Nitrogen Ammonia Nh3 Nitrite No2 And Nitrate No3

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Nitrogen compounds of greatest interest to water quality are those that are biologically available as nutrients to plants or exhibit toxicity to humans or aquatic life. Atmospheric nitrogen (N2) is the primary source of all nitrogen species, but it is not directly available to plants because the N=N triple bond is too strong to be broken by photosynthesis. Atmospheric nitrogen must be converted to other nitrogen compounds before it can become available as a plant nutrient.

The conversion of atmospheric nitrogen to other chemical forms is called nirogen fixation and is accomplished by a few types of bacteria that are present in water, soil, and root nodules of alfalfa, clover, peas, beans, and other legumes. Atmospheric lightning is another significant source of fixed nitrogen because the high temperatures generated in lightning strikes are sufficient to break N2 and O2 bonds making possible the formation of nitrogen oxides. Nitrogen oxides created within lightning bolts are dissolved in rainwater and absorbed by plant roots, thus entering the nitrogen nutrient sub-cycles, (see Figure 3.5). The rate at which atmospheric nitrogen can enter the nitrogen cycle by natural processes is too low to support today's intensive agricultural production. The shortage of fixed nitrogen must be made up with fertilizers containing nitrogen fixed by industrial processes, which are dependent on petroleum fuel. Modern large-scale farming has been called a method for converting petroleum into food.

Ammonia Soluble Water
FIGURE 3.7 Nitrogen cycle. The Nitrogen Cycle

As illustrated in Figure 3.7, in the nitrogen cycle, plants take up ammonia and nitrogen oxides dissolved in soil pore water and convert them into proteins, DNA, and other nitrogen compounds. Animals get their nitrogen by eating plants or other plant-eating animals. Once in terrestrial ecosystems, nitrogen is recycled through repeated biological birth, growth, death, and decay steps. There is a continual and relatively small loss of fixed nitrogen when specialized soil bacteria convert fixed nitrogen back into nitrogen gas (denitrification), which then is released to the atmosphere until it can reenter the nutrient sub-cycles again.

When nitrogen is circulating in the nutrient sub-cycles, it undergoes a series of reversible oxidation-reduction reactions that convert it from nitrogenous organic molecules, such as proteins, to ammonia (NH3), nitrite (NO2-), and nitrate (NO3). Ammonia is the first product in the oxidative decay of nitrogenous organic compounds. Further oxidation leads to nitrite and then to nitrate. Ammonia is naturally present in most surface and wastewaters. Its further degradation to nitrites and nitrates consumes dissolved oxygen.

Organic N-> NH3 +2 > NO2- +2 > NO3-. (3.15)

Ammonia/Ammonium Ion (NH3/NH4+)

In water, ammonia reacts as a base, raising the pH by generating OH- ions, as in Equation 3.16.

The equilibrium of Equation 3.16 depends on pH and temperature, (see Figure 3.6). In a laboratory analysis, total ammonia (NH3 + NH4+) is measured and the distribution between unionized ammonia (NH3) and ionized ammonia (NH4+) is calculated from knowledge of the water pH and temperature at the sampling site. Since the unionized form is far more toxic to aquatic life than the ionized form, field measurements of water pH and temperature at the sampling site are very important. The two forms of ammonia have different mobilities in the environment. Ionized ammonia is strongly adsorbed on mineral surfaces where it is effectively immobilized. In contrast, unionized ammonia is only weakly adsorbed and is transported readily by water movement. If a suspended sediment carrying sorbed NH4+ is carried by a stream into a zone with a higher pH, a portion will be converted to unionized NH3, which can then desorb and become available to aquatic life forms as a toxic pollutant. Unionized ammonia is also volatile and a fraction of it is transported as a gas.

As discussed above, nitrogen passes through several different chemical forms in the nutrient sub-cycle. In order to allow quantities of these different forms to be directly compared with one another, analytical results often report their concentrations in terms of their nitrogen content. For example, 10.0 mg/L of unionized ammonia may be reported as 8.23 mg/L NH3-N (ammonia nitrogen),* or 10.0 mg/L of nitrate reported as 2.26 mg/L NO3-N (nitrate nitrogen).**

Rules of Thumb

1. Ammonia toxicity increases with pH and temperature.

2. At 20°C and pH > 9.4, the equilibrium of Equation 3.16 is to the left, favoring NH3, the toxic form.

3. At 20°C and pH < 9.4, the equilibrium of Equation 3.16 is to the right, favoring NH4+, the nontoxic form.

4. A temperature increase shifts the equilibrium to the left, favoring the NH3 form.

5. NH3 concentrations >0.5 mg NH3-N/L cause significant toxicity to fish.

6. The unionized form is volatile, or air-strippable. The ionized form is nonvolatile.

Changes in environmental conditions can cause an initially acceptable concentration of total ammonia to become unacceptable and in violation of a stream standard. For example, consider the case of a wastewater treatment plant that discharges its effluent into a detention pond that, in turn, periodically releases its water into a stream. The treatment plant is meeting its discharge limit for unionized ammonia when its effluent is measured at the end of its discharge pipe. However, the detention pond is prone to support algal growth. In such a situation, it is not unusual for algae to grow to a level that influences the pond's pH. During daytime photosynthesis, algae may remove enough dissolved carbon dioxide from the pond to raise the pH and shift the equilibrium of Equation 3.16 to the left, far enough that the pond concentration of NH3 becomes higher than the discharge permit limit. In this case, discharges from the pond could exceed the stream standard for unionized ammonia even though the total ammonia concentration is unchanged.

Example 3.9

Ammonia is removed from an industrial wastewater stream by an air-stripping tower. To meet the effluent discharge limit of 5-ppm ammonia, the influent must be adjusted so that 60% of the total ammonia is in the volatile form. To what pH must the influent be adjusted if the wastewater in the stripping tower is at 10°C? Use Figure 3.8.

* Calculated as follows: -atomic wt. of N- x ^c. of NH3 = 14 x 10 mg/L = 8.23 mg/L NH3-N.

molecular wt. of NH3 17

** Calculated as follows: -atomic wt. of N- x conC. of nh3 = — x 10 mg/L = 2.26 mg/L NO3-N.

molecular wt. of NO, 62

Ammonia Curve
FIGURE 3.8 Percent unionized ammonia (NH3) as a function of pH and temperature.

Answer: In Figure 3.8, the 60% unionized ammonia gridline crosses the 10° C curve between pH = 9.7 and 9.8. Thus, the influent must be adjusted to pH = 9.8 or higher to meet the discharge limit.

Criteria and Standards for Ammonia

Typical state standards for unionized ammonia (NH3) are

• Aquatic life: Cold water biota = 0.02 mg/L NH3-N, chronic; warm water biota = 0.06 mg/L NH3-N, chronic; acute standard calculated from temperature and pH.

• Domestic water supply: 0.05 mg N/L total (NH3 + NH4+), for a 30-day average.

Ammonia and other nitrogenous materials in natural waters tend to be oxidized by aerobic bacteria, first to nitrite and then to nitrate. Therefore, all organic compounds containing nitrogen should be considered as potential nitrate sources. Organic nitrogen compounds enter the environment from wild animal and fish excretions, human sewage, and livestock manure. Inorganic nitrates come primarily from manufactured fertilizers containing ammonium nitrate and potassium nitrate.

In oxygenated waters, nitrite is rapidly oxidized to nitrate, so normally there is little nitrite present in surface waters. Manufactured fertilizers are another source of nitrate and ammonia. Both nitrite and nitrate are important nutrients for plants, but they are toxic to fish and humans at high concentrations. Nitrates and nitrites are very soluble, do not adsorb readily to mineral and soil surfaces, and are very mobile in the environment. Consequently, where soil nitrate levels are high, contamination of groundwater by nitrate leaching is a serious problem. Unlike ammonia, nitrites and nitrates do not evaporate and remain in water until they are consumed by plants and microorganisms.

Rules of Thumb

1. Unpolluted surface waters normally contain only trace amounts of nitrite.

2. Measurable groundwater nitrite contamination is more common because of low oxygen concentrations in the soil's subsurface.

3. Nitrate and nitrite leach readily from soils to surface and groundwaters.

4. High concentrations ( >1-2 mg/L) of nitrate or nitrite in surface or groundwater generally indicate agricultural contamination from fertilizers and manure seepage.

5. Greater than 10 mg/L of nitrite and nitrate in drinking water is a human health hazard.

Drinking water standards for nitrate are strict because the nitrate ion is reduced to nitrite ion in the saliva of all humans and in the intestinal tracts of infants during the first six months of life. Nitrite oxidizes iron in blood hemoglobin from Fe2+ to Fe3+. The resulting compound, called meth-emoglobin, cannot carry oxygen. The resulting oxygen deficiency is called methemoglobinemia. It is especially dangerous in infants (blue baby syndrome) because of their small total blood volume.

Criteria and Standards for Nitrate

Typical state standards for nitrate (NO3) are

• Agriculture MCLs: Nitrate = 100 mg/L NO3-N; Nitrite = 10 mg/L NO2-N (1-day average).

• Domestic water supply MCLs: Nitrate = 10 mg/L NO3-N; Nitrite = 1.0 mg/L NO2-N (1-day average).

Example 3.10

COD caused by sodium nitrite disposal

A chemical company wished to dispose of 250,000 gallons of water containing 500 mg/L of nitrite into a municipal sewer system. The manager of the municipal wastewater treatment plant had to determine whether this waste might be detrimental to the operation of his plant.

Under oxidizing conditions that exist in the treatment plant, nitrite is oxidized to nitrate as follows:

The consumption of oxygen shown in Equation 3.17 makes nitrite useful as a rust-inhibiting additive in boilers, heat exchangers, and storage tanks by deoxygenating the water. When nitrite is added to a wastewater stream, it is the same as adding chemical oxygen demand. More oxygen will be needed to maintain aerobic treatment steps at their optimum performance level. In addition, it will produce additional nitrate that may have to be denitrified before it can be discharged.


For a wastewater stream of a total volume of 250,000 gallons that contains 500 mg/L of nitrite, the net weight of nitrite is

500 mg/L x 250,000 gal. x 3.79 L/gal. = 474 x 106 mg = 474,000 g of nitrite.

From Equation 3.17, stoichiometric consumption of oxygen is 1 mole (32 g) for each 2 moles (92 g) of nitrite oxidized, resulting in a 1 to 2.9 ratio of O2 to NO2- by weight. Therefore, 474,000 g of nitrite will potentially consume 163,000 g of dissolved oxygen. Whether or not this represents a significant additional COD depends on the operating specifications of the treatment plant.

Also from Equation 3.17, stoichiometric production of nitrate is 2 moles (124 g) for each 2 moles (92 g) of nitrite oxidized, resulting in a 1.3 to 1 ratio of NO3- to NO2- by weight. Therefore, 474,000 g of nitrite will produce 616,000 g (1,362 lbs) of nitrate that might have to be denitrified before release.

Methods for Removing Nitrogen from Wastewater

After the activated-sludge treatment stage, municipal wastewater generally still contains some nitrogen in the forms of organic nitrogen and ammonia. Additional treatment is required to remove nitrogen from the waste stream.

Air-stripping Ammonia

See Example 3.9. Air-stripping can follow the activated-sludge process. pH must be raised with lime to about 10 or higher to convert all ammoniacal nitrogen to the volatile NH3 form. Scaling, icing, and air pollution are some of the disadvantages of air-stripping; whereas an advantage is that raising the pH precipitates phosphorus in the form of calcium phosphate compounds.


This is a two-step process:

1. Ammonia and organic nitrogen are first biologically oxidized completely to nitrate under strongly aerobic conditions (nitrification). This is achieved by more than normal and extensive aeration of the sewage:

2 NH4 + + 3 O2 N"rosoaas > 4 H+ + 2 NO2- + 2 H2O.

2. Nitrate is then biologically converted to gaseous nitrogen under anaerobic conditions (denitrification). This requires a carbon nutrient source. Water that is low in total organic carbon (TOC) may require the addition of methanol or other carbon source.

4 NO3- + 5 {CH2O} + 4 H+ denfgbacea > 2 N2 (g) + 5 CO2 (g) + 7 H2O.

Break-point Chlorination

The chemical reaction of ammonia with dissolved chlorine results in denitrification by converting ammonia to chloramines and nitrogen gas (see Chapter 6). With continued addition of Cl2, nitrogen gas and a small amount of nitrate are formed. Any chloramine remaining serves as a weak disinfectant and is relatively nontoxic to aquatic life.

Ammonium Ion-exchange

This is a good alternative to air-stripping because an exchange resin, the natural zeolite clinoptilolite, has been developed and is selective for ammonia. NH4+ is exchanged for Na+ or Ca2+ on the resin. The zeolite can be regenerated with sodium or calcium salts.


The removal of biomass, produced in the sewage treatment system by filtering to reduce suspended solids, results in a net loss of nitrogen that has been incorporated in the biomass cell structure.


Sulfide is often present naturally in groundwater as the dissolved anion S2-, especially in natural hot springs. There, it arises from soluble sulfide minerals and anaerobic bioreduction of dissolved sulfates. Sulfide is also formed in surface waters from anaerobic decomposition of organic matter containing sulfur. It is a common product of wetlands and eutrophic lakes and ponds. Sulfide reacts with water to form hydrogen sulfide, H2S, a colorless, highly toxic gas that smells like rotten eggs. The human nose is very sensitive to the odor of low levels of H2S. The odor threshold for H2S dissolved in water is 0.03 to 0.3 |J.g/L.

There are two important sources of H2S in the environment: the anaerobic decomposition of organic matter containing sulfur, and the reduction of mineral sulfates and sulfites to sulfide. Both mechanisms require reducing, or anaerobic conditions, and are strongly accelerated by the presence of sulfur-reducing bacteria. H2S is not formed in the presence of an abundant supply of oxygen.

Blackening of soils, wastewater, sludge, and sediments in locations with standing water, in addition to the odor of rotten eggs, is an indication that sulfide is present. The black material results from a reaction of H2S with dissolved iron and other metals to form precipitated ferrous sulfide (FeS), along with other metal sulfides.

H2S can have two stages of dissociation under reducing conditions in water, depending on the pH:

• At pH = 5, about 99% of dissolved sulfide is in the form of H2S, the unionized form.

H2S is the most toxic and volatile form; HS- and S2- are nonvolatile and much less toxic. H2S > 2.0 |j.g/L constitutes a long-term hazard to fish.

Rules of Thumb

1. Well water smelling of H2S is usually a sign of sulfate-reducing bacteria. Look for a water redox potential <-200 mV and a sulfate (SO42-) concentration in groundwater >100 mg/L.

2. A typical concentration of H2S in unpolluted surface water is <0.25 | g/L.

3. H2S > 2.0 | g/L constitutes a chronic hazard to aquatic life.

4. In aerated water, H2S is bio-oxidized to sulfates and elemental sulfur.

5. Unionized H2S is volatile and air-strippable. The ionized forms, HS- and S2-, are nonvolatile.

Typical state standards for unionized H2S are

• Aquatic life (cold and warm water biota): 2.0 |j.g/L (30-day).



Phosphorus is a common element in igneous and sedimentary rocks and in sediments but it tends to be a minor element in natural waters because most inorganic phosphorus compounds have low solubility. Dissolved concentrations are generally in the range of 0.01-0.1 mg/L and seldom exceed 0.2 mg/L. The environmental behavior of phosphorus is largely governed by the low solubility of most of its inorganic compounds, its strong adsorption to soil particles, and its importance as a nutrient for biota.

Because of its low dissolved concentrations, phosphorus is usually the limiting nutrient in natural waters. The dissolved phosphorus concentration is often low enough to limit algal growth. Because phosphorus is essential to metabolism, it is always present in animal wastes and sewage. Too much phosphorus in wastewater effluent is frequently the main cause of algal blooms and other precursors of eutrophication.

Important Uses for Phosphorus

Phosphorus compounds are used for corrosion control in water supply and industrial cooling water systems. Certain organic phosphorus compounds are used in insecticides. Perhaps the major commercial uses of phosphorus compounds are in fertilizers and in the production of synthetic detergents. Detergent formulations may contain large amounts of polyphosphates as "builders," to sequester metal ions and maintain alkaline conditions. The widespread use of detergents instead of soap has caused a sharp increase in available phosphorus in domestic wastewater.

Prior to the use of phosphate detergents, most wastewater inorganic phosphorus was contributed from human wastes; about 1.5 g/day per person is released in urine. As a consequence of detergent use, the concentration of phosphorus in treated municipal wastewaters has increased from 3-4 mg/L in pre-detergent days, to the present values of 10-20 mg/L. Since phosphorus is an essential element for the growth of algae and other aquatic organisms, rapid growth of aquatic plants can be a serious problem when effluents containing excessive phosphorus are discharged to the environment.

The Phosphorus Cycle

In a manner similar to nitrogen, phosphorus in the environment is cycled between organic and inorganic forms. An important difference is that under certain soil conditions, some nitrogen is lost to the atmosphere by ammonia volatilization and microbial denitrification. There are no analogous gaseous loss mechanisms for phosphorus. Also important are the differences in mobility of the two nutrients. Both exist in anionic forms (NO2/NO3- and H2PO4/HPO42-) which are not subject to retention by cation exchange reactions. However, nitrate anions do not form insoluble compounds with metals and, therefore, readily leach from soil into surface and groundwaters. Phosphate anions are largely immobilized in the soil by the formation of insoluble compounds — chiefly iron, calcium, and aluminum phosphates — and by adsorption to soil particles. Nitrogen compounds leach more readily than phosphorus compounds from soils into ground and surface waters which contribute to a phosphorus-limited algal growth in most surface waters. The critical level of inorganic phosphorus for forming algal blooms can be as low as 0.01 to 0.005 mg/L under summer growing conditions but more frequently is around 0.05 mg/L.

Organic compounds containing phosphorus are found in all living matter. Orthophosphate (PO43-) is the only form readily used as a nutrient by most plants and organisms. The two major steps of the phosphorus cycle, conversion of organic phosphorus to inorganic phosphorus and back to organic phosphorus, are both bacterially mediated. Conversion of insoluble forms of phosphorus, such as calcium phosphate, Ca(HPO4)2, into soluble forms, principally PO43-, is also carried out by microorganisms. Organic phosphorus in tissues of dead plants and animals, and in animal waste products is converted bacterially to PO43-. The PO43- thus released to the environment is taken up again into plant and animal tissue.

Mobility in the Environment

Phosphorus is an important plant nutrient and is often present in fertilizers to augment the natural concentration in soils. Phosphorus is also a constituent of animal wastes. Runoff from agricultural

Rules of Thumb

1. In surface waters, phosphorus concentrations are influenced by the sediments, which serve as a reservoir for adsorbed and precipitated phosphorus. Sediments are an important part of the phosphorus cycle in streams. Bacterially mediated exchange between dissolved and sediment-adsorbed forms plays a role in making phosphorus available for algae and therefore contributes to eutrophication.

2. In streams, dissolved phosphorus from all sources, natural and anthropogenic, is generally present in low concentrations, around 0.1 mg/L or less.

3. The natural background of total dissolved phosphorus has been estimated to be about 0.025 mg-P/L; that of dissolved phosphates about 0.01 mg-P/L.

4. The solubility of phosphates increases at low pH and decreases at high pH.

5. Particulate phosphorus (sediment-adsorbed and insoluble compounds) is about 95% of the total phosphorus in most cases.

6. In carbonate soils, dissolved phosphorus can react with carbonate to form the mineral precipitate hydroxyapatite (calcium phosphate hydroxide), Ca10(PO4)6(OH)2.

areas is a major contributor to total phosphorus in surface waters, where it occurs mainly in sediments because of the low solubility of its inorganic compounds and its tendency to adsorb strongly to soil particles.

Dissolved phosphorus is removed from solution by

• Precipitation

• Strong adsorption to clay minerals and oxides of aluminum and iron

• Adsorption to organic components of soil

Reducing and anaerobic conditions, as in water-saturated soil, increase phosphorus mobility because insoluble ferric iron, to which phosphorus is strongly adsorbed, is reduced to soluble ferrous iron, thereby releasing adsorbed phosphorus. In acid soils, aluminum and iron phosphates precipitate, while in basic soils, calcium phosphates precipitate. The immobilization of phosphorus is therefore dependent on soil properties, such as pH, aeration, texture, cation-exchange capacity, the amount of calcium, aluminum and iron oxides present, and the uptake of phosphorus by plants.

Because of these removal mechanisms for dissolved phosphorus, phosphorus compounds resist leaching, and there is little movement of phosphorus with water drainage through most soils. It is mobilized mainly with erosion sediments. Phosphorus transport into surface waters is controlled chiefly by preventing soil erosion and controlling sediment transport. In most soils, except for those that are nearly all sand, almost all the phosphorus applied to the surface is retained in the top 1 to 2 ft. The adsorption capacity for phosphorus has been estimated for several soils to be in the range of 77 to over 900 lbs/acre-ft of soil profile. Often, the total phosphorus removal capacity for a soil will exceed the planning life of a typical land application project. If the phosphorus-removing capacity of a soil becomes saturated, it usually can be restored in a few months, during which adsorbed phosphorus is precipitated with metals or removed by crops.

Dissolved phosphate species exhibit the following pH-dependent equilibria (see Figure 3.9):

H3PO4 O H2PO4- + H+ o HPO42- + 2 H+ o PO43- + 3 H+ (3.19)

• Below pH 2, H3PO4 is the dominant species.

• Between pH 2 and pH 7, H2PO4- is the dominant species.

• Between pH 7 and pH 12, HPO42- is the dominant species.

• Above pH 12, PO43- is the dominant species.

H3po4 Species

FIGURE 3.9 pH dependence of phosphate species.

FIGURE 3.9 pH dependence of phosphate species.

Phosphate Species

FIGURE 3.10 Forms of immobile phosphorus. General relationships between soil pH and phosphorus reactions are

• In the acid range, dissolved phosphorus is predominantly H2PO4-, and immobile phosphorus is bound with iron and aluminum compounds.

• In the basic range, dissolved phosphorus is predominantly HPO42-, and immobile phosphorus is mainly in the form of calcium phosphate.

• Maximum availability of phosphorus for plant uptake (as well as leaching) occurs between pH 6 and 7.

Soil pH

FIGURE 3.10 Forms of immobile phosphorus. General relationships between soil pH and phosphorus reactions are

• In the acid range, dissolved phosphorus is predominantly H2PO4-, and immobile phosphorus is bound with iron and aluminum compounds.

• In the basic range, dissolved phosphorus is predominantly HPO42-, and immobile phosphorus is mainly in the form of calcium phosphate.

• Maximum availability of phosphorus for plant uptake (as well as leaching) occurs between pH 6 and 7.

Whole-lake experiments [D.W. Schindler, Science, 184, 897 (1974); 195, 260 (1977)] have demonstrated that, even when algal growth in lakes is temporarily limited by carbon or nitrogen instead of phosphorus, natural long-term mechanisms act to compensate for these deficiencies. Carbon deficiencies are corrected by CO2 diffusion from the atmosphere, and nitrogen deficiencies are corrected by changes in biological growth mechanisms. Therefore, even if a sudden increase in phosphorus occurs temporarily causing algal growth to be limited by carbon or nitrogen, eventually these deficiencies are corrected. Then, algal growth becomes proportional to the phosphorus concentration as the system becomes once more phosphorus-limited.

Phosphorus Compounds

Compounds containing phosphorus that are of interest to water quality include:

Orthophosphates (all contain PO43) Trisodium phosphate — Na3PO4 Disodium phosphate — Na2HPO4 Monosodium phosphate — NaH2PO4 Diammonium phosphate — (NH4)2HPO4.

Orthophosphates are soluble and are considered the only biologically available form. In the environment, hydrolysis slowly converts polyphosphates to orthophosphates. Analytical methods measure orthophosphate. To measure total phosphate, all forms of phosphate are chemically converted to orthophosphates (hydrated forms).

Polyphosphates (also called condensed phosphates, meaning dehydrated) Sodium hexametaphosphate — Na3(PO4)6 Sodium tripolyphosphate — Na5P3O10 Tetrasodium pyrophosphate — Na4P2O7

Organic phosphate (biodegradation or oxidation of organic phosphates releases orthophosphates).

Rules of Thumb

1. The critical level of inorganic phosphorus for algae bloom formation can be as low as 0.01 to 0.005 mg/L under summer growing conditions but more frequently is around 0.05 mg/L.

2. Lakes are nitrogen-limited if the ratio of total nitrogen to total phosphorus (N/P) is less than 13, nutrient-balanced if 13 < N/P < 21, and phosphorus-limited if N/P > 21. Exact ranges depend on the particular algae species. Most lakes are phosphorus limited; in other words, additional phosphorus is needed to sustain further algal growth.

3. Different N/P ratios and pH values favor the growth of different kinds of algae.

4. Low N/P ratios favor N-fixing blue-green algae.

5. High N/P ratios, often achieved by controlling phosphorus input by means of additional wastewater treatment, cause a shift from blue-green algae to less objectionable species.

6. Lower pH (or increased CO2) gives green algae a competitive advantage over blue-green algae.

Sedimentary phosphorus occurs in the following forms:

• Phosphate minerals: Mainly hydroxyapatite, Ca5OH(PO4)3.

• Nonoccluded phosphorus: Phosphate ions (usually orthophosphate) bound to the surface of SiO2 or CaCO3. Nonoccluded phosphorus is generally more soluble and more available than occluded phosphorus (below).

• Occluded phosphorus: Phosphate ions (usually orthophosphate) contained within the matrix structures of amorphous hydrated oxides of iron and aluminum and amorphous alumino-silicates. Occluded phosphorus is generally less available than nonoccluded phosphorus.

• Organic phosphorus: Phosphorus incorporated with aquatic biomass, usually algal or bacterial.

Removal of Dissolved Phosphate

Current remedies for phosphate-caused foaming and eutrophication are

• Using lower phosphate formulas in detergents,

• Precipitating the phosphate with Fe3+, Al3+, or Ca2+, and

• Diverting the discharge to a less sensitive location.

Phosphate removal is carried out in a manner similar to water softening by precipitation of Ca2+ and Mg2+. The usual precipitants for removing phosphate are alum (Al2(SO4)3), lime (Ca(OH)2), and ferric chloride (FeCl3). The choice of precipitant depends on the discharge requirements, wastewater pH, and chemical costs.

The pertinent reactions for the precipitation of phosphate with alum, ferric chloride, and lime are

Alum: Al2(SO4)3 + 2 HPO42- 2 AlPO4(s) + 3 SO42- + 2 H+.

Ferric chloride: FeCl3 + HPO42- FePO4(s) + 3 Cl- + H+.

Lime: 3 Ca2+ + 2 OH- + 2 HPO42- Ca3(PO4)2(s) + 2 H2O.

5 Ca2+ + 4 OH- + 3 HPO42- Ca5(OH)(PO4)3(s) + 3 H2O.

Where effluent concentrations of phosphorus up to 1.0 mg/L are acceptable, the use of iron or aluminum salts in a wastewater secondary treatment system is often the process of choice. If very low levels of effluent phosphorus are required, precipitation at high pH by lime in a tertiary unit is necessary. The lowest levels of phosphorus are achieved by adding NaF with lime to form Ca5(PO4)3F (fluorapatite). The operating pH for phosphate removal with lime is usually above 11 because flocculation is best in this range.

If alkalinity is present, aluminum and iron ions are consumed in the formation of metal-hydroxide flocs. This may increase required dosages by up to a factor of 3. Calcium ions react with alkalinity to form calcium carbonate. Thus, the amount of precipitant needed for phosphate precipitation is controlled more by the alkalinity than the stoichiometry of the reaction. In the case of aluminum and iron precipitants, the reaction with alkalinity is not totally wasted because the hydroxide flocs assist in the settling and removal of metal-phosphate precipitates, along with other suspended and colloidal solids in the wastewater.

Biological phosphorus removal can be accomplished by operating an activated sludge process in an anaerobic-aerobic sequence. A number of bacteria respond to this sequence by accumulating large excesses of polyphosphate within their cells in volutin granules. During the anaerobic phase, a release of phosphate occurs. In the aerobic phase, the released phosphate and an additional increment is taken up and stored as polyphosphate, giving a net removal, coincident with organic removal and metabolism. Phosphate can be removed from the waste stream as sludge or through use of a second anaerobic step. During the second anaerobic step, the stored phosphate is released in dissolved form. Then, the bacterial cells can be separated and recycled and the released soluble phosphate removed by precipitation.



The commonly encountered elemental metals may be divided into three general classes:

1. Alkali metals: Li, Na, K (Periodic Table Group 1A).

2. Alkaline metals: Be, Mg, Ca, Sr, Ba (Periodic Table Group 2A).

3. Heavy metals: All metals to the right of the alkali and alkaline metals in the Periodic Table.

Metals in natural waters may be in dissolved or particulate forms.

Dissolved forms are

• Complexes: Zn(OH)42+, Au(CN) 2-, Ca(P2O7)2-, PuEDTA, etc.

• Organometallics: Hg(CH3)2, B(C2H5)3, Al(C2H5)3, etc.

Particulate forms are

• Mineral sediments

• Precipitated oxides, hydroxides, sulfides, carbonates, etc.

• Cations sorbed to sediments

A metal water quality standard may be written for the dissolved, potentially dissolved, total recoverable, or total form.

• Dissolved: Sample is filtered on site through a 0.45-micron filter, then acidified to pH 2 for preservation before analysis. Acidification prevents precipitation of any dissolved metal before analysis. This procedure omits from the analysis metals adsorbed on suspended sediments.

• Potentially dissolved: Sample is acidified to pH 2, held for 72 to 90 hours, then filtered through a 0.45-micron filter and analyzed. This procedure is intended to simulate the possibility that metals bound in suspended sediments might be transported into more acidic conditions and might partially dissolve. It measures the metals dissolved at the time of sampling, in addition to a portion of the metals bound to suspended sediments.

• Total recoverable: Sample is acidified to pH 2 and analyzed without filtering. This procedure measures all metals, dissolved and bound to suspended sediments.

• Total: Sample is "digested" in an acidic solution until essentially all the metals present are extracted into soluble forms for analysis.

General Behavior of Dissolved Metals in Water

Because water molecules are polar, metal cations always attract a hydration shell of water molecules by electrostatic attraction to the positive charge of the cation, as illustrated in Figure 3.11.

M+n _H2°® M(H2O)x+n, x = 6 for most cations. (3.20)

Hydrated metal ions behave as acids by donating protons (H+) to H2O molecules, forming the acidic H3O+ hydronium ion. The process can continue stepwise up to n times to make a neutral metal hydroxide:

M(H2O)5OH+(n-1) + H2O o M(H2O)4(OH)2+(n-2) + H3O+, etc. up to n times. (3.22)

For example, with Fe3+, it takes 3 proton transfer steps to form neutral ferric hydroxide:

The overall reaction is

Group Hydroxide Solubility Hydration

FIGURE 3.11 Water molecules form a hydration shell around dissolved metal cations. Molecules in the hydration shell can lose a proton to bulk water molecules, as indicated by the arrow, leaving a hydroxide group bonded to the metal. This way, the hydrated metal behaves as an acid. Eventually, the metal may precipitate as a hydroxide compound of low solubility.

FIGURE 3.11 Water molecules form a hydration shell around dissolved metal cations. Molecules in the hydration shell can lose a proton to bulk water molecules, as indicated by the arrow, leaving a hydroxide group bonded to the metal. This way, the hydrated metal behaves as an acid. Eventually, the metal may precipitate as a hydroxide compound of low solubility.

With each step, the hydrated metal is progressively deprotonated, forming polyhydroxides and becoming increasingly insoluble. At the same time, the solution becomes increasingly acidic. Eventually, the metal precipitates as a low solubility hydroxide. The degree of acidity induced by metal hydration is greatest for cations of high charge and small size. All metal cations with a charge of +3 or more are moderately strong acids. This process is one source of acidic water draining from mines.

Rule of Thumb

Only polyvalent cations (e.g., Fe3+, Zn2+, Mn2+, Cr3+) attract water molecules strongly enough to act as acids by causing the release of H+ from water molecules in the hydration sphere. Monovalent cations, such as Na+, do not act as acids at all.

Lowering the pH (increasing H3O+) shifts the equilibrium of Equation 3.26 to the left, tending to dissolve any solid metal hydroxide that has precipitated. Raising the pH (adding OH) consumes H3O+ and shifts the equilibrium of Equation 3.26 to the right, precipitating an insoluble metal hydroxide. However, if the pH is raised too high, precipitated metal hydroxides can redissolve (see Figure 3.12). At high pH values, a metal hydroxide may form complexes with OH- anions to become a negatively charged ion having increased solubility. For example, precipitated Fe(OH)3 can react with OH- anions as follows:

Solubility Curve

FIGURE 3.12 Solubilities of some metals precipitating as hydroxides vs. pH.

FIGURE 3.12 Solubilities of some metals precipitating as hydroxides vs. pH.

The negatively charged polyhydroxide anions are more soluble because their ionic charge attracts them strongly to polar water molecules. As shown in Figure 3.10, the high value of pH where solubility begins to increase again varies from metal to metal. Hard and alkaline water provides a buffer against pH changes. In hard or alkaline water, the tendency of metals to make water acidic is diminished.

Example 3.11: Effect of Dissolved Metal on Alkalinity

A sample of groundwater contains a high concentration of dissolved iron, about 20 mg/L. At the laboratory, alkalinity is measured to be 150 mg/L as CaCO3. Does this laboratory measurement of alkalinity accurately represent the groundwater alkalinity?

Answer: Soluble inorganic iron is in the ferrous form, Fe2+. When a groundwater sample is exposed to air, oxygen oxidizes Fe2+ to the ferric form, Fe3+. This process is often enhanced by aerobic iron bacteria. Depending on the pH, hydrated Fe3+ can lose protons from its hydration sphere to any bases present, forming ferric hydroxide species and making the solution more acidic. Loss of protons from the hydration sphere is not a significant process for hydrated ferrous iron Fe2+. The acidic behavior of hydrated Fe3+ occurs to a greater extent at a higher pH. Equation 3.29 represents the overall reaction converting dissolved ferrous iron to precipitated ferric hydroxide:

4 Fe(H2O)62+ + O2 o 4 Fe(OH)3(s) + 14 H2O + 8 H+. (3.29)

Fe(OH)3 is a yellow to red-brown precipitate often seen on rocks and sediments in surface waters with high iron concentrations.

The concentration of H+ formed by Equation 3.29 can be up to 2 times the Fe2+ concentration, depending on the final pH. Each H+ released will neutralize a molecule of base, consuming some alkalinity, by reactions such as

We will assume the pH is high enough that the equilibrium of Equation 3.29 goes essentially to completion to the right side, a worst case scenario from the viewpoint of affecting the alkalinity. The atomic weights of hydrogen and iron are 1 g/mol and 56 g/mol, respectively. If the equilibrium of Equation 3.29 is completely to the right, one mole (56 g) of Fe3+ will make 2 moles of H+ (2 g). At the time of sampling, the concentration of dissolved Fe2+ was about 20 mg/L and all is eventually oxidized to Fe3+. The molar concentration of iron is

0.020 g/L = o.ooo36 mol/L, or 9.35 mmol/L. 56 g/mol

The moles of H+ produced are two times the moles of iron:

Moles of H+ = 2 x 0.36 mmol/L = 0.72 mmol/L, or 1 mg/mmol x 0.72 mmol/L = 0.72 mg/L.

We must now determine what effect this quantity of H+ will have on the alkalinity. Alkalinity is measured in terms of a comparable quantity of CaCO3. The molecular weight of CaCO3 is 100, and it dissolves to form the doubly charged ions Ca2+ and CO32-. Alkalinity is a property of the CO32- anion, which reacts to accept 2 H+ cations:

therefore, 0.72 mmol/L of H+ will react with 0.072/2 = 0.36 mmol/L of CO32-, and 0.36 mmol/L of CaCO3 are required as a source of the CO32-. From the definition of alkalinity, the change in alkalinity is equal to the change in concentration of CaCO3, in mg/L.

0.36 mmol/L of CaCO3 = 0.36 mmol/L x 100 mg = 36 mg/L = change in alkalinity

1 mmol

Groundwater alkalinity at time of sampling = Laboratory measured alkalinity + Alkalinity lost by Equation 3.29

The original alkalinity of the groundwater before exposure to air was

This example illustrates the pH buffering effect of alkalinity. The addition of H+ to the solution by Equation 3.29 does not change the pH greatly as long as some alkalinity remains because the added H+ is taken up by carbonate species in the water.



The general term solids refers to matter that is suspended (insoluble solids) or dissolved (soluble solids) in water. Solids can affect 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 pretreatment to prevent scale formation. Water high in suspended solids may harm aquatic life by causing abrasion damage, clogging fish gills, harming spawning beds, and reducing photosynthesis by blocking sunlight penetration, among other consequences. On the other hand, hard water (caused mainly by dissolved calcium and magnesium compounds) reduces the toxicity of metals to aquatic life.

Total solids (sometimes called residue) are the solids remaining after evaporating the water from an unfiltered sample. It includes two subclasses of solids that are separated by filtering (generally with a filter having a nominal 0.45-micron or smaller pore size):

1. Total suspended solids (TSS, sometimes called filterable solids) in water are organic and mineral particulate matter that do not pass through a filter. They may include silt, clay, metal oxides, sulfides, algae, bacteria, and fungi. TSS is generally removed by floccu-lation and filtering. TSS contributes to turbidity, which limits light penetration for photosynthesis and visibility in recreational waters.

2. Total dissolved solids (TDS, sometimes called nonfilterable solids) are substances that would pass through a 0.45 micron filter but will remain as residue when the water evaporates. They may include dissolved minerals and salts, humic acids, tannin, and pyrogens. TDS is removed by precipitation, ion-exchange, and reverse osmosis. In natural waters, the major contributors to TDS are carbonate, bicarbonate, chloride, sulfate, phosphate, and nitrate salts. Taste problems in water often arise from the presence of high TDS levels with certain metals present, particularly iron, copper, manganese, and zinc.

The difference between suspended and dissolved solids is a matter of definition based on the filtering procedure. Solids are always measured as the dry weight, and careful attention must be paid to the drying procedure to avoid errors caused by retained moisture or loss of material by volatilization or oxidation.

Rules of Thumb

1. TSS is detrimental to fish health by decreasing growth, disease resistance, and egg development.

2. Suspended solids should be restricted so they do not reduce the maximum depth of photosynthetic activity by more than 10% from the seasonally established norm.

3. Water with a TDS < 1200 mg/L generally has an acceptable taste. Higher TDS can adversely influence the taste of drinking water and may have a laxative effect.

4. In water to be treated for domestic potable supply, a TDS < 650 mg/L is a preferred goal.

5. For drinking water, recommended TDS is < 500 mg/L; the upper limit is 1000 mg/L.

TDS and Salinity

TDS and salinity both indicate dissolved salts. Table 3.1 offers a qualitative comparison between the terms.


Comparison of TDS and Salinity

TDS Degree of Salinity

1000 - 3000 mg/L Slightly saline

3000 - 10,000 mg/L Moderately saline

10,000 - 35,000 mg/L Very saline

>35,000 mg/L Briny

Specific Conductivity and TDS

Specific conductivity is directly related to TDS and serves as a check on TDS measurements.

Rules of Thumb

1. Conductivity units are ^mhos/cm or nSiemens/cm:

2. TDS in mg/L can be estimated from a measurement of specific conductivity.

3. For seawater (NaCl-based)

4. For groundwater (carbonate or sulfate-based) TDS (mg/L) « (0.55 to 0.7) x (Sp. Cond. in |iS/cm).

5. If the ratio -^^ is demonstrated to be consistent, the simpler specific conductivity measure-

ment may sometimes be substituted for TDS analysis. TDS Test for Analytical Reliability

A calculated value for TDS may be used for judging the reliability of a sample analysis if all the important ions have also been measured. The TDS concentration should be equal to the sum of the concentrations of all the ions present plus silica. You can use either of the following equations to calculate TDS from an analysis or to check on the validity of analytical results. All concentrations are in mg/L.

TDS = sum of cations + sum of anions + silica or

TDS = 0.6(alkalinity) + Na+ + K+ + Ca2+ + Mg2+ + Cl- + SO42- + SiO3.

In any given analysis, it is unlikely that all the ions have been measured. Frequently, only the major ions (Na+, K+, Ca2+, Mg2+, Cl-, HCO 3-, SO42-) are necessary for the calculations, as other ion concentrations are likely to be insignificant by comparison.

Use the following guidelines for checking accuracy of a TDS analysis:

1. TDSmeas should always be equal to or somewhat larger than TDScalc because a significant ion contributor might not have been included in the calculation.

2. An analysis is acceptable if the ratio of measured-to-calculated TDS is in the range measured TDS

calculated TDS

3. If TDSmeas < TDScalc, the sample should be reanalyzed.

4. If TDSmeas > 1.2 x TDScalc, the sample should be reanalyzed, perhaps with a more complete set of ions.


Temperature affects all water uses.

• The solubility of gases such as oxygen and carbon dioxide decreases as water temperature increases.

• Biodegradation of organic material in water and sediments is accelerated with increased temperatures, increasing the demand on dissolved oxygen.

• Fish and plant metabolism depends on temperature.

Most chemical equilibria are temperature dependent. Important environmental examples are the equilibria between ionized and unionized forms of ammonia, hydrogen cyanide, and hydrogen sulfide.

Temperature regulatory limits are set to maintain a normal pattern of diurnal and seasonal fluctuations, with no changes deleterious to aquatic life. Maximum induced change is limited to a 3° C increase over a 4-hour period, lasting for 12 hours maximum.

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