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  In addition to providing updated material in the rapidlydeveloping area of environmental chemistry, this edition emphasizes several major concepts that are proving essential to the practice of environmental chemistry at thebeginning of the new millennium. These include the concept of the anthrosphere as a distinct sphere of the environment and the practice of industrial ecology,sometimes known as “green chemistry” as it applies to chemical science.


2.1 The Anthrosphere 2.2 Technology and the Anthrosphere 2.3 Infrastructure 2.4 Dwellings 2.5 Transportation 2.6 Communications 2.7 Food and Agriculture 2.8 Manufacturing 2.9 Effects of the Anthrosphere on Earth 2.10 Integration of the Anthrosphere into the Total Environment 2.11 The Anthrosphere and Industrial Ecology 2.12 Environmental Chemistry 3.1 Water Quality and Quantity 3.2 The Properties of Water, a Unique Substance 3.3 The Characteristics of Bodies of Water 3.4 Aquatic Life 3.5 Introduction to Aquatic Chemistry 3.6 Gases in Water © 2001 CRC Press LLC 3.7 Water Acidity and Carbon Dioxide in Water 3.8 Alkalinity 3.9 Calcium and Other Metals in Water 3.10 Complexation and Chelation 3.11 Bonding and Structure of Metal Complexes 3.12 Calculations of Species Concentrations 3.13 Complexation by Deprotonated Ligands 3.14 Complexation by Protonated Ligands 3.15 Solubilization of Lead Ion from Solids by NTA 3.16 Polyphosphates in Water 3.17 Complexation by Humic Substances 3.18 Complexation and Redox Processes 4.1 The Significance of Oxidation-Reduction Phenomena 4.2 The Electron and Redox Reactions 4.3 Electron Activity and pE 4.4 The Nernst Equation 4 5 Reaction Tendency: Whole Reaction from Half-Reactions 4.6 The Nernst Equation and Chemical Equilibrium 4.8 Reactions in Terms of One Electron-Mole 4.9 The Limits of pE in Water 4.10 pE Values in Natural Water Systems4.11 pE-pH Diagrams 4.12 Corrosion 5.1 Chemical Interactions Involving Solids, Gases, and Water 5.2 Importance and Formation of Sediments 5.3 Solubilities 5.4 Colloidal Particles in Water 5.5 The Colloidal Properties of Clays 5.6 Aggregation of Particles 5.7 Surface Sorption by Solids 5.8 Ion Exchange with Bottom Sediments 5.9 Sorption of Gases—Gases in Interstitial Water 6.1 Aquatic Biochemical Processes 6.2 Algae 6.3 Fungi 6.4 Protozoa 6.5 Bacteria 6.6 The Prokaryotic Bacterial Cell 6.7 Kinetics of Bacterial Growth 6.8 Bacterial Metabolism 6.9 Microbial Transformations of Carbon 6.10 Biodegradation of Organic Matter 6.11 Microbial Transformations of Nitrogen 6.12 Microbial Transformations of Phosphorus and Sulfur 6.13 Microbial Transformations of Halogens and Organohalides © 2001 CRC Press LLC 6.14 Microbial Transformations of Metals and Metalloids 6.15 Microbial Corrosion 7.1 Nature and Types of Water Pollutants 7.2 Elemental Pollutants 7.3 Heavy Metals 7.4 Metalloids 7.5 Organically Bound Metals and Metalloids 7.6 Inorganic Species 7.7 Algal Nutrients and Eutrophication 7.8 Acidity, Alkalinity, and Salinity 7.9 Oxygen, Oxidants, and Reductants 7.10 Organic Pollutants 7.11 Pesticides in Water 7.12 Polychlorinated Biphenyls 7.13 Radionuclides in the Aquatic Environment 8.1 Water Treatment and Water Use 8.2 Municipal Water Treatment 8.3 Treatment of Water for Industrial Use 8.4 Sewage Treatment 8.5 Industrial Wastewater Treatment 8.6 Removal of Solids 8.7 Removal of Calcium and Other Metals 8.8 Removal of Dissolved Organics 8.9 Removal of Dissolved Inorganics 8.10 Sludge 8.11 Water Disinfection 8.12 Natural Water Purification Processes 8.13 Water Reuse and Recycling


9.1 The Atmosphere and Atmospheric Chemistry 9.2 Importance of the Atmosphere 9.3 Physical Characteristics of the Atmosphere 9.4 Energy Transfer in the Atmosphere 9.5 Atmospheric Mass Transfer, Meteorology, and Weather 9.6 Inversions and Air Pollution 9.7 Global Climate and Microclimate 9.9 Acid-Base Reactions in the Atmosphere 9.10 Reactions of Atmospheric Oxygen 9.11 Reactions of Atmospheric Nitrogen 9.12 Atmospheric Carbon Dioxide 9.13 Atmospheric Water 10.1 Particles in the Atmosphere 10.2 Physical Behavior of Particles in the Atmosphere 10.3 Physical Processes for Particle Formation © 2001 CRC Press LLC 10.4 Chemical Processes for Particle Formation 10.5 The Composition of Inorganic Particles 10.6 Toxic Metals 10.7 Radioactive Particles 10.8 The Composition of Organic Particles 10.9 Effects of Particles 10.10 Water as Particulate Matter 10.11 Control of Particulate Emissions 11.1 Inorganic Pollutant Gases 11.2 Production and Control of Carbon Monoxide 11.3 Fate of Atmospheric CO 11.4 Sulfur Dioxide Sources and the Sulfur Cycle 11.5 Sulfur Dioxide Reactions in the Atmosphere 11.6 Nitrogen Oxides in the Atmosphere 11.7 Acid Rain 11.8 Ammonia in the Atmosphere 11.9 Fluorine, Chlorine, and Their Gaseous Compounds 11.10 Hydrogen Sulfide, Carbonyl Sulfide, and Carbon Disulfide 12.1 Organic Compounds in the Atmosphere 12.2 Organic Compounds from Natural Sources 12.3 Pollutant Hydrocarbons 12.4 Aryl Hydrocarbons 12.5 Aldehydes and Ketones 12.6 Miscellaneous Oxygen-Containing Compounds 12.7 Organohalide Compounds 12.8 Organosulfur Compounds 12.9 Organonitrogen Compounds 13.1 Introduction 13.2 Smog-Forming Automotive Emissions 13.3 Smog-Forming Reactions of Organic Compounds in the Atmosphere 13.4 Overview of Smog Formation 13.5 Mechanisms of Smog Formation 13.6 Reactivity of Hydrocarbons 13.7 Inorganic Products from Smog 13.8 Effects of Smog 14.1 Anthropogenic Change in the Atmosphere 14.2 Greenhouse Gases and Global Warming 14.3 Acid Rain 14.4 Ozone Layer Destruction 14.5 Photochemical Smog 14.6 Nuclear Winter 14.7 What Is to Be Done? © 2001 CRC Press LLC 15.1 Introduction 15.2 The Nature of Solids in the Geosphere 15.3 Physical Form of the Geosphere 15.4 Internal Processes 15.5 Surface Processes 15.6 Sediments 15.7 Clays 15.8 Geochemistry 15.9 Groundwater in the Geosphere 15.10 Environmental Aspects of the Geosphere 15.11 Earthquakes 15.12 Volcanoes 15.13 Surface Earth Movement 15.14 Stream and River Phenomena 15.15 Phenomena at the Land/Ocean Interface 15.16 Phenomena at the Land/Atmosphere Interface 15.17 Effects of Ice 15.18 Effects of Human Activities 15.20 Water Pollution and the Geosphere 15.21 Waste Disposal and the Geosphere 16.1 Soil and Agriculture 16.2 Nature and Composition of Soil 16.3 Acid-Base and Ion Exchange Reactions in Soils 16.4 Macronutrients in Soil 16.5 Nitrogen, Phosphorus, and Potassium in Soil 16.6 Micronutrients in Soil 16.7 Fertilizers 16.8 Wastes and Pollutants in Soil 16.9 Soil Loss and Degradation 16.10 Genetic Engineering and Agriculture 16.11 Agriculture and Health 17.1 Introduction and History 17.2 Industrial Ecosystems 17.3 The Five Major Components of an Industrial Ecosystem 17.4 Industrial Metabolism 17.5 Levels of Materials Utilization 17.6 Links to Other Environmental Spheres 17.7 Consideration of Environmental Impacts in Industrial Ecology 17.8 Three Key Attributes: Energy, Materials, Diversity 17.9 Life Cycles: Expanding and Closing the Materials Loop 17.10 Life-Cycle Assessment 17.11 Consumable, Recyclable, and Service (Durable) Products 17.12 Design for Environment 17.13 Overview of an Integrated Industrial Ecosystem 17.14 The Kalundborg Example 17.15 Societal Factors and the Environmental Ethic © 2001 CRC Press LLC


18.1 Introduction 18.2 Minerals in the Geosphere 18.3 Extraction and Mining 18.4 Metals 18.5 Metal Resources and Industrial Ecology 18.6 Nonmetal Mineral Resources 18.7 Phosphates 18.8 Sulfur 18.9 Wood—A Major Renewable Resource 18.10 The Energy Problem 18.11 World Energy Resources 18.12 Energy Conservation 18.13 Energy Conversion Processes 18.13 Petroleum and Natural Gas 18.14 Coal 18.15 Nuclear Fission Power 18.16 Nuclear Fusion Power 18.17 Geothermal Energy 18.18 The Sun: An Ideal Energy Source 18.19 Energy from Biomass 18.20 Future Energy Sources 18.21 Extending Resources through the Practice of Industrial Ecology 19.1 Introduction 19.2 Classification of Hazardous Substances and Wastes? 19.3 Sources of Wastes 19.4 Flammable and Combustible Substances 19.5 Reactive Substances 19.6 Corrosive Substances 19.7 Toxic Substances 19.8 Physical Forms and Segregation of Wastes 19.9 Environmental Chemistry of Hazardous Wastes 19.10 Physical and Chemical Properties of Hazardous Wastes 19.11 Transport, Effects, and Fates of Hazardous Wastes 19.12 Hazardous Wastes and the Anthrosphere 19.13 Hazardous Wastes in the Geosphere 19.14 Hazardous Wastes in the Hydrosphere 19.15 Hazardous Wastes in the Atmosphere 19.16 Hazardous Wastes in the Biosphere


20.1 Introduction 20.2 Waste Reduction and Minimization 20.3 Recycling © 2001 CRC Press LLC 20.4 Physical Methods of Waste Treatment 20.5 Chemical Treatment: An Overview 20.6 Photolytic Reactions 20.7 Thermal Treatment Methods 20.8 Biodegradation of Wastes 20.9 Land Treatment and Composting 20.10 Preparation of Wastes for Disposal 20.11 Ultimate Disposal of Wastes 20.12 Leachate and Gas Emissions 20.13 In-Situ Treatment 21.1 Biochemistry 21.2 Biochemistry and the Cell 21.3 Proteins 21.4 Carbohydrates 21.5 Lipids 21.6 Enzymes 21.7 Nucleic Acids 21.8 Recombinant DNA and Genetic Engineering 21.9 Metabolic Processes 21.10 Metabolism of Xenobiotic Compounds 22.1 Introduction to Toxicology and Toxicological Chemistry 22.2 Dose-Response Relationships 22.3 Relative Toxicities 22.4 Reversibility and Sensitivity 22.5 Xenobiotic and Endogenous Substances 22.6 Toxicological Chemistry 22.7 Kinetic Phase and Dynamic Phase 22.8 Teratogenesis, Mutagenesis, Carcinogenesis, and Effects on the Immune and Reproductive Systems 22.9 Health Hazards 23.1 Introduction 23.2 Toxic Elements and Elemental Forms 23.3 Toxic Inorganic Compounds 23.4 Toxicology of Organic Compounds 24.1 General Aspects of Environmental Chemical Analysis 24.2 Classical Methods 24.3 Spectrophotometric Methods 24.4 Electrochemical Methods of Analysis 24.5 Chromatography 24.6 Mass Spectrometry © 2001 CRC Press LLC 24.7 Analysis of Water Samples 24.8 Automated Water Analyses 25.1 Introduction 25.2 Sample Digestions 25.3 Analyte Isolation for Organics Analysis 25.4 Sample Cleanups 25.5 Immunoassay Screening of Wastes 25.6 Determination of Chelating Agents 25.7 Toxicity Characteristic Leaching Procedures 26.1 Atmospheric Monitoring 26.2 Sampling 26.3 Methods of Analysis 26.4 Determination of Sulfur Dioxide 26.5 Nitrogen Oxides 26.6 Analysis of Oxidants 26.7 Analysis of Carbon Monoxide 26.8 Determination of Hydrocarbons and Organics 26.9 Analysis of Particulate Matter 26.10 Direct Spectrophotometric Analysis of Gaseous Air Pollutants 27.1 Introduction 27.2 Indicators of Exposure to Xenobiotics 27.3 Determination of Metals 27.4 Determination of Nonmetals and Inorganic Compounds 27.5 Determination of Parent Organic Compounds 27.6 Measurement of Phase 1 and Phase 2 Reaction Products 27.7 Determination of Adducts 27.8 The Promise of Immunological Methods 28.1 Introduction 28.2 Elements 28.3 Chemical Bonding 28.4 Chemical Reactions and Equations 28.5 Solutions 29.1 Organic Chemistry 29.2 Hydrocarbons 29.3 Organic Functional Groups and Classes of Organic Compounds 29.4 Synthetic Polymers © 2001 CRC Press LLC


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  The part of the geosphere that is directlyinvolved with environmental processes through contact with the atmosphere, the To a large extent, the strong interactions among living organisms and the various spheres of the abiotic environment are best described by cycles of matter thatinvolve biological, chemical, and geological processes and phenomena. Environmental chemistry may be defined as the study of the sources, reactions,transport, effects, and fates of chemical species in water, soil, air, and living .environments, and the effects of technology thereon Environmental chemistry is not a new discipline.


  It occurs in all spheres of the environment—in the oceans as a vast reservoirof saltwater, on land as surface water in lakes and rivers, underground as groundwater, in the atmosphere as water vapor, in the polar icecaps as solid ice, andin many segments of the anthrosphere such as in boilers or municipal water distribution systems. Geological science uses chemistry in the form of geochemistry to explain the nature and behavior of geological materials, physics toexplain their mechanical behavior, and biology to explain the mutual interactions 3 between the geosphere and the biosphere.


  Virtually all of the biosphere is containedby the geosphere and hydrosphere in the very thin layer where these environmental spheres interface with the atmosphere. It is believed that organisms were responsible forconverting Earth’s original reducing atmosphere to an oxygen-rich one, a process that also resulted in the formation of massive deposits of oxidized minerals, such as 2 3 2 atmosphere, thus preventing runaway greenhouse warming of Earth’s surface.

2 In so doing, plants and algae function as autotrophic organisms, those that utilize

solar or chemical energy to fix elements from simple, nonliving inorganic material into complex life molecules that compose living organisms. The opposite process,biodegradation, breaks down biomass either in the presence of oxygen (aerobic respiration), {CH O} + O (g) CO + H O (1.4.2) → 2 2 2 2 or absence of oxygen (anaerobic respiration):2{CH O} CO (g) + CH (g) (1.4.3) → 2 2

4 Both aerobic and anaerobic biodegradation get rid of biomass and return carbon

  Bioengineering of organisms withrecombinant DNA technology and older techniques of selection and hybridization are causing great changes in the characteristics of organisms and promise to result ineven more striking alterations in the future. The habitats oceanic marine environment is characterized by saltwater and may be divided broadly into the shallow waters of the continental shelf composing the neritic zoneand the deeper waters of the ocean that constitute the oceanic region.


  Therefore, the following importantpoints related to electromagnetic radiation should be noted: 8 λ and frequency ( , Greek “nu”) as illustrated below: νAmplitude Shorter wavelength.higher frequency Wavelength = c ν λ 1 where is in units of cycles per second (s- , a unit called the hertz, Hz) ν and is in meters (m).λ quanta or photons. This time period has seen a transition from the almost exclusive use of energy captured by photosynthesis and utilized as biomass (food to provide musclepower, wood for heat) to the use of fossil fuel petroleum, natural gas, and coal for about 90 percent, and nuclear energy for about 5 percent, of all energy employedcommercially.


  As part of the carbon cycle, atmospheric carbon in CO is fixed as 2 biomass; as part of the nitrogen cycle, atmospheric N is fixed in organic matter. Many are exogenic cycles in which the element in question spends part of the cycle in the atmosphere—O for oxygen, N for nitrogen, CO for carbon.

2 Carbon Cycle

  Some of the carbon is 3 2 large amount of carbon is present in minerals, particularly calcium and magnesium carbonates such as CaCO . An important aspect of the carbon cycle is that it is the cycle by which solar energy is transferred to biological systems and ultimately to the geosphere andanthrosphere as fossil carbon and fossil fuels.

2 Solubilization and chemical processes

  Atmospheric carbon dioxide is fixed as organic matter by photosynthesis, and2 organic carbon is released as CO by microbial decay of organic matter.2 The Nitrogen Cycle As shown in Figure 1.6 , nitrogen occurs prominently in all the spheres of the environment. The N molecule is very stable so that breaking it down into atoms that can 2 be incorporated with inorganic and organic chemical forms of nitrogen is the limiting step in the nitrogen cycle.

2 Inorganic nitrates

  The production of gaseous N and N O by microorganisms and the evolution of 2 2 these gases to the atmosphere completes the nitrogen cycle through a process called denitrification . It involves the interchange of oxygen between the elemental form of gaseous O , 2 contained in a huge reservoir in the atmosphere, and chemically bound O in CO , 2 H O, and organic matter.

3 A particularly important aspect of the oxygen cycle is stratospheric ozone, O

  3 As discussed in Chapter 9, Section 9.9, a relatively small concentration of ozone in the stratosphere, more than 10 kilometers high in the atmosphere, filters out ultra- violet radiation in the wavelength range of 220-330 nm, thus protecting life on Earthfrom the highly damaging effects of this radiation. In the geosphere, phosphorus is held largely in poorly soluble minerals, such as hydroxyapatite a calcium salt, deposits ofwhich constitute the major reservoir of environmental phosphate.

2 Inorganic SO - in both soluble

  Among the significant species involved in the sulfur cycle are gaseous hydrogen sulfide, H S; mineral sulfides, such as PbS, sulfuric acid, H SO , the main constitu- 2 2 4 ent of acid rain; and biologically bound sulfur in sulfur-containing proteins. Insofar as pollution is concerned, the most significant part of the sulfur cycle is the presence of pollutant SO gas and H SO in the atmosphere.


  A reasonable definition of a pollutant is a substance present in greater than natural concentration as a result of human activity that has a net detrimental effectupon its environment or upon something of value in that environment. Pollution of Various Spheres of the Environment Pollution of surface water and groundwater are discussed in some detail in Chapter 7, Particulate air pollutants are covered in Chapter 10, gaseous inorganic air pollutants in Chapter 11, and organic air pollutants and associated photochemicalsmog in Chapters 12 and 13.


  Some of the major ways in which modern technology has contributed to envi- ronmental alteration and pollution are the following: Despite all of the problems that it raises, technology based on a firm foundation of environmental science can be very effectively applied to the solution of environ-mental problems. Then place each of the following with the appropriate arrow, indicating the direction of its movement with a notation such as At Hy: (a) → Iron ore used for steel making, (b) waste heat from coal-fired electricity generation, (c) hay, (d) cotton, (e) water from the ocean as it enters the hydrologiccycle, (f) snow, (g) argon used as an inert gas shield for welding.

4 How is this process related to aerobic respiration?

  Define cycles of matter and explain how the definition given relates to the defini- tion of environmental chemistry. Describe the role of organisms in the nitrogen cycle.


  During most of its time on Earth, humankind made little impact on the planet, and its small, widely scattered anthrospheric artifacts—simple huts or tents fordwellings, narrow trails worn across the land for movement, clearings in forests to grow some food—rested lightly on the land with virtually no impact. However, withincreasing effect as the industrial revolution developed, and especially during the last century, humans have built structures and modified the other environmental spheres,especially the geosphere, such that it is necessary to consider the anthrosphere as a separate area with pronounced, sometimes overwhelming influence on theenvironment as a whole.


  Among the major advances during this century were widespread use of steam power, steam-powered railroads, thetelegraph, telephone, electricity as a power source, textiles, the use of iron and steel in building and bridge construction, cement, photography, and the invention of theinternal combustion engine, which revolutionized transportation in the following century. Since about 1900, advancing technology has been characterized by vastly increased uses of energy; greatly increased speed in manufacturing processes,information transfer, computation, transportation, and communication; automated control; a vast new variety of chemicals; new and improved materials for newapplications; and, more recently, the widespread application of computers to manu- facturing, communication, and transportation.


  Some of the 1 major components of the infrastructure of a modern society are the following: In general, the infrastructure refers to the facilities that large segments of a population must use in common in order for a society to function. Whereas advances in technology and the invention of new machines and devices enabled rapid advances in the development of the infrastructure during the 1800s andearly 1900s, it may be anticipated that advances in electronics and computers will have a comparable effect in the future.


  In relatively affluent societies the quality of living The construction and use of modern homes and the other buildings in which people spend most of their time place tremendous strains on their environmentalsupport systems and cause a great deal of environmental damage. It has been pointed out that all too often the design and operation of modern homes and other buildings takes place “out of thecontext” of the surroundings and the people who must work in and occupy the 2 buildings.


  A new development that is just beginning to reshape the way humans move, where they live, and how they live, is the growth of a telecommuter society com-posed of workers who do their work at home and “commute” through their com- puters, modems, FAX machines, and the Internet connected by way of high-speedtelephone communication lines. For example, such a capability enables detection of perturbations in environmentalsystems, analysis of the data to determine the nature and severity of the pollution problems causing such perturbations, and rapid communication of the findings to allinterested parties.


  Thismay consist of the synthesis of a chemical from raw materials, casting of metal or plastic parts, assembly of parts into a device or product, or any of the other thingsthat go into producing a product that is needed in the marketplace. A key component of an automated system is the control system, which regulates the response of components of asystem as a function of conditions, particularly those of time or location.


  Such materials have accumulated in the anthrosphere in painted and coated surfaces, such as organotin-containing paints used to prevent biofouling on boats; under and adjacent to airport runways; under and along highway paving; buried in old factory sites; in landfills;and in materials dredged from waterways and harbors that are sometimes used as landfill on which buildings, airport runways and other structures have been placed. The relatively harmonious relationship between the anthrosphere and the rest of the environment began to change markedly with the introduction of machines, par-ticularly power sources, beginning with the steam engine, that greatly multiplied the capabilities of humans to alter their surroundings.


  The company also advises purchasers of haz-ardous chemicals on ways to reduce inventories of such chemicals, and in some cases even ways to produce the materials on site to reduce the transportation of Scenario Creation to Avoid Environmental Problems A major concern with rapidly developing new technologies, of which transgenic biotechnology is a prime example, is the emergence of problems, often related topublic concern, that were unforeseen in the development of the technologies. G., the job of a panel on scenario creation is “to think the unthinkable and speak the unspeakable, not to say what we think willor should happen.” That statement is supported by the fact that scenario creation was developed by the Rand Corporation in the 1950s under sponsorship by U.


  Soil chemistry is covered in Chapter 16, “Soil Environmental Chemistry.” Human activities have such a profound effect on the environment that it is convenient to invoke a fifth sphere of the environment called the “anthrosphere.”Much of the influence of human activity on the environment is addressed inChapters 1-16, particularly as it relates to water and air pollution. What largely caused or marked the change between the “relatively harmonious relationship between the anthrosphere and the rest of the environment” that char-acterized most of human existence on Earth, and the current situation in which the anthrosphere is a highly perturbing, potentially damaging influence?


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  The chem-istry of water exposed to the atmosphere is quite different from that of water at the The study of water is known as hydrology and is divided into a number of subcategories. For example, disturbance of land by conversion of grasslands or forests to agricultural land or intensification of agricultural production may reduce The water that humans use is primarily fresh surface water and groundwater, the sources of which may differ from each other significantly.


  Hydrogen bonds are a special type of bond that can form between the hydrogen in one water molecule and the oxygen in another water molecule. Because of this high heat capacity, a relatively large amount × × of heat is required to change appreciably the temperature of a mass of water; hence, a body of water can have a stabilizing effect upon the temperature of nearbygeographic regions.


  As a result of exposure to the atmosphere and (during daylight hours) because of the photosynthetic activity of algae, the epilimnion containsrelatively higher levels of dissolved oxygen and generally is aerobic. During the overturn, the chemical and physical characteristics ofthe body of water become much more uniform, and a number of chemical, physical, and biological changes may result.


  The living organisms (biota) in an aquatic ecosystem may be classified as either autotrophic or heterotrophic. Algae are the most important autotrophic aquaticorganisms because they are producers that utilize solar energy to generate biomass from CO and other simple inorganic species.

2 Heterotrophic organisms utilize the organic substances produced by autotrophic

  Decomposers (or reducers) are a subclass of the heterotrophic organismsand consist of chiefly bacteria and fungi, which ultimately break down material of biological origin to the simple compounds originally fixed by the autotrophicorganisms. Biochemical oxygen demand, BOD, discussed as a water pollutant in Section 7.9, refers to the amount of oxygen utilized when the organicmatter in a given volume of water is degraded biologically.


  Oxidation-reduction reactions and equilibria are discussed in Chapter 4, and details of solubility calculations andinteractions between liquid water and other phases are given in Chapter 5. Biological processes play a key pH of water by removing aqueous CO , thereby converting an HCO 2 3 3 2+ ion; this ion in turn reacts with Ca in water to precipitate CaCO .

2 Gas exchange with the atmosphere -

3 2HCO + h2 ν 3 2 2 3 - - -

2 Acid-base HCO + OH CO + H O

  Thus, an exact description of the chemistry of a natural water system based upon acid-base,solubility, and complexation equilibrium constants, redox potential, pH, and other chemical parameters is not possible. Though not exact, nor entirely realistic, such models can yield useful generalizations and insightspertaining to the nature of aquatic chemical processes, and provide guidelines for the description and measurement of natural water systems.


Dissolved gases—O for fish and CO for photosynthetic algae—are crucial to the 2 2 welfare of living species in water. Some gases in water can also cause problems, such as the death of fish from bubbles of nitrogen formed in the blood caused by exposure

2 Lake Nyos in the African country of Cameroon asphyxiated 1,700 people in 1986

  The solubilities of gases in water are calculated with Henry’s Law, which states that the solubility of a gas in a liquid is proportional to the partial pressure of thatThese calculations are discussed in some detail in gas in contact with the liquid. Oxygen is produced by the photosynthetic action of algae, but this process is really not an efficient means ofoxygenating water because some of the oxygen formed by photosynthesis during the daylight hours is lost at night when the algae consume oxygen as part of theirmetabolic processes.

2 The weight of organic material required to consume the 8.3 mg of O in a liter of

  2 water in equilibrium with the atmosphere at 25˚C is given by a simple stoichiometric 2 the microorganism-mediated degradation of only 7 or 8 mg of organic material can completely consume the oxygen in one liter of water initially saturated with air at25˚C. At higher temperatures, the decreased solubility of oxygen, combined with the increased respiration rate of aquatic organisms, frequentlycauses a condition in which a higher demand for oxygen accompanied by lower solubility of the gas in water results in severe oxygen depletion.


species act as acids in water by releasing H ion, others act as bases by accepting H , and the water molecule itself does both. An important species in the acid-basechemistry of water is bicarbonate ion, HCO 3 base: (3.7.1) 3 ←→ 3 3 ←→ 2

2 Acidity as applied to natural water and wastewater is the capacity of the water to

  The term free mineral acid is applied to strong acids such as H SO and HCl 2 4 in water. Whereas total acidity is determined by titrationwith base to the phenolphthalein endpoint (pH 8.2), free mineral acid is determined by titration with base to the methyl orange endpoint (pH 4.3).

5 Some industrial wastes, such as spent steel pickling liquor, contain acidic metal ions

  The acidity of such wastes must be measured in calculating the amount of lime or other chemicals required to neutralize the acid. Because of the 2 presence of carbon dioxide in air and its production from microbial decay of organic matter, dissolved CO is present in virtually all natural waters and wastewaters.

2 Rainfall from even an absolutely unpolluted atmosphere is slightly acidic due to the

  Carbon dioxide, and its ionization products, bicarbonate 2 2 ion (HCO -) have an extremely important influence upon 3 3 the chemistry of water. Algae in water utilize dissolved CO in the synthesis of biomass.

2 CO (water) (atmosphere) (3.7.4)

  As a consequence of the low level of atmospheric CO , water totally lacking in alkalinity 2 3 2 and CO - greatly increases the solubility of carbon dioxide. As water seeps through layers of decaying organic matter while infiltrating the ground, it maydissolve a great deal of CO produced by the respiration of organisms in the soil.

2 Later, as water goes through limestone formations, it dissolves calcium carbonate

  Although CO in water is often represented as H CO , the equilibrium constant 2 2 3 for the reactionCO (aq) + H O H CO (3.7.7) 2 2 ←→ 2 3 3 is only around 2 10- at 25˚C, so just a small fraction of the dissolved carbon × dioxide is actually present as H CO . For CO in aqueous solution, the 2 2 diagram is a series of plots of the fractions present as CO , HCO -, and CO - as a 2 3 3 function of pH.

5 M. The carbon dioxide dissociates partially in

(3.7.18) 2 2 ←→ 3 3 a 1 2 ] ][H ][HCO [H 7

3 K = = = 4.45 10- (.7.19)

× 1a 5 ][CO 1.146 10- × 2 7 1 6 / 2 [H ] = [HCO 10- 4.45 10- ) 10- × × × = 2.25 × 3 pH = 5.65This calculation explains why pure water that has equilibrated with the unpolluted atmosphere is slightly acidic with a pH somewhat less than 7.


CO + H O (3.8.1) → 3 2 2 HCO(3.8.2) → 3 3 (3.8.3) →

2 Other, usually minor, contributors to alkalinity are ammonia and the conjugate bases of phosphoric, silicic, boric, and organic acids

2 contributors to it are HCO -, CO -, and OH-: 3 3 [alk] = [HCO - ] + 2[CO -] + [OH]- - [H ] (3.8.4) 3

3 Alkalinity generally is expressed as phenolphthalein alkalinity, corresponding to

titration with acid to the pH at which HCO 3 (pH 8.3), or total alkalinity, corresponding to titration with acid to the methyl

2 It is important to distinguish between high basicity, manifested by an elevated

  Expressing alkalinity in terms of mg/L of CaCO can, however, lead to confusion, and the 3 Contributors to Alkalinity at Different pH Values Natural water typically has an alkalinity, designated here as “[alk],” of 1.00 10- × 3 equivalents per liter (eq/L), meaning that the alkaline solutes in 1 liter of the water 3 will neutralize 1.00 10- moles of acid. Substitution× 3 3 3 into the expression for K shows that at pH 7.00 and [HCO 10- M, the 1 3a 4 value of [CO (aq)] is 2.25 10- M, somewhat higher than the value that arises × 2 from water in equilibrium with atmospheric air, but readily reached due to the presence of carbon dioxide from bacterial decay in water and sediments.

2 HCO , CO -, and OH- are

3 3 14 4 10- × Kw = 1.00 = 1.00 10- (3.8.7) × [OH-] = [H ]10- × 1.00 andK -] [HCO 2 2 3a (3.8.8)[CO -] = 3 ][H 4 Solving these three equations gives [HCO 10- M and [CO × 3 × 3 4 10- M, so the contributions to the alkalinity of this solution are the following: 4.64 10- eq/L from HCO × 3 4 2 2 2.18 10- = 4.36 10- eq/L from CO × × × 3 4 1.00 10- eq/L from OH- × 3 alk = 1.00 10- eq/L ×Dissolved Inorganic Carbon and Alkalinity The values given above can be used to show that at the same alkalinity value the concentration of total dissolved inorganic carbon, [C], 2 [C] = [CO ] + [HCO - ] + [CO -] (3.8.9) 2 3 3 varies with pH. At pH 7.00, 4 3 3 [C] = 2.25 10- + 1.00 10- + 0 = 1.22 10- (3.8.10) × × × 7pH whereas at pH 10.00, 4 4 4 [C] = 0 + 4.64 10- + 2.18 10- = 6.82 10- (3.8.11) × × ×pH

10 The calculation above shows that the dissolved inorganic carbon concentration at

  This pH-dependent difference in dissolved inorganic carbon concentration represents a significant potential source of carbon for algae growing in water whichfix carbon by the overall reactions CO + H O + h {C O (3.8.12) ν → Η Ο} + 2 2 2 2 andHCO - + H O + h {C OH- + O (3.8.13) ν → Η Ο} + 3 2 2 2 As dissolved inorganic carbon is used up to synthesize biomass, {CH O}, the 2 water becomes more basic. The amount of inorganic carbon that can be consumed 3 consumed from 1.00 L of water having an alkalinity of 1.00 10- eq/L is × [C] 1 L - [C] 1 L = × × 7 10pH pH 3 4 4 1.22 10- mol - 6.82 10- mol = 5.4 10- mol (3.8.14) × × × 4 This translates to an increase of 5.4 10- mol/L of biomass.

2 Assuming no input of additional CO , at higher alkalinity more biomass is produced

  2 for the same change in pH, whereas at lower alkalinity less is produced. Because of this effect, biologists use alkalinity as a measure of water fertility.

2 The increased solubility of carbon dioxide in water with an elevated alkalinity can

be illustrated by comparing its solubility in pure water (alkalinity 0) to its solubility in 3 3 water initially containing 1.00 10- M NaOH (alkalinity 1.00 10- eq/L). The × × number of moles of CO that will dissolve in a liter of pure water from the 2 atmosphere containing 350 ppm carbon dioxide isSolubility = [CO (aq)] + [HCO - ] (3.8.15) 2 3 Substituting values calculated in Section 3.7 gives 5 6 5 Solubility = 1.146 10- + 2.25 10- = 1.371 10- M× × × 3 The solubility of CO in water, initially 1.00 10- M in NaOH, is about 100-fold× 2 higher because of uptake of CO by the reaction 2 2 ←→ 3 so thatSolubility = [CO (aq)] + [HCO -] 2 3 5 3 3 = 1.146 10- + 1.00 10- = 1.01 10- M (3.8.17) × × ×


  A bare 2+ metal ion, Ca for example, cannot exist as a separate entity in water. In order to secure the highest stability of their outer electron shells, metal ions in water arebonded, or coordinated, to other species.

2 These all provide means through which metal ions in water are transformed to more

  Because of reactions such as these and the formation of dimeric species, 4+3+ such as Fe (OH) , the concentration of simple hydrated Fe(H O) ion in water is 2 2 2 6 vanishingly small; the same holds true for many other hydrated metal ions dissolved in water. A good example is acid mine water (seeChapter 7), which derives part of its acidic character from the acidic nature of hydrated iron(III): Fe(H O) Fe(OH) (s) + 3H + 3H O (3.9.5) 2 6 ←→ 3 2 H O 2+ 4 + 2Fe(H O) OH + 2H O (3.9.6) → (H O) Fe Fe(H O) 2 4 2 4 2 5 2 O HAmong the metals other than iron(III) forming polymeric species with OH- as a bridg- ing group are Al(III), Be(II), Bi(III), Ce(IV), Co(III), Cu(II), Ga(III), Mo(V), Pb(II),Sc(II), Sn(IV), and U(VI).

3 Calcium in Water

  The chemistry of calcium, although complicated enough, is simplerthan that of the transition metal ions found in water. Calcium is a key element in many geochemical processes, and minerals constitute the primary sources of calciumion in waters.

3 Calcium ion, along with magnesium and sometimes iron(II) ion, accounts for

  Temporary hardness is due to the presence of calcium and bicarbonate ions in water and may be eliminated by boiling the water: Ca + 2HCO CaCO (s) + CO (g) + H O (3.9.7) 3 ←→ 3 2 2 Increased temperature may force this reaction to the right by evolving CO gas, and a 2 white precipitate of calcium carbonate may form in boiling water having temporary hardness. The carbon dioxide that water may gain by equili- bration with the atmosphere is not sufficient to account for the levels of calcium → 2 2 2 2 accounts for the very high levels of CO and HCO 2 3 important in aquatic chemical processes and geochemical transformations.

3 K = [H -] = 4.45 10- (.9.10)

×a1 ][CO 2 3 11

3 K = [H 10- (.9.11)

= 4.69 × a2 [HCO 3 and the solubility product of calcium carbonate (calcite): 2+ 2

9 K = [Ca ][CO 10- (3..12)

3sp The reaction between calcium carbonate and dissolved CO is 2 3 2 2 ←→ 3 for which the equilibrium expression is the following: 2+ 2 [Ca ][HCO K K 5sp 3 a

1 K' = = = 4.24 0- (3.9.4)

  × [CO 2 ] K 2a The stoichiometry of Reaction 3.9.13 gives a bicarbonate ion concentration that is twice that of calcium. Substitution of the value of CO concentration into the expres- 2 4 2+ 4 sion for K' yields values of 4.99 10- M for [Ca ] and 9.98 10- for [HCO × × - ].

3 K K = = 2.09 10- (.9.15)

  2 ][CO 2 a value of 5.17 10- M is obtained for [H ] (pH 8.29). The alkalinity is essentially × 2 3 equal to [HCO -], which is much higher than [CO -] or [OH-].

3 Factors such as nonequilibrium conditions, high CO concentrations in bottom

  2 regions, and increased pH due to algal uptake of CO cause deviations from these 2 values. Nevertheless, they are close to the values found in a large number of natural water bodies.


  The properties of metals dissolved in water depend largely upon the nature of metal species dissolved in the water. In addition to 3+ the hydrated metal ions, for example, Fe(H O) and hydroxy species such as 2 6 2+ FeOH(H O) discussed in the preceding section, metals may exist in water revers- 2 5 ibly bound to inorganic anions or to organic compounds as metal complexes.

4 Fe(CN)

  This phenomenon is called complexation; the species that binds with the metal ion, CN- in the example above, is called a ligand, and theproduct in which the ligand is bound with the metal ion is a complex, complex ion, or coordination compound. In the example above, the cyanide ion is a unidentate ligand, which means that it possesses only one site that bonds to a metal ion.

2 CH

2 N H C

2 O C

  The ligands found in natural waters and wastewaters contain a variety of functional groups which can donate the electrons required to bond the ligand to a 2 metal ion. Chelates formed by the strong chelating agent ethylenediaminetetraacetate (EDTA, structure illustrated at the beginning of Section 3.13) have been shown to greatly 60 increase the migration rates of radioactive Co from pits and trenches used by theOak Ridge National Laboratory in Oak Ridge, Tennessee, for disposal of 3 intermediate-level radioactive waste .


  However, in solution, ligands of many complexes exchange rapidly between solution and the coordinationsphere of the central metal ion. The coordination number of a metal atom, or ion, is the number of ligand electron-donor groups that are bonded to it.

3.9.6. Selectivity and Specificity in Chelation

  One example of such a chelating agent is ferri- chrome, synthesized by and extracted from fungi, which forms extremely stablechelates with iron(III). It has been observed that cyanobacteria of the Anabaena species secrete appreciable quantities of iron-selective hydroxamate chelating agents 6 during periods of heavy algal bloom.


  The stability of complex ions in solution is expressed in terms of formation constants . These can be stepwise formation constants (K expressions) representing the bonding of individual ligands to a metal ion, or overall formation constants ( β expressions) representing the binding of two or more ligands to a metal ion.

3 K = = .9

(Stepwise formation constant)× 1 2+ ][Zn ][NH 3 2+ 2+ ZnNH + NH Zn(NH ) (3.12.3) 3 3 ←→ 3 2 2+ ] 2 [Zn(NH ) 3 10 (3.12.4)

2 K = = .1

  × 2 2+ ][ZnNH ][NH 3 3 2+2+ Zn + 2NH Zn(NH ) (3.12.5) 3 ←→ 3 2 2+ ][Zn(NH ) 4 3 2 = K = 8.2 10 (3.12.6) β = K × (Overall formation constant) 2 1 2 2 2+ ][Zn ][NH 3 2+ 2+ (For Zn(NH ) , = K K K and for Zn(NH ) , = K K K K .) β β 3 3 3 1 2 3 3 4 4 1 2 3 4 The following sections show some calculations involving chelated metal ions in aquatic systems. Because of their complexity, the details of these calculations may be beyond the needs of some readers, who may choose to simply consider the results.


  At pH values of 11 or above, EDTA is essentially all in the 4 completely ionized tetranegative form, Y O C C C C O H H H H N C C N(Y ) H H H H O C C C C O O H H OConsider a wastewater with an alkaline pH of 11 containing copper(II) at a total level of 5.0 mg/L and excess uncomplexed EDTA at a level of 200 mg/L (expressedas the disodium salt, Na H C H O N •2H O, formula weight 372). The × 2 formation constant K of the copper-EDTA complex CuY 1 2 [CuY 10 (3.13.1) × 1 = = 6.3 2+ 4 [Cu ][Y -]The ratio of complexed copper to uncomplexed copper is 2 4 4 18 15 [CuY= 5.4 10- 6.3 10 = 3.3 10 (3.13.2) × × × × 12+ ][Cu and, therefore, essentially all of the copper is present as the complex ion.

1 It is seen that in the medium described, the concentration of hydrated copper(II)

  Any phenomenon in solution that depends upon the concentration of the hydrated copper(II) ion (such as aphysiological effect or an electrode response) would be very different in the medium 2+ mg/L were present as Cu in a more acidic solution and in the absence of complexing agent. The phenomenon of reducing the concentration of hydrated metalion to very low values through the action of strong chelating agents is one of the most important effects of complexation in natural aquatic systems.


  Generally, complexing agents, particularly chelating compounds, are conjugate bases of Brönsted acids; for example, glycinate anion, H NCH CO -, is the conjugate 2 2 2 3 2 2 with metal ions for a ligand, so that the strength of chelation depends upon pH. In order to understand the competition between hydrogen ion and metal ion for a ligand, it is useful to know the distribution of ligand species as a function of pH.

3 The T

  2 nitrogen atom, as shown in Figure 3.10 . Note the similarity of the NTA structure to that of EDTA, discussed in Section 3.13.

2 T-]

1 1a a [H T] 3 H T- H + HT(3.14.3) 2 ←→ 3 ][HTK = [H -] = 1.12 10- pK = 2.95 (3.14.4) ×a 2 a 2 [H T-] 2

3 HT H + T

  3 2 2 As was shown for the CO /HCO -/CO - system in Section 3.7 and Figure 3.8 , 2 3 3 fractions of NTA species can be illustrated graphically by a diagram of the distribution-of-species with pH as a master (independent) variable. The key pointsused to plot such a diagram for NTA are given in Table 3.2 , and the plot of fractions of species ( values) as a function of pH is shown in Figure 3.11 .

2 The HT - species has an extremely wide range of predominance, however, spanning the entire normal pH range of ordinary fresh waters

Table 3.2. Fractions of NTA Species at Selected pH Values 23 pH value α α α αH T H T - HT T32 pH below 1.00 1.00 0.00 0.00 0.00 0.49 0.49 0.02 1 0.00 pH = pKa 1 pH = / ) 0.16 0.68 0.16 0.00(pKa 1 + pKa 2 2 0.02 0.49 0.49 2 0.00 pH = pKa 1 pH = / ) 0.00 0.00 1.00 0.00(pKa 2 + pKa 3 2 0.00 0.00 0.50 3 0.50 pH = pKa pH above 12 0.00 0.00 0.00 1.00 H T3 1.023 HT

0.8 H T-

  2 0.6αx 0.4 0.2 0.0 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 pH Figure 3.11. Plot of fraction of species α x as a function of pH for NTA species in water.


  A major concern regarding the widespread introduction of strong chelating agents such as NTA into aquatic ecosystems from sources such as detergents or electro-plating wastes is that of possible solubilization of toxic heavy metals from solids through the action of chelating agents. The extent of solubilization of heavy metals depends upon a number offactors, including the stability of the metal chelates, the concentration of the complexing agent in the water, pH, and the nature of the insoluble metal deposit.

2 Since [PbT-]/[HT -] is approximately 0 to 1, most of the NTA in solution is present

  The atomic weight of lead is 207, so the concentration of lead in solution is approximately 20 mg/L. In this example it is assumed that 25 mg/L of trisodium NTA is in equi- 2 librium with PbCO at pH 7.00 and a calculation is made to determine whether the 3 2 lead will be complexed appreciably by the NTA.

3 K' = = 4.69 10- = 10. (..9)

×a2 pK'a 2 3 [HCO where the acid dissociation constants of the carbonate species are designated as K'a todistinguish them from the acid dissociation constants of NTA. Figure 3.8 shows that within a pH range of about 7 to 10 the predominant carbonic species is HCO -; there- 3 2 fore, the CO - released by the reaction of NTA with PbCO will go into solution as 3 3 HCO 3 2 PbCO (s) + HT PbT - + HCO (3.15.12) 3 ←→ 3 This reaction and its equilibrium constant are obtained as follows: 2 PbCO (s) Pb + CO (3.15.13) 3 ←→3 2+ 2 13 K = [Pb ][CO -] = 1.48 10- (3.15.14)×sp 3

3 Pb + T PbT -

(3.15.6) ←→ 11 K = [PbT-] = 2.45 10 (3.15.7) ×f 2+ 3 ][T -][Pb 2 3 + HT H + T(3.15.4) ←→ ][T 10- (3.15.5)

11 K = [H -] = 5.25

× 3a 2

2 CO - + H

HCO (3.15.15) 3 ←→ 3 3 1 = [HCO(3.15.16) = 1 11 K' [CO ] 3 × 2a

2 PbCO (s) + HT PbT - + HCO (3.15.1)

3 ←→ 3 2 3 3 sp Ka f K = = K = 4.06 10- (3.15.17) × 2 [HT -] K' 2a From the expression for K, Equation 3.15.17, it may be seen that the degree to whichPbCO is solubilized as PbT - depends upon the concentration of HCO 3 3 this concentration will vary appreciably, the figure commonly used to describe natural 3 waters is a bicarbonate ion concentration of 1.00 10- , as shown in Section 3.9. × Using this value the following may be calculated: × 10= 40.6 [PbT-] = K = 4.06 (3.15.18) 3 10 3

1.00 Thus, under the given conditions, most of the NTA in equilibrium with solid PbCO

  As in the previous example, at a trisodiumNTA level of 25 mg/L, the concentration of soluble lead(II) would be approximately 20 mg/L. Effect of Calcium Ion upon the Reaction of Chelating Agents with Slightly Soluble Salts2+ Chelatable calcium ion, Ca , which is generally present in natural waters and wastewaters, competes for the chelating agent with a metal in a slightly soluble salt,such as PbCO .

11 K of NTA, 5.25 10- . The fraction of NTA bound as CaT- depends upon the

  × 3a 2+ 2+ 3 concentration of Ca and the pH. Typically, [Ca ] in water is 1.00 10- M.

3 K" =

(3.15.23) [CaT-][HReaction 3.15.22 may be obtained by subtracting Reaction 3.15.19 from Reaction3.15.12, and its equilibrium constant may be obtained by dividing the equilibrium constant of Reaction 3.15.19 into that of Reaction 3.15.12: 2 PbCO (s) + HT PbT - + HCO (3.15.12) 3 ←→ 3 2 3 3 s a f K = = K = 4.06 10- (3.15.17) × 2 2a ←→ 3 K' = 10- (3.15.20) ] = 7.75 × 2+ 2 [Ca][HT PbCO (s) + CaT- + H Ca + HCO (3.15.22) 3 ←→ 3 2 10- × K" = K = 4.06 = 5.24 (3.15.24)

3 K'

× 10- 7.75 3 2+ at pH 7.00, a concentration of HCO - of 1.00 10- , a concentration of Ca of 1.00 × 3 3 10- , and in equilibrium with solid PbCO , the distribution of NTA between the lead × 3 complex and the calcium complex is: [PbT -] 10- × × = [H 5.24 = 0.524]K" = 1.00 2+ 3 3 [CaT-] [Ca ] [HCO - ] 1.00 10 - 1.00 10- × × × 3 It may be seen that only about 1/3 of the NTA would be present as the lead chelate, 2+ whereas under the identical conditions, but in the absence of Ca , approximately all of the NTA in equilibrium with solid PbCO was chelated to NTA. Since the fraction 3 of NTA present as the lead chelate is directly proportional to the solubilization ofPbCO , differences in calcium concentration will affect the degree to which NTA 3 solubilizes lead from lead carbonate.


  The effect is to reduce the equilibrium concentration of calcium ion and prevent the precipitation of calcium car-bonate in installations such as water pipes and boilers. 3 The simplest form of phosphate is orthophosphate, PO -: 4 O 3-P O O O 3 4 3a1 a2 a third hydrogen ion is so difficult to remove from orthophosphate, as evidenced by the 3 very high value of pK , very basic conditions are required for PO - to be present at 3 4a significant levels in water.

4 Pyrophosphate ion, P O -, is the first of a series of unbranched chain poly-

2 7 phosphates produced by the condensation of orthophosphate: 3 4 2PO - + H O P O - + 2OH- (3.16.1) 4 2 ←→ 2

7 A long series of linear polyphosphates may be formed, the second of which is

  5 triphosphate ion, P O -. These species consist of PO tetrahedra with adjacent 3 10 4 tetrahedra sharing a common oxygen atom at one corner.


  In the a 2 a 3 a 4 case of triphosphoric acid, H P O , the first two pK values are small, pK is 2.30, 5 3 10 a a 3 pK is 6.50, and pK is 9.24. When linear polyphosphoric acids are titrated with a 4 a 5 base, the titration curve has an inflection at a pH of approximately 4.5 and another inflection at a pH close to 9.5.

10 Stepwise ionization of

O O O(3.16.2) P O O O O 3

10 Each P atom in the polyphosphate chain is attached to an -OH group that has one

  Therefore, one mole of triphosphoric acid, H P O , 5 3 10 loses three moles of hydrogen ion at a relatively low pH (below 4.5), leaving the 3 2 3 10 9.5), an additional two moles of “end hydrogens” are lost to form the P O 3 10 species. Titration of a linear-chain polyphosphoric acid up to pH 4.5 yields the number of moles of phosphorus atoms per mole of acid, and titration from pH 4.5 to pH 9.5yields the number of end phosphorus atoms.

4 Researchers have found evidence that algae and other microorganisms catalyze

  Therefore, there is muchless concern about the possibility of polyphosphates binding to heavy metal ions and transporting them than is the case with organic chelating agents such as NTA orEDTA, which must depend upon microbial degradation for their decomposition. The different chelating abilities of chain and ring phosphates are due to structural hindrance of bonding by the ring polyphosphates.


  If a material containing humic substances is extracted with strong base, and the resulting solution is acidified, theproducts are (a) a nonextractable plant residue called humin; (b) a material that precipitates from the acidified extract, called humic acid; and (c) an organic materialthat remains in the acidified solution, called fulvic acid. These substances contain a carbon skeleton witha high degree of aromatic character and with a large percentage of the molecular weight incorporated in functional groups, most of which contain oxygen.

3 O CO H O

  The binding of metal ions by humic substances is one of the most important environmental qualities of humic substances. It is now generally believed that these suspected carcin- ogens can be formed in the presence of humic substances during the disinfection ofraw municipal drinking water by chlorination (see Chapter 8).


  If V mL of acid of normality N are required to 3 p titrate V mL of sample to the phenolphthalein endpoint, what is the formula for s the phenolphthalein alkalinity as mg/L of CaCO ? Of the following components in the wastewater,the one that would not cause an increase in alkalinity due either to the component itself or to its reaction with limestone, is (a) NaOH, (b) CO , (c) HF, (d) HCl, (e) 2 all of the preceding would cause an increase in alkalinity.

3 After going through this aquifer for some distance and reaching equilibrium, the

  If the wastewater stream in Problem 10 were 0.100 M in NTA and contained other solutes that exerted a buffering action such that the final pH were 9.00, what would 2 be the equilibrium value of HT - concentration in moles/liter? Exactly 1.00 10- mole of CaCl , 0.100 mole of NaOH, and 0.100 mole of Na T × 2 3 2+ were mixed and diluted to 1.00 liter.

3 HCO - at equilibrium, what is the value of the ratio of the concentration of NTA

  The study of water is known as , , is the branch of the science dealing with the characteristics of fresh water, and thescience that deals with about 97% of all Earth’s water is called . List or describe one of each of the following unique properties of water related to(a) thermal characteristics, (b) transmission of light, (c) surface tension, (d) solvent properties.


__________________________ __________________________


  The overall reaction is 2+ 2+ Cd + Fe Cd + Fe(4.1.1) → This reaction is the sum of two half-reactions, a reduction half-reaction in which cadmium ion accepts two electrons and is reduced, 2+ Cd + 2e- Cd(4.1.2) → and an oxidation half-reaction in which elemental iron is oxidized: 2+ Fe Fe + 2e-(4.1.3) → When these two half-reactions are added algebraically, the electrons cancel on both sides and the result is the overall reaction given in Equation 4.1.1. Reduction of insolubleiron(III) to soluble iron(II), Fe(OH) (s) + 3H + e- Fe + 3H O (4.1.5) → 3 2 in a reservoir contaminates the water with dissolved iron, which is hard to remove in 4 3 NH + 2O NO + H O (4.1.6) → 4 2 3 2 converts ammonium nitrogen to nitrate, a form more assimilable by algae in the water.


In order to explain redox processes in natural waters it is necessary to have an understanding of redox reactions. In a formal sense such reactions can be viewed as

3 Suppose that the solution is treated with elemental hydrogen gas over a suitable

  As the reaction goes to the right, the hydrogen is oxidized as it changes from an oxidation state (number) of 0 in elemental H to a higher oxi- 2 dation number of +1 in H . 2+ 3+ + If the initial activities of H , Fe , and Fe were of the order of 1 (concentra- 2 2 3+ 2+ the left half-cell, Fe would be reduced to Fe in the right half-cell, and ions would migrate through the salt bridge to maintain electroneutrality in both half-cells.


  In this book, for the most part, pE and pE are used instead of E and E to more clearly illustrate redox equilibria in aquatic systems over many orders of magnitude 0.0591F EpE = 2.303RT = E (at 25˚C) (4.3.2) 0.0591F where R is the molar gas constant, T is the absolute temperature, and F is the Faraday constant. Just as pH is defined as pH = - log(a )(4.3.3) + H+ where a is the activity of hydrogen ion in solution, pE is defined as H pE = - log(a )(4.3.4) e- where a is the activity of the electron in solution.

2 Whereas it is relatively easy to visualize the activities of ions in terms of concentra-

  tion, it is harder to visualize the activity of the electron, and therefore pE, in similar terms. For example, at 25˚C in pure water, a medium of zero ionic strength, the 7 7 hydrogen ion concentration is 1.0 10 - M, the hydrogen-ion activity is 1.0 10 - , × × and the pH is 7.0.


  Referring to Figure 4.2 , if the Fe ion concentration is 2+ increased relative to the Fe ion concentration, it is readily visualized that the potential and the pE of the right electrode will become more positive because the 3+ higher concentration of electron-deficient Fe ions clustered around it tends to draw 3+ 2+ electrons from the electrode. 3+ 3 2+ 5 If, for example, the value of [Fe ] is 2.35 10- M and [Fe ] = 7.85 10 - M,× × the value of pE is 3 10- × pE = 13.2 + log 2.35 = 14.7 (4.4.4) 5× 10- 7.85 3+ 2+ As the concentration of Fe increases relative to the concentration of Fe , the value 2+ of pE becomes higher (more positive) and as the concentration of Fe increases 3+ relative to the concentration of Fe , the value of pE becomes lower (more negative).


  This reaction, which can be carried outin two separate half-cells as shown in Figure 4.3 , can be obtained by subtracting the lead half-reaction, Equation 4.5.5, from the copper half-reaction, Equation 4.5.3: 2+ (4.5.3)Cu + 2e- Cu pE = 5.71 ←→2+ (4.5.5) ←→2+ 2+ Cu + Pb Cu + Pb pE = 7.84 (4.5.7) ←→ The positive values of pE for this reaction, 7.84, indicate that the reaction tends to go to the right as written. 2+ 2+ If the activities of Cu and Pb are not unity, the direction of the reaction and value of pE are deduced from the Nernst equation.


  Imagine that instead of the cell being set up to measure the potential between the copper and lead electrodes, the voltmeter, V, wasremoved and the electrodes directly connected with a wire so that the current might flow between them. In predicting or explaining the behavior of an aquatic system, it is helpful to be able to predict the useful energy that can be extracted fromchemical reactions in the system, such as microbially mediated oxidation of organic matter to CO and water, or the fermentation of organic matter to methane by 2 anaerobic bacteria in the absence of oxygen.

5 Fe / O / H O Fe(OH) (s) + 2H pE = 7.6 (4.8.4)

  4 2 2 2 ←→ 3 From Equation 4.7.2, the standard free energy change for a reaction is G = - 2.303nRT(pE )(4.7.2) ∆ which, for a one electron-mole reaction is simplyG = - 2.303RT(pE ) (4.7.2) ∆ Therefore, for reactions written in terms of one electron-mole, a comparison of pE values provides a direct comparison of G values. Furthermore, the measured H potential will be more positive (more oxidizing) in an oxidizing medium, such as the aerobic surface layers of a lake, than in a reducing medium, such as the anaerobicbottom regions of a body of water.


  The condition under which oxygen from the oxidation of water has a pressure of1.00 atm can be regarded as the oxidizing limit of water whereas a hydrogen pressure of 1.00 atmosphere may be regarded as the reducing limit of water. 1 4/ (4.9.4) O 2 pE = 20.75 - pH(4.9.5) The pE-pH relationship for the reducing limit of water taken at P = exactly 1 H2 atm is given by the following derivation: H + e- / H pE = 0.00 (4.9.6) 2←→ 2 (4.9.7) pE = - pH(4.9.8) For neutral water (pH = 7.00), substitution into Equations 4.9.8 and 4.9.5 yields -7.00 to 13.75 for the pE range of water.


  24 reaction is 1 1 / CO + H + e- / + CH / H O (4.10.2) 8 8 4 4 2 2 ←→ for which the Nernst equation is1 /8 + P [H ] CO2 + pE = 2.87 + log = 2.87 + log[H ] = 2.87 - 7.00 = -4.13 (4.10.3)1 /8 + P [H ] CH4 1 / 4 7 × ×O 2 72 from which the pressure of oxygen is calculated to be 3.0 10- atm. For example, if a metal is beingconsidered, several different oxidation states of the metal, hydroxy complexes, and different forms of the solid metal oxide or hydroxide may exist in different regionsdescribed by the pE-pH diagram.

12 K = [Fe ] = 8.0

10 (4.11.3) ×sp [H ] Fe(OH) (s) + 3H Fe + 3H O (4.11.4) 3 ←→ 23+

3 K ' = [Fe ] = 9.1

  10 (4.11.5) ×sp [H ](The constants K and K ' are derived from the solubility products of Fe(OH) and sp sp 2 3 calculations.) Note that the formation of species such as Fe(OH) , Fe(OH) , and 2 3 is not considered. For the half-reaction 2+ Fe + 2e- Fe pE = - 7.45 (4.11.17) ←→2+ the Nernst equation gives pE as a function of [Fe ] 2+ 1 pE = - 7.45 + / log[Fe ] (4.11.18) 25 2+ For iron metal in equilibrium with 1.00 10- M Fe , the following pE value is × obtained: 5 1 pE = - 7.45 + / log 1.00 10- = - 9.95 (4.11.19) × 2 Examination of Figure 4.4 shows that the pE values for elemental iron in contact2+ with Fe is below the reducing limit of water.


  In additionto its multibillion dollar annual costs due to destruction of equipment and structures, corrosion introduces metals into water systems and destroys pollution control equip-ment and waste disposal pipes; it is aggravated by water and air pollutants and some kinds of hazardous wastes (see corrosive wastes in Chapter 19, Section 19.6). The area corroded is the anode, where the following oxidation reactionoccurs, illustrated for the formation of a divalent metal ion from a metal, M: 2+ M M + 2e-(4.12.1) → Several cathodic reactions are possible.

2 Oxygen may also be involved in cathodic reactions, including reduction to hydrox-

ide, reduction to water, and reduction to hydrogen peroxide: 4OH- (4.12.3) Ο Η Ο → 2 2 2(4.12.4) Ο → Η Ο 2 2OH- + (4.12.5) Ο Η Ο → Η Ο 2 2 2

2 Oxygen may either accelerate corrosion processes by participating in reactions such

  Morgan, Aquatic Chemistry: Chemical Equilibria, 3rd ed., John Wiley and Sons, Inc., New York, and Rates in Natural Waters 1995. Christ, Solutions, Minerals, and Equilibria, Harper and Row, New York, 1965.

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