration of the sodium in surface water is due to industrial effluent, agricultural activities and seawater intrusion in coastal area. Sodium is commonly measured where the water is used for the drinking and agricultural purposes. The elevated sodium level in certain soil types result in the degradation of soil fertility and it also results in restriction of water movement, which leads to affects the plant growth. The high concentration of sodium in water bodies limits the biological diversity due to the osmotic pressure. Soil permeability can be harmed by high calcium ratio. (Weiner, 2000)

Potassium ranks seventh among the elements in order of abundance. Potassium (as K⁺) is found in low concentrations in natural water, because the rocks containing potassium is resistance to weathering. But potassium from industrial waste and agricultural run off contributes its level in natural freshwater bodies. Potassium usually found in the ionic form and it is easily soluble. It is readily incorporated into mineral structure and accumulated by aquatic biota, as it is an essential nutritional element. The normal concentration of potassium in natural waters is usually less than 10 mg/L, whereas concentrations as high as 100-25000 mg/L can occur in hot springs and brines (Weiner, 2000).

The estimation of sodium and potassium on the emission spectroscopy deals with the excitation of electrons from ground state to higher energy state and come back to original position with emission of light. The trace amount of sodium and potassium can be determined by flame photometry method. The sample is sprayed in to a gas flame; the excitation takes place in controlled condition. The desired spectral line is isolated by appropriate silt arrangement in light dispersing device. The intensity of the light is measured by phototube potentiometer. The intensity of light at appropriate wavelength is approximately proportional to the concentration of the element.

INSTRUMENT: Systronics Flame Photometer 128 with a compressor and a gas supply.

Reagents

1. Stock sodium solution: Dissolve 2.542g of Sodium chloride of dried (at 140o C) was in 1000ml-distilled water (1ml = 1mg of sodium).
2. Intermediate sodium solution: Dilute 10 ml of the stock sodium solution to 100 ml of water. (1 ml = 100µg Na)
3. Stock potassium solution: Dissolve1.907g of dried Potassium chloride, in 1000ml of distilled water. (1ml = 1mg of potassium)
4. Intermediate potassium solution: Dilute 10 ml of the stock potassium solution to 100 ml of water. (1 ml = 100µg K)
5. Standard sodium and potassium solution: Prepare the standard solution as per the required concentration, from intermediate solution.

Procedure

Follow the instructions given by the manufacturers of flame photometry. Start the electrical supply and compressor, stabilise the air supply to appropriate level. Switch on the gas fuel and adjust the cone-shaped blue coloured flame by aspirating distilled water. Calibrate the instrument with standard solutions and follow it with sample analysis.

Calculation

            Mg/L of sodium = Test absorbance * Slope * D

                                               Sum of concentration of standards               
                   Where, Slope =
                                               Sum of absorbance of standards

            Mg/L of Potassium = Test absorbance * Slope * D

                                               Sum of concentration of Standards
                   Where, Slope =
                                               Sum of absorbance of Standards

                                   ml sample + ml distilled water
D = Dilution ratio =   
                                                ml sample

8.11 ALKALINITY: In natural water, that are not highly polluted alkalinity is more commonly found than acidity. Alkalinity is the good indicator of presence of dissolved inorganic carbon (Bicarbonates and carbonate anions). Some of the minor contributions to alkalinity are come from ammonia, phosphates, borates, silicates and other basic substances.

Alkalinity is beneficial to aquatic animals and plants, because it buffers the both natural and human induced pH changes. Water with high alkalinity generally has a high concentration of dissolved inorganic carbon (in the form of HCO₃- and CO₃²-), which can be converted to biomass by photosynthesis (Kegley and Andrews, 1998)

Titrating a basic water sample with acid to pH 8.3 measures phenolphthalein alkalinity. Phenolphthalein alkalinity primarily measures the amount of carbonate ion (CO32-) present in the sample. Titrating with acid to pH 3.7 measures methyl orange alkalinity or total alkalinity. Total alkalinity measures the neutralizing effects of all the bases present.

Reagents

1. Standard sulphuric acid (0.02N):  Dilute 200 ml of 0.1N H₂SO₄ with distilled water to 1000ml to obtain standard 0.02N H₂SO₄.
2. Phenolphthalein indicator: Dissolve 0.5g phenolphthalein in 1:1 95 % ethanol and distilled water. Add 0.05 N CO₂ free NaOH solution drop wise until faint pink colour appears.
3. Methyl orange indicator: Dissolve 0.5 g of methyl orange in 1000ml CO₂ free distilled water.

Procedure

Measure a suitable volume of sample in 250 ml conical flask. Add 2-3 drops of phenolphthalein indicator. If the pink colour develops titrate against 0.02 N H₂SO₄, till the colour disappears, which is the characteristic of pH 8.3.Note down the volume of H2SO4 consumed. Add 2-3 drops of methyl orange and continue titration with H₂SO₄ till the yellow colour changes to orange, which is the characteristic of pH 4.5. Note down the additional amount of H2SO4 required. In case pink colour does not appear after addition of phenolphthalein continue with addition of methyl orange.

Calculation

Calculate total phenolphthalein and methyl orange alkalinity as follows and express in mg/L as CaCO₃,
                                                                 A*1000
      P Alkalinity m mg/L as CaCO3 =      
                                                               ml sample

                                                               B*1000
     T Alkalinity m mg/L as CaCO3 =     
                                                             ml sample

Where, A= ml of H₂SO₄ required to raise pH up to 8.3, B= ml of H₂SO₄ required to raise pH up to 4.5

8.12 CHLORIDES: Chloride (Cl-) enters surface waters by the atmospheric deposition of oceanic aerosols with the weathering of some sedimentary rocks (mostly rock salt deposits), industrial effluents and agricultural run-off. High concentrations of chloride can make waters unpalatable and therefore unfit for drinking or livestock watering. Chloride is abundant anions found in the wastewater and is a good indicator of pollution sources. Chloride gives water a salty taste detectable at a level of 250 ppm with sodium as cation, but with Ca and Mg cations, the salty taste is not detectable until the chloride concentration reaches up to 1000 ppm (Weiner, 2000).

All chlorides salts are very soluble expect for chloride salts of lead (PbCl₂), silver (AgCl₂) and mercuric (Hg₂Cl₂, HgCl₂). Chloride is not absorbed by soils and little part moves with water and there is no retardation. Consequently, it eventually moves to closed basins (as Great salt lake in Utah) or to the ocean. (Kegley and Andrews, 1998)

Estimation of chloride by argentometric method        

In a neutral or slightly alkaline solution, potassium chromate can indicate the endpoint of the silver nitrate titration of chloride. Silver chloride is quantitatively precipitated before red silver chromate is formed.

            1) Ag++Cl-            AgCl (White precipitate)
            2) 2Ag++CrO-4                AgCrO4  (Red precipitate)

Reagents

1. Potassium chromate indicator: Add 50g K₂CrO₄ in little amount of water and dilute to 1000ml.
2. Standard Silver nitrate (0.0141N): Dilute 70.5 mL of 0.1N AgNO₃ in distilled water to 500ml. (1ml=0.5mg Cl=500µg Cl)
3. Standard Sodium Chloride (0.0141N): Dissolve 824.0mg NaCl (dried at 140oC) and dilute to 1000ml. (1ml= 0.5mg Cl=500µg Cl)

Procedure

Use a 100ml of sample or a suitable portion diluted to 50ml. Directly titrate samples in the pH range 7-10.Adjust sample pH to 7 to 10 with H2SO4 or NaOH if it is not in this range. Add 10 ml of K2CrO4 indicator solution. Titrate with standard AgNO3 titrant to a pinkish yellow colour end point. Standardize AgNO3 against standard NaCl. For better accuracy titrate distilled water (50ml) in the same way to establish reagent blank.

Calculation

                                    (A-B)*N*35.45*1000
Chlorides mg/L =     
                                          ml sample

Where, A = ml Ag NO3 required for sample. B = ml Ag NO3 required for blank. N = Normality of AgNO3 used.

8.13 NITRATES: Nitrate ion (NO3¯) is the common form of combined nitrogen found in natural waters. It may be bio chemically reduced to nitrite (NO2¯) by denitrification processes, usually under anaerobic conditions. The nitrite ion is rapidly oxidized to nitrate. Natural sources of nitrate to surface water include igneous rocks, land drainage and plant growth and decay. The natural concentration of nitrate is 0.1mg/L; it may be enhanced by sewage, industrial effluents and nitrate fertilizers (Chapman, 1996)

Ammonia and other nitrogenous material in natural waters tend to be oxidized by aerobic bacteria first to nitrite and then nitrates. So all organic compounds containing nitrogen are potential source of nitrates. In oxygenated to nitrite (NO2¯) is rapidly oxidized to nitrate (NO3¯) so normally the amount of nitrite is very low. High concentrations (>1-2 mg/L) of nitrate or nitrite in surface or ground water generally indicate agricultural contamination from fertilizers and manure.

NITRIFICATION

                                 (O)                                        (R)
Ammonia                  Nitrite                Nitrates
                               ADB                                        FAB

ADB = Aerobic Denitrifying Bacteria.  Ex: Nitrosomes
FAB = Facultative Anaerobic Denitrifying Bacteria. Ex: Pseudomonas

Estimation of nitrates by electrode screening method

Principle: The NO3¯ ions electrode is a selective sensor that develops a potential across a thin, porous, inert membrane that holds in a place a water-immiscible liquid ion exchanger. The electrode responds only to NO3¯ ion activity between about 10¯5 and 10¯1 M (0.14 to 1400 mg NO3¯ -N/L). The lower limit of detection is determined by the small but finite solubility of the liquid ion exchanger.

Apparatus: Nitrate ion electrode

Reagents

1. Nitrate free water: Use distilled or deionized water of highest purity to prepare all solutions and dilutions.
2. Stock nitrate solution: Dry potassium nitrate in an oven at 105°C for 24 h. dissolve 0.7218 g in water and dilute to 1000 mL. Preserve with 2 mL CHCl3/L; 1.00 mL = 100 µg NO3¯ -N.
3. Standard nitrate solution: Dilute 1.0, 10 and 50 mL stock nitrate solution to 100 mL with water to obtain standard solutions of 1.0, 10 and 50 mg NO3¯ -N/L, respectively.
4. Buffer solution: Dissolve 6.66 g Al2(SO4) 3.18H2O, 3.12 g Ag2SO4,  1.24 g H3BO3, and 1.94 g sulfamic acid (H2NSO3H), in about 400 mL water. Adjust to pH 3.0 by slowly adding 0.10N NaOH. Dilute to 1000 mL.
5. Sodiun hydroxide, NaOH, 0.1N.

Procedure

Prepare calibration curve by transfer 10 mL of 1 mg NO3¯ -N standard to a 50 mL beaker, add 10 mL 10 mL buffer solution, and stir for a constant time (2 or 3 min.). Immerse electrode tips and record millivolt reading after 1 min. Remove electrodes, rinse, and blot dry. Repeat for 10 mg NO3¯ -N/L and 50 mg NO3¯ -N/L standards and plot a graph.

Measurement of sample: Transfer 10 mL sample to a 50 mL beaker, add 10 mL buffer solution, and stir for a constant time (2 or 3 min.). Immerse electrode tips in sample and record potential reading after 1 min. Read concentration from calibration curve.

8.14 PHOSPHATES: Phosphates are an essential nutrient for living organisms and exist in water bodies on both forms as dissolved and particulate species. Phosphorus concentration is limiting factor for algae growth. Artificial increase in phosphorous concentration in aquatic system results in eutropication. In natural water and wastewater phosphorous occurs as orthophosphates, polyphosphates and organically bound phosphates (Chapman, 1996)

Orthophosphates applied to agricultural land as a fertilizer, it find its way to surface water by runoff. The condensed phosphate like pyrophosphates and polyphosphates are comes to environment from the heavy-duty detergents (Sodium tri polyphosphate Na5P3O10). Biological processes are the only source for organic phosphates, such as weathering of rock and decomposition of organic matter.

Estimation of phosphorous by stannous chloride Method

In the acidic condition, orthophosphates reacts with ammonium molybdate to form molybdophosphoric acid. It is further reduced to intensely coloured molybdenum blue colour by adding the reducing agent stannous chloride. The intensity of the coloured complex is measured at 690nm, which is directly proportional to the concentration of phosphates present in the sample.

PO4+12(NH4)2MoO4+24H+         (NH4)PO4.12MoO3+21NH4++12H2O

(NH4)3PO4.12 MoO3+Sn2+                Molybdenum blue +Sn4+

Reagents

1. Standard phosphate solution: Dissolve 219.5 mg of anhydrous KH2PO4 in 1000ml of water (1 ml = 0.5 mg PO4).
2. Ammonium molybdate solution:
1. Dissolve 25 g ammonium molybdate in 175 ml distilled water.
2. Add cautiously 280 ml concentrated sulphuric acid to 400ml water and cool.
Mix two solutions (a) and (b) and make up to 1000ml.
3. Stannous Chloride: Dissolve 2.5 g SnCl2. 2 H2O in the 100ml glycerol. Heat in water bath to ensure complete dissolution. Prepare stannous chloride freshly.

Procedure

Take the appropriate aliquot of sample (as per the expected concentration of phosphates) in the Nessler tube. Add 2ml of ammonium molybdate and mix well. Add 0.5 ml of stannous chloride and make up to the mark with distilled water. Prepare the blank using the distilled water instead of sample. Measure the intensity of light path at 690 nm.

Calculation

            Mg/L of phosphates = Test absorbance * Slope * Dilution factor

                                    Sum of concentration of Standards
 Where, Slope =        
                                    Sum of absorbance of Standards

                                              ml sample + ml distilled water
 D = Dilution ratio =           
                                                        ml sample

8.15 SULPHATES: Sulphate is naturally present in surface water as SO₄2-. It arises from atmospheric deposition of oceanic aerosols and leaching of sulphur compounds, either sulphate minerals such as gypsum or sulphide minerals such as pyrite from sedimentary rocks. Under anaerobic condition, sulphates serve as an oxygen source for bacteria that reduce dissolved sulphate to sulphide, which is then volatilised to the atmosphere as H2S. This process is common in anaerobic regions of wetlands and lakes fed by surface and ground waters with high sulphate levels (Chapman, 1996).

Industrial discharge and air emission of sulphur-di-oxide and sulphur-tri-oxide are the sources of sulphate in surface water. Mine drainage contributes large amount of sulphate. Sulphate concentration normally varies between 10-80 mg/L in most of the surface waters, although they may reach up to thousands of milligram per litre near industrial effluent. High sulphate content is present in well water and surface water in arid zone, because of sulphate minerals.

Estimation of sulphates by barium chloride precipitation method

Sulphate ions in the sample are precipitated as barium sulphate BaSO4 in acidic media (HCl) with barium chloride. The absorption of light by this precipitated suspension is measured by spectrophotometer at 420nm.

Ba⁺⁺ + SO₄²-           BaSO₄ (White precipitate)

Reagents

1. Conditioning reagent: Mix 50 ml glycerol with a solution containing 30ml concentrated hydrochloric acid, 300ml of distilled water, 100ml 95% ethyl alcohol or isopropyl alcohol and 75 g Nacl.
2. Barium chloride: Barium chloride crystals of 20-30 meshes.
3. Standard sulphate solution: Dissolve 147.9 mg anhydrous Na₂SO₄ and dilute to 1000ml. (1ml=100 mg SO₄).

Procedure

Take suitable volume of sample in nessler tube. Add 5ml of conditioning reagent and mix well. Stir the content with help of a stirrer. Add a spatula of barium chloride and continue the stirring for one or more minutes. Note down the absorbance at 420 nm after 2 minutes. Prepare the standard graph with standard sulphate solution of 0-40mg/L ranges.

Calculation

Mg/L of nitrate = Test absorbance * Slope * Dilution factor

                               Sum of concentration of Standards
 Where, Slope =   
                                Sum of absorbance of Standards

                                     ml sample + ml distilled water
 D = Dilution ratio =   
                                                  ml sample

8.16 FLUORIDE: Fluoride may occur naturally in water, in rare instances the naturally occurring fluoride concentration may approach 10 mg/L; such waters should be defluoridated. Accurate determination of fluoride has increased in importance with the growth of the practice of fluoridation of water supplies as a public health measure. Maintenance of an optimal fluoride concentration is essential in maintaining effectiveness and safety of the fluoridation procedure (APHA, 1985).

Fluorides have dual significance in water supplies. High concentration causes dental fluorosis and lower concentration (<0.8 mg/L) causes dental caries. A fluoride concentration of approximately 1mg/L in drinking water is recommended. They are frequently found in certain industrial processes resulting in fluoride rich wastewaters. Significant sources of fluoride are found in coke, glass and ceramic, electronics, pesticide and fertiliser manufacturing, steel and aluminium processing and electroplating industries. It is calculated by SPADNS method (http://wgbis.ces.iisc.ernet.in/energy/monograph1/Frames.html).

Principle

The colorimetric method of estimating fluoride is based on the reaction of fluorides (HF) with zirconium SPADNS solution and the 'lake' (colour of SPADNS reagent), which is greatly influenced by the acidity of the reaction mixture. Fluoride reacts with the dye ‘lake’, dissociating (bleaching) the dye into a colourless complex anion (ZrF6 2-). As the amount of fluoride increases, the colour produced becomes progressively higher or of different hue.

Apparatus required: Spectrophotometer and lab glassware.

Reagents

1. Standard fluoride solution:
1. Stock: 221.0mg of AR grade sodium fluoride was dissolved in distilled water and made up to 1000ml to give 1ml = 100 m g of F-
2. Working Standard:100ml of the stock fluoride was diluted to 1000ml to give 1ml = 10 mg of fluoride.
2. SPADNS solution: Dissolve 958 mg SPADNS, sodium 2-(parasulfophenylazo)-1,8-dihydroxy-3,6-napthalene disulfonate, also called 4,5-dihydroxy-3-(parasulfophenylazo)-2,7-napthalenedisulfonic acid trisodium salt, in distilled water and dilute to 500 mL.
3. Zirconyl-acid reagent: Dissolve 133 mg zirconyl chloride octahydrate, ZrOCl2.8H2O, in about 25 mL distilled water. Add 350 mL concentrated HCl and dilute to 500 mL with distilled water.
4. Sodium arsenite solution: Dissolve 5.0 g NaAsO2 and dilute to 1 L with distilled water. (Toxic)

Procedure

Take suitable volume (50 mL) of sample in nessler tube. If the sample contains residual chlorine, remove it by adding 1 drop (0.05 mL) NaAsO2 solution. Add 5ml each of SPADNS solution and Zirconyl-acid reagent mix well. Note down the absorbance at 570 nm. Prepare the standard graph with standard fluoride solution of 0 - 1.40 mg F -/L ranges.

Calculation
When the prepared 0 mg F -/L standard is used to set the photometer, alternatively calculate fluoride concentrations as follows:

                                                          A0 - Ax
                                    mg F -/L=   
                                                          A0 – A1

Where:
A0 = absorbance of the prepared 0 mg F -/L standard,
A1 = absorbance of the prepared 1.0 mg F -/L standard, and
Ax = absorbance of the prepared sample.

8.17 Dissolved Oxygen

Oxygen is essential to all forms of aquatic life, including those organisms responsible for the self-purification processes in natural waters. DO depend upon the physical, chemical and biochemical activities in the water body. The analysis for DO is a key test in water pollution and waste treatment control. Low levels of DO frequently indicate a high concentration of decaying organic matter in the water. As bacteria digest organic matter, they use up oxygen, leaving little for the other aquatic creatures. The factors affecting the DO are (1) temperature, as temperature increases, the amount of oxygen (or any gas) dissolved in water decreases. (2) Partial pressure of O2 in contact with water, at high altitudes, less oxygen is dissolved in water because the partial pressure of oxygen in the atmospheric is low (3) Salinity the solubility of gases O₂, CO₂ decreases with increasing salinity (Weiner, 2000).

D.O (mg/L)	Water quality
Above 8.0	Good
6.5-8.0	Slightly polluted
4.5-6.5	Moderately polluted
4.0-4.5	Heavily polluted
Below 4.0	Severely polluted

Estimation of dissolved oxygen by Winkler’s iodometric method

When manganous sulphate is added to the sample containing alkaline potassium iodide, manganous hydroxide is formed, which is oxidized by the dissolved oxygen in the sample to basic manganic oxide. The basic manganic oxide liberates iodine equivalent to that of dissolved oxygen originally present in the sample. The liberated iodine is titrated with standard solution of sodium thiosulphate using starch as the indicator.

REAGENTS

Manganous sulphate solution: Dissolved 100 g of manganous sulphate in 200ml of distilled water and heated to dissolve salt and filtered after cooling.

1. Alkaline potassium iodide solution: Dissolved 100 g of KOH and 50 g of KI in 200ml of preboiled distilled water.
2. Sodium thiosulphate 0.025 N: Dilute 62.50 ml of 0.1N sodium thiosulphate in a preboiled distilled water and made up to the volume of 250ml.
3. Starch indicator solution: Dissolved 1 g of starch in 100ml of distilled water and warmed for complete dissolution.
4. Concentrated sulphuric acid.

Procedure

The sample was collected in 125ml BOD bottle carefully without allowing air bubbles. Added 1ml of manganous sulphate and 1ml of alkali iodide – azide reagent. A brown precipitate of basic manganic oxide formed was allowed to settle. Added 1ml of concentrated sulphuric acid and mixed well until the precipitate dissolved. About 25ml of the solution was taken and titrated against sodium thiosulphate until a straw yellow colour appeared. Few drops of starch indicator was added and titrated again until the blue colour disappeared (Manivasakam, 1997).

CALCULATION

                                                 (ml*N) of sodium thiosulphate * 8 * 1000
Dissolved oxygen, mg/L =  
                                                                  V2 [(V1 - V)/V1]
Where,
V1 = Volume of sample bottle
V2 = Volume of contents titrated
V = Volume of MnSO4 and KI added (2ml)

8.18 Biological Analysis

8.18.1 Faecal Coliform Bacteria: In general, coliform bacteria can be divided into a fecal and a non-fecal group.  Fecal coliform bacteria occur normally in the intestines of humans and other warm-blooded animals.  They are discharged in great numbers in human and animal wastes.  Their absence from water is thus a good indication that the water is bacteriologically safe for human consumption.

Standard Total Coliform Multiple Tube Tests
The total coliform bacteria test includes both Escherichia and Aerobacter coliform bacteria groups.  Aerobacter and some Escherichia can grow in soil.  Therefore, not all coliforms found in the total coliform test come from human wastes.  Escherichia coli (E. coli) apparently are all of fecal origin.  The Coliform bacteria that ferment lactose with gas formation with in 48 h at 35°C (APHA, 1985)

Medium composition

Peptone, Dipotassium hydrogen phosphate, Ferric ammonium citrate, sodium thiosulphate, 1ml Teepol and 50 ml distilled water.

Procedure

The medium was taken in to three set of test tubes 10, 1.0 and 0.1 mL respectively, each set had 5 test tubes. The tubes are then sterilized and incubated at 35 ± 0.5º C for 24 to 48 hours. The water sample was added to each set of test tubes 40, 9 and 1 mL respectively. Each of which contains lactose broth or lauryl tryptose broth and an inverted tube.  The tubes are then incubated at 35 ± 0.5º C and each test tube were examined for growth, gas or acidic reaction after 24 and 48 hours. Production of gas in the test tubes within 48 ± 3 h constitutes a positive reaction. The numbers of positive results for each set of proportion were noted down. Then the density of organisms were determined from the MPN index (908: V) given in the APHA standard methods.