Water Quality: Natural waters

Transcription

1 Water Quality: Natural waters Important parameters: Water temperature is typically measured by either a mercury or spirit thermometer, or (increasingly frequently), with a thermistor thermometer. Thermistors are small electronic devices whose electrical resistance varies directly with temperature. ph is defined as the negative logarithm of the hydrogen ion concentration of an aqueous solution. In a practical sense, it is actually a measure of hydrogen ion activity, which is not quite the same as hydrogen ion concentration, since ion activity varies with temperature. Thus, ph readings must be corrected for temperature. The most usual methods for measuring ph are electrometric (with a ph meter), and colorimetric (with ph indicator dyes and a color comparitor). ph papers (hydrion papers) are not accurate enough to use for determining water ph. The ph of natural waters varies from about 3.0 to perhaps 10.0 or so, but usually falls between 6.0 and 8.0. It's important to remember that ph is a logarithmic scale ranging from 0-14 with 7.0 representing a "neutral" ph. Because the scale is logarithmic, a ph of 6.0 represents 10 times as much hydrogen ion as a ph of 7.0, and a ph of 5.0 represents 100 times as much hydrogen ion as a ph 0f 7.0. Also, because of the logarithmic nature of the scale, ph values cannot be averaged in the normal manner. A log-average calculation procedure must be used. In actuality, since hydrogen ions are "naked" protons, and thus very reactive, they probably do not exist for very long in water because they react with something (usually a water molecule). It is therfore more accurate to refer to hydrogen ions in aqueous solution as hydronium ions: 2H2O -----> H3O + + OH - Alkalinity is defined as the ability of a water to neutralize a standard acid, usually 0.02 N (N/50) sulfuric acid. Classically, alkalinity is usually measured in a two step titrimetric procedure, the first step involving titration to the phenolphthalein endpoint (ph = 8.3), and the second step involving titration to the methyl orange endpoint (ph = 3.7). Phenolphthalein and methyl orange are indicator dyes that change color at different ph's. Because the color change at the endpoint is sharper, a bromcresol green-methyl red dye mixture is often substituted for methyl orange now. The titration may be done using a ph meter instead of dyes to determine the endpoints. Alkalinity values are reported as: (1) phenolphthalein alkalinity (2) methyl orange alkalinity (3) total alkalinity. Alkalinity in most natural waters is usually (but not always) due to the bicarbonate ion (HCO3 - ). Other sources of alkalinity include the carbonate ion(co3-2 ), the hydroxide ion

2 (OH - ) and a few other minor ions. One can use the alkalinity figures to determine the relative contribution of bicarbonate, carbonate, and hydroxide ions by the following table: Result of titration HYDROXIDE alkalinity CARBONATE alkalinity BICARBONATE alkalinity P = T P < 1/2 T 0 2P T - 2P P = 1/2 T 0 2 (T-P) 0 P> 1/2 T 2P-T 2 (T-P) 0 P = T T 0 0 P = Phenolphthalein Alkalinity ; T = Total Alkalinity T (Total Alkalinity ) = phenolphthalein alkalinity + methyl orange alkalinity Click HERE for a discussion of titrimetric methods of chemical analysis. Color in natural waters is generally due either to metals (eg. iron or manganese) or to dissolved organic compounds (especially humic acids). Color may be measure as a bulk property, that is, the color of lakes may be measured; or as a property of an individual water sample. In polluted waters, the color of an individual sample may be significant, and may be determined with the use of a colorimeter or visual color comparitor and reported in standard color units. Color standards are usually glass filters (visual comparitors) or standard solutions of potassium chloroplatinate (colorimetric measurements). The color of lakes is usually determined by a visual comparitor using the Forel-Ule Color Scale. Acidity is defined as the ability of water to neutralize a standard base, usually a 0.02N (N/50) sodium hydroxide solution. Phenolphthalein and methyl orange (or bromcresol green-methyl red) are used as the indicators. Acidity is generally reported as (1) free acidity, and (2) total acidity. Acidity in natural waters is generally slight to non-existent. If present, it is generally due to sulfuric acid produced by the oxidation of metallic sulfide minerals such as pyrite. In waters polluted by acid mine drainage or industrial wastes, acidity may be significant. Dissolved oxygen (D.O.) in natural waters is introduced by either diffusion from, or turbulent mixing with, air; or by the photosynthesis of aquatic plants. It varies inversely with (1) temperature, (2) the dissolved solids content of the water, and (3) with altitude. Oxygen is not very soluble in water and the maximum concentration of 02 in pure water at 0 o C (in equilibrium with air at sea level) is 14.6 mg/liter. Because of this low solubility, it is often a limiting factor to aquatic organisms. Oxygen values are usually near

3 saturation for turbulent streams, but may be well below saturation levels in the bottom waters of lakes, or slowly moving organically-polluted streams. In still or slowly miving water in which photosynthetic plants are very active, dissolved oxygen values may be well above saturation. Table 1: Influence of temperature and dissolved solids on Oxygen solubility in water* Temperature 0 mg/l Cl - 5,000 mg/l Cl - 10,000 mg/l Cl - 15,000 mg/l Cl - 20,000 mg/l Cl - 0 o C o C o C o C o C o C (77 o F) o C o C o C * Oxygen solubility expressed as saturation values (in mg/l). Source: Standard Methods for the Examination of Water and Wastewater. APHA (1980) Measurement of dissolved oxygen: Dissolved oxygen in water can be measured electrometrically with a dissolved oxygen meter, or by chemical analysis; usually the Winkler method or some variation thereof. The Winkler method employs the following reagents: Alkaline iodide-azide reagent. An aqueous solution of potassium hydroxide, potassium iodide, and sodium azide. The azide is added to eliminate interference from nitrate ion. Aqueous manganous sulfate solution Concentrated sulfuric acid N sodium thiosulfate titrant (classical) or Phenylarsine oxide (PAO) titrant (modern) Aqueous starch solution indicator Reactions: (1) 2Mn +2 + O2 + 4OH > 2MnO(OH)2 [precipitates as a white floc] (2) MnO(OH)2 [white floc] + 6I - + 6H + [from sulfuric acid] > Mn +2 [floc dissolves] + 2I3 - [yellow] + 3H2O

4 (3) 2H2O + I3 - [yellow] + PAO > 2HI + I - [colorless] + PAO(OH)2 (4) starch + I > deep blue complex ion Carbon dioxide in natural waters is due either to solution of atmospheric carbon dioxide from air (minor), from the respiration of aquatic organisms (major source), or from volcanic activity ( may be a major source locally). The chemistry of CO2 in water is fairly complex because it can exist in a number of different chemical states: free carbon dioxide (CO2), carbonic acid (H2CO3), bicarbonate ion (HCO3 - ), carbonate ion (CO3-2 ). These are reversibly interconvertible via the following reactions: (1) CO2 [free carbon dioxide] + H2O <-----> H2CO3 [carbonic acid] (2) H2CO3 <-----> H + + HCO3 - [bicarbonate ion] (3) HCO3 - <-----> H + + CO3-2 [carbonate ion] ph plays an important role in determining which of these forms predominate in a given water. Low ph tends to favor free CO2 while high ph tends to favor the carbonate ion. At the near-neutral ph of most natural waters, the bicarbonate ion predominates. Streams, except those that receive significant groundwater inputs, seldom have any free carbon dioxide present in their water. Lakes and ponds (and oceans) on the other hand, may have appreciable free carbon dioxide present, especially in their bottom waters. Measurement of free carbon dioxide: The free carbon dioxide content of natural water may be estimated by a titrimetric procedure using N/44 sodium hydroxide as a titrant and phenolphthalein as an indicator. Free carbon dioxide may also be estimated by calculation using the ph and total alkalinity of the water, and by a nomographic method. Because of the difficulties involved with the titrimetric method, the nomographic or calculation methods are generally used. Hardness is defined as the ability of a water to combine with a standard soap solution. Soaps react with divalent cations to form insoluble precipitates of metal stearates. The major divalent cations of most natural waters are calcium (Ca +2 ) and magnesium (Mg +2 ),

5 so hardness is usually a measure of the sum of these two ions. Other divalent cations such as strontium and ferrous iron may also contribute to hardness. Measurement of hardness: Hardness is usually measured titrimetrically. While standard soap solutions were originally used as a titrant in hardness titrations, these are difficult to prepare and are subject to degradation, so the tetrasodium salt of ethylenediaminetetraacetic acid (EDTA) is generally used as a titrant with a dye such as Calmagite used as an indicator. Like alkalinity, hardness is generally expressed in terms of equivalent calcium carbonate. While alkalinity and hardness are quite different things, they are very strongly positively correlated in most waters, that is; if alkalinity is high, so is hardness. Hardness is sometimes divided into (1) carbonate, and (2) non-carbonate hardness: Carbonate hardness is due to divalent cations associated with the bicarbonate (HCO3 - ) ion. Non-carbonate hardness is associated with divalent cations associated with the sulfate (SO4-2 ), nitrate (NO3 - ), or chloride (Cl - ) ions. Since sulfate is usually the most common of these three in freshwaters, non-carbonate hardness is sometimes referred to as "sulfate hardness". Non-carbonate hardness may be calculated by subtracting alkalinity from total hardness. Turbidity is defined as the ability of a water to scatter light. This ability is caused by the presence of small suspended particles (or some large molecules) in the liquid. Since turbidity is an optical property, it is measured by optical methods; usually employing an electro-optical instrument called a nephelometer (or turbidimeter). Turbidity values are generally expressed in Nephelometric Turbidity Units (NTU). An older method for measuring turbidity was the "Jackson Candle Turbidimeter", and turbidities measured with this instrument were expressed in Jackson Turbidity Units (JTU). Turbidity in natural waters may be due to suspended clay particles, bacteria, or planktonic organisms. Salinity is a water quality parameter generally applied to seawater and brines which have a rather high total dissolved solids content (seawater contains approx. 35,000 mg/l of dissolved solids). Salinity is generally expressed as percent (% = parts per hundred) or parts per thousand ( o /oo). Salinity can be measured hydrometrically (with a hydrometer), conductometrically (with a conductivity meter), refractometrically (with a refractometer), or "argentometrically" by precipitation of the halides (mostly chloride and bromide ions) with a silver nitrate solution. Salinity as such is not measured in freshwater, but in seawater and estuarine waters it is an important and frequently measured parameter. Total filtrable residue (dissolved solids or TDS) is a measure of the total amount of dissolved and colloidial material in a water, that is; the material that will pass through a

6 standard filter. In practice, a water sample is passed through a filter and the filtrate is evaporated to dryness and the residue remaining is weighed (gravimetrically determined). Total nonfiltrable residue (suspended solids) is defined as the amount of solid particulate material suspended in a water. In practice, a sample of known volume is passed through a pre-weighed filter and the filter and the material retained on it is then dried and weighed. The difference in weight represents the weight of the solid material suspended in the water. Such material may be either organic or inorganic, usually a mixture of both. Settlable solids is a measure of the amount of suspended material in a water that will settle-out under the influence of gravity in an unagitated container. It is usually measured with a conical glass container called an Imhoff Cone. Very small suspended particles in water will generally not settle-out because of random thermal agitation by water molecules. Settlable solids measurements are not generally routinely made on natural waters, but are an important consideration in some wastewaters. Settlable solids values are generally reported as a volume (ie. ml/l) Specific Conductance is a measure of a water to conduct an electrical current. Since pure water is a poor conductor of electricity, the degree to which a water conducts an electric current varies directly with the number of dissolved ions present. Conductance can therefore be strongly correlated with total dissolved solids, since in most natural waters, the dissolved solids exist in ionic form. (1) K = Specific Conductance (µmhos/cm 2 ) / Total Dissolved Solids (mg/l) Where K is an empirically determined constant relating the TDS and SC of a specific water. Equation (1) can be rearranged to yield the following equation for calculating the TDS of a water sample: (2) Total Dissolved Solids (mg/l) = Specific Conductance (µmhos/cm 2 ) / K Specific conductance is generally measured with a small electical device (a conductivity meter) that has a pair of electrodes separated by a small, water-filled gap. The conductance of water is expressed in micromhos (µmhos). A mho is a reciprocal ohm, a unit of electrical resistance. Specific conductance is sometimes used as a check on the accuracy of water analyses in which all of the major ions are determined. Figure 1: Idealized relationship between specific conductance and total ionic solids

7 Ionic species: Anions Bicarbonate ion (HCO3 - ) is more common in waters draining from watersheds that contain carbonate rocks such as limestone and dolomite, but carbonates may also be present in other sedimentary rocks such as shales and sandstones. Carbonates are far less common in igneous and metamorphic rocks, although they may occur here as well, especially in the marbles (metamorphic) and the carbonatites (igneous). Bicarbonate ion is the most common anion in most natural freshwaters, and is the chief source of alkalinity in such waters. Bicarbonate ion levels are generally expressed in terms of calcium carbonate equivalent and are determined as part of the alkalinity procedure. Bicarbonate is an important source of carbon for higher aquatic plants and algae, although the aquatic mosses cannot use it as a carbon source. When bicarbonate is used by aquatic plants, it tends to raise water ph because a hydroxide ion is released in the process of extracting carbon dioxide from the bicarbonate ion. (1) HCO > CO2 + OH - Carbonate ion (CO3-2 ) is far less common in natural waters than is the bicarbonate ion, because it exists only at higher ph values (ph > 8.3). Nevertheless, there may be appreciable carbonate ion present in "soda" lakes and some thermal springs. This ion is also generally determined as part of the alkalinity procedure. Hydroxide ion (OH - ). Water ionizes weakly to form hydrogen (hydronium) ions and hydroxide ions: H2O -----> H+ + OH - [or 2H2O -----> H3O + + OH - ] Therefore, (OH - ) is naturally present in all waters, but generally at molar concentrations of less than 10-6 ( 1 part per million). Hydroxide ion concentration is generally expressed as poh, which is the complement of ph. poh is defined as the negative log of the hydroxide ion concentration. Thus, since the ion product constant for water is 1 x 10-14,

8 if the ph of a solution is 4, the poh must be 10. Conversely, if the ph is 10, the poh must be 4. Nitrate ion (NO3 - ) is usually not present in large amounts in unpolluted waters, but may reach high values in some groundwaters, and in waters polluted by agricultural or industrial wastes. The nitrate ion represents the most highly oxidized state of Nitrogen (N +5 ) found in water. Nitrogen may take other forms in aquatic systems as well. The most common of these are the nitrite ion (NO2 - ) and ammonia (NH3 or NH4 + ). Other forms such as hydroxylamine (NH2OH), and organic forms (such as amino acids and proteins) occur as well. Under oxidizing conditions, ammonia is oxidized to nitrite and then to nitrate. Under reducing conditions, nitrate is likely to be reduced to nitrite, and then to nitrogen gas. All of these oxidations and reductions are largely carried-out by specific bacteria in water. Some aquatic bacteria can use nitrate as a terminal oxygen acceptor in respiration (in place of oxygen which more usually plays this role), and thus carry out essentially aerobic respiration in anaerobic environments Groundwaters which contain significant amounts of nitrate ( >10 mg/l) may constitute a health hazard to young children, in which nitrate ingestion may produce a condition called methemoglobinemia which may cause mental retardation or even death. Nitrite ion (NO2 - ) in water is usually due to the oxidation of ammonia or the reduction of nitrate ion (see above). Nitrite levels in unpolluted waters are generally low. Ammonium ion (NH4 + ) is the most common form of ammonia at the ph of most natural waters, although significant ammonia gas (NH3) may exist in water at higher ph. Ammonia is quite soluble in water and reacts with it to form ammonium hydroxide (NH4OH) which dissociates to form ammonium ions and hydroxide ions: NH3 + H2O -----> NH4OH -----> NH4 + + OH - Raising the ph of such a system increases the hydroxide ion concentration which drives the above reaction to the left, favoring ammonia gas rather than the ammonium ion. Sulfate ion (SO4-2 ) may be fairly abundant in seawater and in waters of arid regions where sulfate minerals such as gypsum (CaSO4. 2H2O) may be found in rocks and soil. Gypsum and other sulfate minerals tend to be fairly soluble in water, and thus are usually found only in arid regions. Sulfates may also be produced by the oxidation of sulfide minerals, such as iron pyrite (FeS2), or of hydrogen sulfide gas (H2S). Waters containing acid mine drainage or industrial wastes may contain rather high levels of sulfate. Some bacteria can use sulfate ion as a terminal electron acceptor in respiration as well (see similar discussion under nitrate above). When so used, sulfate is reduced to elemental sulfur or to hydrogen sulfide gas. Chloride ion (Cl - ) is generally not abundant in natural freshwaters, but in regions where rainfall is affected by seaspray (ie. coastal areas) or salty dusts (ie. downwind from

9 deserts), chloride ion levels may be significant. Road salt and oilfield wastewaters may contribute high levels of chloride (and sodium) ions to streams, lakes, and groundwater. Phosphorus ionic species (PO4-3, or HPO4-2, or H2PO4 - ) are important biologically because Phosphorus, while an important macronutrient element, is not very abundant in the Earth's rocks. Usually, phosphate is measured as the orthophosphate ion (PO4-3 ), with other phosphate forms (meta- or polyphosphates) and organic phosphorus compounds being converted to orthophosphate by chemical pretreatment. Phosphorus occurs naturally as the mineral apatite in igneous rocks, and as "phosphate rock", a chemical sedimentary rock. Phosphates are also used widely in cleansing agents and in water and wastewater treatment. In natural waters, phosphates are often precipitated as insoluble metallic phosphates ( eg. iron and aluminum phosphates). There are also frequently precipitated as a component of organic detritus. Orthophosphate forms can be thought of as ionization products of phosphoric acid: H3PO4 (phosphoric acid) -----> H + + H2PO > H + + HPO > H + + PO4-3 Silica ionic and colloidial forms. Silica can exist in the following forms in water: HSiO3 - SiO3-2 Si (OH)6-2 H4SiO4 (hydrosilicic acid) SiO2 (colloidial silica) The form in which silica is found in water is highly dependent on ph. Because of its low solubility at normal (6-8) ph values, silica is generally a minor component of most natural waters, but is an important nutrient for diatoms (important members of the phytoplankton), freshwater sponges, and several other members of aquatic ecosystems. Ionic species: Cations Sodium ion (Na + ) is commonly present in natural fresh waters at fairly low levels. In coastal regions, streams and lakes tend to carry somewhat higher levels of sodium as a result of rainout or dryfall of sodium chloride from atmospherically transported sea spray. Some igneous rocks may yield fair amounts of sodium upon weathering, but most sodium compounds are so soluble that they are rapidly leached from surface rocks and soils. In arid regions, sodium may accumulate in playas and salt lakes, and sodium is usually both abundant and the dominant cation in such waters. Sodium levels may be elevated in the

10 waters of less arid regions downwind from deserts because of wind transport of sodium chloride dusts from playa margins and salt flats. Potassium ion (K + ) usually occurs in natural waters in concentrations one-half to onetenth that of sodium. Even though potassium is fairly abundant in many feldspars, and feldspars are among the most common silicate minerals, the potasium released by the weathering of such minerals appears to be rather quickly re-incorporated into clay minerals which largely resist further weathering. Potassium then, tends to remain scarce in most natural waters. Calcium ion (Ca +2 ) is usually the most common cation in freshwater, but is less abundant than either sodium or magnesium in seawater and in most brackish waters. Calcium carbonate is a common constituent of many sedimentary rocks, but has a quite low solubility in water unless there is considerable carbon dioxide present. Under these conditions, the following reaction occurs: H2O + CO2 + CaCO > Ca HCO3 - [both ions in solution] But this is a reversible reaction and if carbon dioxide is lost from the water, either to the atmosphere or to aquatic plants, the equilibrium is shifted to the left and calcium carbonate is precipitated as marl (in freshwater) or limestone (in seawater). Magnesium ion (Mg +2 ) exceeds calcium in abundance in seawater, but is far less common that calcium in freshwater. Iron ionic species: Fe +2 [ferrous iron] or Fe +3 [ferric iron] are usually not abundant in natural waters because such ions tend to be rapidly precipitated as highly insoluble iron oxides and hydroxides. Ferrous iron is more soluble that ferric, but this ion tends to be rapidly oxidized to the insoluble ferric state in well oxygenated waters. Iron may be fairly abundant in solution in anoxic waters, such as the bottom waters of polluted or eutrophic lakes. All ionic species of iron are more soluble at low ph. Other ions: In waters under special geologic situations, such as those in contact with unusual rocks, those in regions of recent volcanic activity, or those influenced by thermal waters (hot springs), other ionic species such as borate (BO4-3 ; H4BO4 - ), fluoride (F - ), Aluminate (Al2O4-2 ), Strontium (Sr +2 ), or other less common ions may occur in significant quantities. Click Here for a table of representative freshwater analyses. Click Here for a table of Average Seawater analysis.

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