Applications of Environmental Aquatic Chemistry: A Practical Guide - Chapter 4 pptx

4 Behavior of Metal Species in the Natural Environment 4.1 METALS IN WATER 4.1.1 B ACKGROUND A casual glance at the periodic table shows that most of the elements (about three- fourths) are metals or metalloids.* It often happens in environmental literature that little or no distinction is made between metals and metalloids, especially for the metalloids arsenic, selenium, and antimony. To discuss their chemical behavior, the elemental metals may be divided into three general classes: 1. Alkali metals: Li, Na, K, Rb, Cs, and Fr (periodic table group 1A). 2. Alkaline metals: Be, Mg, Ca, Sr, Ba, and Ra (periodic table group 2A). 3. Metals not in the alkali or alkaline groups include the transition metals (all the group B periodic table metals), the metals and metalloids in groups 3A through 6A, and metals whose classifications are not based primarily on periodic table groups, the so-called trace or heavy metals. y Heavy metals in surface waters can be from natural or anthropogenic sources. Currently, anthropogenic inputs of metals exceed natural inputs. Living organisms require trace amounts of some heavy metals, including cobalt, copper, iron, manga- nese, molybdenum, vanadium, strontium, and zinc. Excessive levels of essential metals, however, can be detrimental to the organism. Nonessential heavy metals of * Metalloids are those elements in periodic table groups 3A through 6A that have electrical and chemical properties intermediate between those of metals and nonmetals. They are B, Si, Ge, As, Sb, Te, and Po. For regulatory purposes, it is sometimes useful to group metals and metalloids together, as when they share the same analytical method (e.g., ICP, ion-coupled plasma spectroscopy). y The term ‘‘heavy metals’’ is often encountered in texts and reports, usually meaning metals with atomic numbers equal to or greater than Cu (at. no. 29), especially metals exhibiting toxicity. However, the term heavy metals has no precise definition and its use is inconsistent (Duffus, 2002). Another designation often used is trace metals, generally used for those metals found in the earth’s crust with average concentrations less than 1%. Nearly all the metals are included in this class, the exceptions (with average crustal concentrations greater than 1%) being Na, K, Ca, Mg, Fe, and Al. ß 2007 by Taylor & Francis Group, LLC. particular concern because of their toxicity are cadmium, chromium, mercury, lead, arsenic, and antimony. When metal atoms combine chemically with other metal atoms, the result is a metal substance, either pure elemental metals or alloys. When metal atoms combine with nonmetal atoms, nonmetallic compounds result that range from ionic salts like sodium chloride to volatile, inflammable liquid organometallic compounds like dimethylmercury. Metals and metal-containing compounds in natural waters may be in dissolved, colloidal, or particulate forms, depending on water quality parameters of pH, redox potential, and the presence of other dissolved species such as sulfide or carbona te which can form compounds with metal ions. This section treats how the solubility of metals in water depends on these different factors. Dissolved forms* are . Cations: Ca 2þ ,Fe 2þ ,K þ ,Al 3þ ,Ag þ , etc. . Complexes y : Zn(OH) 2þ 4 , Au(CN) À 2 , Ca(P 2 O 7 ) 2À , PuEDTA, etc. . Organometallics: Hg(CH 3 ) 2 , B(C 2 H 5 ) 3 ,Al(C 2 H 5 ) 3 , etc. Particulate forms are . Mineral sediments . Precipitated oxides, hydroxides, sulfides, carbonates, silicates, etc. . Cations and complexes sorbed to mineral sediments (clays, oxides, hydrox- ides, sulfides, carbonates, silicates, etc.) and organic mat ter. The behavior of metals in natural waters may be described in terms of how they become distributed between dissolved and solid species. Metal species undergo * It is shown in this chapter that cations in water always attract a hydration shell of water molecules because of electrostatic attractions. Although a cation dissolved in water is often written as an elemental ion, e.g., Al 3þ , it actually is a cluster consisting of the metal ion enclosed within progressively larger surrounding shells of water molecules. The water molecules in the innermost shell are the most strongly attracted to the metal cation and the cluster could be written, for example, as Al(H 2 O) 3þ n , where n is the number of water molecules in the inner hydration shell, from 1 to about 6. One simpler way to indicate a dissolved cation is append the suffix ‘‘aq’’ (for aqueous) to it, as in Al 3þ (aq). In this text, and many others, whenever a dissolved cation is written without any indication of its hydration shell, e.g., Al 3þ , the more accurate designation is to be assumed, e.g., Al 3þ (aq). y A complex is a dissolved chemical species formed by the association of a cation with one or more anions or neutral species (such as water) that contain nonbonding electron pairs. The cation has room for one or more electron pairs in its valence shell and the anions or neutral species it connects with (called ligands) have nonbonding electron pairs in their outer shells that can fill the cation electron shell vacancies. Remember that chemical bonds consist of electron pairs positioned between the bonded atoms. A complex differs from a covalent compound in that the ligand brings both electrons of the bonding pair to the bond, while covalent compounds are formed when the connected atoms contribute one electron each. Cation–ligand bonds are called coordinate bonds, and the complexes formed are called coordination compounds. ß 2007 by Taylor & Francis Group, LLC. continuous changes between diss olved, preci pitated, and sorbed -to-sediment forms. The rates of adsorp tion , desorption , and precipita tion proces ses depend on pH, redox potential , water che mistry, and the compo sition of bott om and suspen ded sedim ents. Adsorpt ion of dissolved metal speci es to sediment s remo ves the metal from the water column and stores it in the sediment s, where it is less biological ly avail able. Desorpt ion returns the metal to the water colum n, where it becom es biological ly available again and where water flow may carry the metal to a new location where sorpt ion and preci pitation c an recur. Metals may be desorb ed from sedim ents if the water undergo es incre ases in sali nity, decreas es in redox potential , or decreases in pH. Sorption and desorp tion proces ses are discussed further in Cha pter 5. In the water environment, nonradioactive metals are of greatest environmental and health concern when in diss olved forms, where they are more mobile and more biologically available than are particulate forms, although ingestion and inhalation of particulates containing metals also can be a serious health hazard. Radioactive metals are hazardous because of their ionizing emissions as well as their chemical toxicity and may be harmful in both dissolved and particulate forms, even without metal species entering the body. 4.1.2 MOBILITY OF METALS IN THE WATER ENVIRONMENT Two properties, solubility and the tendency to sorb to soil particles (sorption coefficient) largely govern the mobility of metals in the water environment. Metals can take many forms in environmental soil=water systems and the mobility of each form can depend in a different way on environmental conditions. The important metal forms are 1. Solid elemental metal precipitates. a. These may be particles of colloid size or larger. Colloids remain suspended in water and are mobilized by water movement. Larger particles may settle out and require stronger flows to move them as sediments. 2. Solid metal compounds formed by weathering of minerals and by reactions of dissolved metal cations with water and other dissolved species such as carbonate, fluoride, and sulfate. a. These may be colloid size or larger. 3. Dissolved metal cations. 4. Dissolved metal compounds, such as carbonate and hydroxy complexes. 5. Metal species sorbed to solid soil and sediment surfaces. a. Sorption proces ses may be reversible to some degree, resulting in a retardation of dissolved metal movement relative to water flow, or irreversible, resulting in immobilization of metal species, except for erosion mechanisms. b. Dissolved or colloidal solids may become sorbed to solid surfaces. ß 2007 by Taylor & Francis Group, LLC. 4.1.3 GENERAL BEHAVIOR OF DISSOLVED METALS IN WATER It can be mis leading to think in term s of the solub ility of elemental metals. For exa mple, to say that ‘‘iron is more solub le und er reduci ng condit ions than under ox idizing condition s, ’’ d oes not c all atte ntion to the fact that it is not elem ental iron that is more solub le; it is the iron compo unds that can be formed which may (or may no t) be more solub le under reducing condit ions. Reaction s of metal catio ns with water (hydro lysis) are the usual criteria for asses sing whether a metal is soluble or insolubl e under certain redox and pH con ditions. When a metal such as iron is said to be insol uble under oxidi zing con ditions and soluble under reduci ng condition s, what actual ly is meant is that the co mpounds formed by reaction of the metal cation with water und er oxidizing con ditions are insolubl e; u nder reducing cond itions iron does not hydrol yze in water and can rema in as a dissolved cation (wi th a hydrat ion shell). The same is true for pH con ditions; met als tend to be less solub le at h igh pH because their cati ons often react in high pH water to form low-sol ubility hydroxides and o xides. At low pH, wher e hy droxide concent rations are low (see Cha pter 3), they may rema in as solub le hy drated cations. 4.1 .3.1 Hydrol ysis Rea ctions The simp lest form of a d issolved meta l is an elemental catio n, such as Fe 3þ or Zn 2þ . How ever, elemental cati ons cannot exist as such in water solut ions. Any charged speci es in solution will inte ract with other charged or polar species because of elect rical forces. Bec ause water molecules are polar (C hapter 2), metal cati ons alw ays attr act a multilayere d hydrat ion shell of water molecules by electrost atic attr action of the negative end of the water mol ecules (the oxygen end, see Cha pter 2) to the positive c harge of the cation, as descri bed by Equati on 4.1 and illu strated in Figure 4.1. M þn ! H 2 O M(H 2 O) þn x (4:1) RULES OF THUMB 1. Dissolved forms of metals move with surface water and groundwater flows. 2. Metals in particulate form can be transported with sediments by wind and in moving water. 3. Both dissolved and particulate forms of metals may sorb to organic soil solids, where they can be immobilized or carried along with eroding soils. 4. Metals pose the greatest environmental risks when particulate metals encounter environmental conditions that increase their solubility. ß 2007 by Taylor & Francis Group, LLC. where M is a metal cation n is the number of positive charges on the cation x is the maximum number of water molecules in the innermost hydration shell x is 6 for most cations Depending on the strength of the electrostatic attraction between the cation and the water molecules, the water molecules closest to the metal cation (in the innermost hydration shell) may bond as a ligand, forming a metal–water complex (see second footnote o n page 110). The strength of the electrostatic attraction depends on the magnitude of the cation charge, the cation radius, and, to a lesser extent, the electro- negativity of the metal. The strongest bonds are formed with cations having the smallest radius, greatest positive charge, and electronegativity greater than 1.8 (Wulfsburg, 1987). 4.1.3.2 Hydrated Metals as Acids Hydrated met al ions can behave as acids by releasing protons (H þ ) from their water ligands that then become attached to the surrounding free H 2 O molecules, forming acidic hydrated protons, H 3 O þ and H 5 O þ 2 (see Secti on 4.2.1). The stronger the bond between the water ligand and the metal cation, the more readily a proton is released to surrounding water mol ecules and the more acidic is the hydrated metal cation. The process can continue stepwise up to n times to make a neutral metal hydroxide.* H H H H H H O O H H O H H H H + O H H O H H H H M n + H O O O O O H H H FIGURE 4.1 Water molecules form a hydration shell around dissolved metal cations. Mole- cules 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. In this way, the hydrated metal behaves as an acid. Eventually, the metal may precipitate as a hydroxide compound of low solubility. * For simplicity, hydrated protons will be designated only by the most common form, H 3 O þ . Because solvent water molecules are normally not included when balancing a chemical reaction, the stoichiom- etry of acid-base reactions remains unchanged, regardless of how many water molecules are shown attached to an acidic proton, i.e., whether the hydrated proton is designated by H þ , H 3 O þ , H 5 O þ 2 , etc. ß 2007 by Taylor & Francis Group, LLC. M(H 2 O) þn 6 þ H 2 O $ M(H 2 O) 5 OH þ(nÀ1) þ H 3 O þ (4:2) M(H 2 O) 5 OH þ(nÀ1) þ H 2 O $ M(H 2 O) 4 (OH) 2 þ(nÀ2) þ H 3 O þ (4:3) For example, with Fe 3þ , it takes three proton transfer steps to form neutral ferric hydroxide: Fe(H 2 O) 3þ 3 þ H 2 O $ Fe(H 2 O) 2 OH 2þ þ H 3 O þ (4:4) Fe(H 2 O) 2 OH 2þ þ H 2 O $ Fe(H 2 O)(OH) þ 2 þ H 3 O þ (4:5) Fe(H 2 O)(OH) þ 2 þ H 2 O $ Fe(OH) 3 (s) þ H 3 O þ (4:6) Adding Equations 4.4 through 4.6 gives the overall reaction: Fe(H 2 O) 3þ 3 þ 3H 2 O $ Fe(OH) 3 (s) þ 3H 3 O þ (4:7) With each step, the hydrated metal is progressively deprotonated, forming polyhydroxides and becoming increasingly insoluble. At the same time, the solution becomes increasingly acidi c due to the formation of more H 3 O þ . Eventually, the metal may precipitate as a low-solubility hydroxide. The degree of acidity induced by metal hydration is the greatest for cations having the greatest electronegativity, which are those of high charge and small size. All metal cations with a charge of þ3 or more are moderately strong acids. This process is one source of acidic water draining from mines. RULES OF THUMB 1. Only polyvalent cations (e.g., Fe 3þ ,Zn 2þ ,Mn 2þ ,andCr 3þ ) have large enough charges to attract water molecules strongly enough to act as acids, by causing the release of H þ from water molecules in the hydration sphere. Monovalent cations, such as Na þ , do not act as acids at all. 2. The interactions of metal cations with water, Equations 4.2 through 4.7, cause the solubility of met al species in water to be dependent on pH and redox potential. a. Low pH (high H 3 O þ concentration and high acidity) increases metal solubility by shifting the equilibria of Equations 4.2 through 4.7 to the left, decreasing the formation of less soluble metal polyhydroxides. b. High pH (low H 3 O þ concentration and low acidity) decreases metal solubility by shifting the equilibria of Equations 4.2 through 4.7 to the right, increasing the formation of less soluble metal polyhydroxides. ß 2007 by Taylor & Francis Group, LLC. 4.1.4 INFLUENCE OF P H ON THE SOLUBILITY OF METALS All the reactions (Equations 4.2 through 4.7) are revers ible, with H 3 O þ on the right side. This means that the equilibri a of these react ions shift to the left if the concen- tration of H 3 O þ is incre ased (by adding more acid) and to the right if it is de creased (by adding a base). Thu s, the formati on of metal hydroxi des by hyd ration of met al cations is sensiti ve to the solut ion pH. Conside ring the o verall reaction, Equ ation 4.7, we see that low ering the pH (incr easing the con centratio n of H 3 O þ ) shifts the equilibri um of Equation 4.7 to the left, tending to diss olve any solid metal hydroxide that has preci pitated. Raising the pH (incr easing the concent ration of OH À ) con- sumes H 3 O þ and shifts the equil ibrium of Equation 4.7 to the right, precipita ting more insolubl e met al hydroxide. Thus, on e may say that the met al becomes more soluble at lower p H and less solub le at higher pH, even though what ac tually o ccurs is that the hydrated met al forms less solub le hydroxi de at higher pH. However , if the pH is rais ed too high, preci pitated met al hydroxi des can redis- solve (see Figure 4.2). At high pH values , a metal hydroxi de may form complexes with OH À anions to becom e a negative ly charged ion having increased solub ility. For example, preci pitated Fe (OH) 3 can react wi th OH À anions as follows: Fe(O H) 3 þ OH À $ Fe(OH) À 4 ( 4: 8) Fe(OH ) À 4 þ OH À $ Fe (OH) 2 À 5 ( 4: 9) Nega tively c harged polyhy droxi de anions are more solub le because their ionic charge attracts them stro ngly to polar water molecules . As shown in Figure 4.2, c. Low redox potential s (reduc ing condition s, low-to -zero diss olved oxygen (DO) level s, where electron donors are more comm on than elect ron accept ors) increase the solub ility of many met als (see Secti on 4. 1.5) by promoti ng low er oxidat ion numbe rs for met al cati ons (lower positive charge, e.g., Fe 2þ rathe r than Fe 3þ ). Fo r cati ons with lower positive charge, the equilibria of Equ ations 4 .2 throu gh 4.7 are maintai ned more stro ngly to the left, resulting in less formation of low-solubility polyhydroxides. d. High redox potentials (oxidizing conditions, high DO levels, where electron acceptors are more common than electron donors) decrease the solubility of many metals by promoting higher oxidation num- bers for metal cations (higher positive charge, e.g., Fe 3þ rather than Fe 2þ ). For cations with higher positive charge, the equilibria of Equations 4.2 through 4.7 are maintained more strongly to the right, resulting in greater formation of low-solubility polyhydroxides. 3. The presence of dissolved species such as sulfide or carbonate, which form low-solubility compounds with metal cations, can largely negate the above generalizations by competing with hydroxide formation. RULES OF THUMB (Continued) ß 2007 by Taylor & Francis Group, LLC. 0.0001 0.001 0.01 0.1 1 10 100 3456789101112 pH Fe(OH)3 Al(OH)3 Ni(OH) 2 Fe(OH) 2 Cd(OH) 2 Cu(OH) 2 Pb(OH) 2 Cd(OH) 2 AgOH Ni(OH) 2 Zn(OH) 2 Cu(OH) 2 Cr(OH) 3 Al(OH) 3 Fe(OH)3 Cr(OH)3 Fe(OH) 2 FIGURE 4.2 Theoretical solubilities of some metal hydroxides versus pH. ß 2007 by Taylor & Francis Group, LLC. the high value of pH, wher e solubility begins to increase again, varie s from met al to metal. Al kaline wat er provi des a buffer agains t pH changes . In a lkaline water, the tenden cy of metals to make wat er acidic is dimini shed by reactions like Equ ations 4.11 through 4.13. EXAMPLE 1 EFFECT OF D ISSOLVED M ETAL ON A LKALINITY 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 for CaCO 3 . Is this laboratory measurement of alkalinity likely to accurately represent the groundwater alkalinity? Answer : Soluble inorganic iron is in the ferrous form, Fe 2þ . Because of its small charge, loss of protons from the hydration sphere is not a signi ficant process for hydrated ferrous iron Fe 2 þ , denoted in Equation 4.10 by Fe(H 2 O) 2 þ 6 . However, when a groundwater sample is exposed to air, oxygen (an electron acceptor) dissolving from the atmosphere can oxidize Fe 2þ to the ferric form, Fe 3þ . This process is often enhanced by aerobic iron bacteria (see reaction step 2 in Figure 4.3). Depending on the pH, hydrated Fe 3þ can lose protons from its hydration sphere to any bases present, including water molecules and hydroxyl ions (OH À ), forming ferric hydroxide species and making the solution more acidic. The acidic behavior of hydrated Fe 3þ occurs to a greater extent at higher pH. Equation 4.10 represents the overall oxidation reaction that converts dissolved ferrous iron to precipitated ferric hydroxide: 4Fe(H 2 O) 2þ 6 þ O 2 $ 4Fe(OH) 3 (s) þ 14H 2 O þ 8H þ (4:10) Fe(OH) 3 is a yellow to red-brown precipitate often seen on rocks and sediments in surface waters with high iron concentrations. The molar concentration of H þ formed by Equation 4.10 can be up to 2 times the Fe 2þ molar concentration, depending on the final pH. Each H þ released will neutralize a molecule of base, consuming some alkalinity, by reactions such as H þ þ OH À $ H 2 O(4:11) H þ þ HCO À 3 $ H 2 CO 3 (4:12) H þ þ CO 2À 3 $ HCO À 3 (4:13) We will assume a worst-case scenario with respect to affecting the alkalinity, where the pH is high enough that the equilibrium of Equation 4.10 goes essentially to completion to the right side as additional oxygen dissolves from the atmosphere. The atomic weights of hydrogenandironare1and56g=mol, respectively. If the equilibrium of Equation 4.10 is completely to the right, 1 mole (56 g) of Fe 2þ will produce 2 moles of H þ (2 g). At th e time of sampling, the concentration of dissolved Fe 2þ (as Fe(H 2 O) 2þ 6 ) was about 20 mg=L and all is eventually oxidized to Fe 3þ (as Fe(OH) 3 ).Themolarconcentrationofironis 0:020 g=L 56 g=mol ¼ 0:00036 mol=Lor0:36 mmol=L ß 2007 by Taylor & Francis Group, LLC. By Equation 4.10, the moles of H þ produced are two times the moles of iron: Moles of H þ ¼ 2 Â 0: 36 mmol=L ¼ 0: 72 mmol =L Grams of H þ ¼ 0: 72 mmol =L Â 1mg=mmol ¼ 0: 72 mg=L We must now determine what effect this quantity of H þ will have on the alkalinity. Alkalinity is measured in terms of a comparable quantity of CaCO 3 . The molecular weight (MW) of CaCO 3 is 100 g=mol, and it dissolves to form the doubly charged ions Ca 2þ and CO 2À 3 . Alkalinity is a property of the CO 2À 3 anion, which consumes acidity by accepting 2 H þ cations: 2H þ þ CO 2À 3 $ H 2 CO 3 (4: 14) Therefore, 0.72 mmol=LofH þ will react with 0.72=2 ¼ 0.36 mmol=Lof CO 2À 3 , and 0.36 mmol=L of CaCO 3 is required as a source of the CO 2À 3 .* From the definition of alkalinity, the change in alkalinity is equal to the change in concentration of CaCO 3 ,in milligram per liter. 0: 36 mmol=L of CaCO 3 ¼ 0: 36 mmol= L Â 100 mg 1 mmol ¼ 36 mg =L ¼ change in alkalinity Groundwater alkalinity at time of sampling ¼ Lab: measured alk : þ alk: lost by Equation 4: 10 The maximum possible value for the original alkalinity of the groundwater before exposure to air was 150 mg=L þ 36 mg=L ¼ 186 mg =L as CaCO 3 The laboratory alkalinity measurement was lower than the actual groundwater alkali- nity, with a maximum error of about 21%. In the above example, the pH was assumed high enough to maintain the equilibrium of Equation 4.10 entirely to the right. Under these conditions, essentially all the H þ added to the solution will react by Equation 4.13 to form H 2 CO 3 . Thus, there would be no signi ficant net change in aqueous H þ (as H 3 O þ ) and little corresponding change in pH. In practice, the changes in concentrations brought about by Equations 4.10 through 4.13 will never cause the equilibria of the reactions to go completely to the right. Thus, the H þ added will never react completely with the alkalinity and there will always be at least a small net increase in the H 3 O þ concentration, accompanied by a corresponding small decrease in pH. This example illustrates the pH buffering effect of alkalinity. The addition of H þ to the solution by hydration of metal ions, as in Equation 4.10, will not change the pH greatly as long as some alkalinity remains, because the added H þ is taken up by carbonate species in the water. This is also true for the addition of H þ from other sources, such as mineral and organic acids. * Equation 4.14 can be obtained by adding Equations 4.12 and 4.13. ß 2007 by Taylor & Francis Group, LLC. [...]... rock or earth material with a negative ABP is likely to have acidic leachate If the ABP is À5 or more negative, the earth material has an acid-neutralizing deficiency of 5.0 tons CaCO3=1000 tons material, and may be considered a potentially hazardous waste RULES OF THUMB 1 If the ABP is positive, leachate from the sample is likely to be basic 2 If the ABP is negative, leachate is likely to be acidic 3... 4. 3.1 TREATMENT OF TRACE METALS IN URBAN STORMWATER RUNOFF Stormwater runoff carries solid and dissolved forms of metals as well as other chemical pollutants, including soil sediments and various kinds of debris, such as paper, plastic, garbage, leaf, and plant litter Because a major constraint on stormwater treatment systems is that they provide passive or near-passive treatment, the biggest challenge... waste materials If precipitation passing through the soil cap were made more acidic because of iron pyrite in the soil, it might mobilize metals or other contaminants in the wastes and create a new hazard The ABP is calculated by ABP ¼ (alkalinity) À (31:25) (wt% pyritic sulfur) (4: 19) where the ABP units are tons of acid-neutralizing substances (as CaCO3-equivalents) per 1000 tons of solid material Any... the ABP is À5, or more negative, the earth material is likely to be acidic enough to be defined as a potentially hazardous waste 4. 5 CASE STUDY 3 4. 5.1 IDENTIFYING METAL LOSS AND GAIN MECHANISMS IN A STREAM This study is modeled after parts of the report by Balistrieri et al (1995) Background: Measurements of pH and metal concentrations in a stream receiving acid rock drainage indicated that the stream... precipitate a fraction of the dissolved metals as hydroxides and oxides However, at many locations, incidental metal removal may not be adequate, if nonpoint source contamination by metals is an important secondary (and sometimes primary) cause of water quality impairment Ideally, a stormwater vault collects storm runoff, retains the sediments, and immobilizes the pollutants so that they do not enter surface... concentrations This is true for any conservative constituent In this example, flow downstream of the confluence of stream and the mine drainage serves to mix their chloride concentrations and each location downstream represents a particular mixing ratio, where the chloride can only change by evaporation of stream water or by new water inputs Figure 4. 4 is a line diagram of a stream receiving mine drainage Sampling... stormwater treatment is the removal of dissolved pollutants by means that require minimal or no operator control or external power sources The general methods available for water treatment are control of the redox potential and pH, addition of chemical reactants to aid precipitation, and length of detention time in vaults or basins In a passive or near-passive treatment system The redox potential... precipitated and desorption of metals sorbed to sediments Encourage biodegradation of dissolved and solid organic substances, such as pesticides and petroleum products Unfortunately, no single set of operating conditions can realize all these goals simultaneously For example, maintaining adequate oxidizing conditions (greater than about 2 ppm of DO) in a stormwater vault encourages biodegradation of organic matter... appropriate for passive treatment because it normally requires chemical additions The pH in a stormwater treatment system is usually determined by the prevailing environmental conditions, and normally is in the range of 6–9 Addition of chemical reactants other than oxygen is an approach that requires regular maintenance and, therefore, is not an option in passive treatment systems A passive treatment... production of acid rock drainage can be a rapid, self-propagating, cyclic process that is accelerated by low pH and the presence of iron-oxidizing bacteria The process will continue as long as oxygen, pyrite, and water are present ß 2007 by Taylor & Francis Group, LLC 4. 4.1.2 Non-iron Metal Sulfides Do Not Generate Acidity The abiotic oxidation by dissolved O2 of non-iron metal sulfides does not generate significant . of solid Fe(OH) 3 and is available to react via Equation 4. 17. 4. 4.1.1 Summary of Acid Formation in Acid Rock Drainage The steps of acid formation in acid rock drainage are summarized below and illustrated. and inhalation of particulates containing metals also can be a serious health hazard. Radioactive metals are hazardous because of their ionizing emissions as well as their chemical toxicity and. this chapter that cations in water always attract a hydration shell of water molecules because of electrostatic attractions. Although a cation dissolved in water is often written as an elemental ion,