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Earth Sciences
John Ferris

General Principles 1.0 Introduction The field of aqueous microbial geochemistry is highly interdisciplinary with scientific roots not only in biology, but also geology and physical chemistry. A major objective is to understand the central role of microbial processes in regulating the chemical composition of natural and contaminated aqueous systems including groundwaters, lakes and rivers, estuaries, and oceans. Even broader associations exist in terms of bioremediation, global elemental cycling, environmental change, and even possibilities for life beyond Earth. As an interdisciplinary science, aqueous microbial geochemistry is challenging in that it demands acquisition and application of a broad range of scientific knowledge. For individuals new to the field, the learning curve may appear to be frightfully steep; however, there are many common threads between different scientific disciplines and aqueous geochemistry, few of which are stronger or more compelling than the fundamental characteristics of natural waters. These are defined in physical, chemical, and biological terms that are not always used commonly across disciplines. One should not be surprised that lexicon is a confounding issue for interdisciplinary science. Natural philosophy has always been disguised by the passion of ardent observers, so much the better to obfuscate the obvious and lay claim to otherwise common scientific principles. For this reason, this chapter is focused on providing a modicum of terminology that usefully serves physical and biological scientists alike. 1.1 Chemical Speciation and Fractionation Natural waters, whether they are pristine or contaminated, are not simple aqueous solutions. At first blush, it is easy to acknowledge that a wide range of dissolved and solid materials are likely present in a water sample. Moreover, dissolved and solid concentrations are apt to be quite different depending on the provenance and environmental quality of the water sample. These simple considerations raise a number of problematic issues that relate to how dissolved and solid materials in natural waters are distinguished in a consistent manner across scientific disciplines. The different physicochemical forms adopted by an individual element or its compounds in aqueous solution are referred to as species. In this context, substances that differ in isotopic composition, molecular conformation, oxidation state, or in the nature of their ionic or covalently bound substituents, can be regarded as distinct chemical species. Further to this concept, the term speciation is used to indicate the specific distribution of chemical species that are present in a given sample. The physical activity of identifying and measuring different chemical species is speciation analysis. In conjunction with use of speciation to describe chemical species distributions, 1 fractionation is the classification of materials according to physical (e.g., size, solubility) or chemical (e.g., bonding, reactivity) properties. There is no perfect way to distinguish precisely between dissolved and particulate fractions in a water sample. This is because the size distributions of -10eous -9mponents vary in a continuous manner from ions and molecules in the 10 to 10 m range to bacteria and algae in the 10 to 10 m range. Materials between 10 and 10 m are9 -7 defined as colloids; however, use of the term nanoparticle has grown in popularity to describe a wide variety of colloidal materials. In natural waters, colloids are characterized by extreme diversity including organic macromolecules, biological debris, clay minerals, iron and manganese oxides, and even viruses. The most common approach in water analysis to separate dissolved and particulate fractions is to filter samples through membrane filters with pore sizes of 0.2 to 0.5 μm. Constituents of the dissolved fraction must be small enough to survive filtration, whereas the particulate solid fraction is retained by the filter. There are some obvious problems with this operation definition of dissolved and particulate fractions, especially in relation to the fractionation of colloidal particles; however, aggregation of colloids to form larger agglomerates and adherence to other particles like bacteria cells or algae enhances the fractionation efficiency of membrane filtration. Alternatively, filtration can be done using filters with nanometer-size pores or centrifugation can be used to sediment the particulate materials. 1.2 Concentrations Concentration is a term that refers to the quantity of a substance in a water sample. A number of different ways are used to express concentrations depending foremost on the field of application. Because of the interdisciplinary nature of physical biogeochemistry, all methods used to report concentrations are apt to be encountered. An especially common approach to expressing concentrations is based on reporting the quantity of a substance in a fixed quantity of solution. In this context, the quantities used are typically weight and/or volume. These different measurement possibilities give rise to several concentration scales including weight per weight (w/w), volume per volume (v/v), and weight per volume (w/v). The corresponding concentration values are reported generally in “parts per” format. For example, concentrated solutions are often specified in terms of parts per hundred, which is equivalent to percentage. In more dilute solutions, parts per thousand (ppt) or parts per million (ppm) are used, whereas parts per billion (ppb) is applied to describe concentrations in extremely dilute solutions. Mole concentrations are expressed in terms of molality m (mol kg of solvent) or -1 molarity M (mol L of solution). Molarity is the most commonly used mole 1Total dissolved solids (TDS) is another term commonly used in reference to the dissolved fraction of a water sample. 2 concentration scale, whereas molality is used almost exclusively to describe the physicochemical properties of solutions because it is based on a temperature-independent mass rather than a volume. An important point to recognize about molarity is that the volume of solvent and the volume of a solution are different. The former quantity can be determined from the molarity only if the densities of both the solution and pure solvent are known. One essential chemical criterion that must be respected for any solution is the principle of electroneutrality. This can be evaluated by first converting molar concentrations of ions to electrical charge equivalent concentrations. In general, the conversion of molar concentrations to equivalent concentrations is expressed by eq/L = z*M where z is the absolute charge of the ion under consideration . Example 1.1 A solution contains 0.1 mM Fe .+ The oxidation state of ferric iron corresponds to a molar charge concentration of 3 equivalents (eq)/mole (or 3 meq/mmole) so the equivalent concentration of the solution is 0.1 mM = 0.1 mmole/L * 3 meq/mmole = 0.3 meq/L Fe 3+ For electroneutrality to be respected, the summed equivalent concentrations of positively charged cations and negatively charged anions in solution must be equal (i.e., charge balance). a b ∑ cationsaeq/L =∑ anions bq/L 1 1 a b cations eq/L − anions eq/L = 0 ∑1 a ∑1 b The concept of charge balance is a useful way to evaluate whether a chemical analysis of a water sample is reasonably complete. This can be expressed in terms of a charge balance error (CBE) 3 a b ∑ cations aq/L − ∑ anions eb/L CBE = 1 1 ×100 a b ∑ cations aq/L + ∑ anions eb/L 1 1 In terms of the CBE, a good chemical analysis of a water sample will yield a value close to zero. Example 1.2 Consider a solution containing 2 mM FeCl and 53mM Na SO . The s2lut4on has 2 mM Fe , 6 mM Cl, 10 mM Na , and 5 mM SO . The charge balance is 4 6 meq/L Fe + 10 meq/L Na+ = 6 meq/L Cl + 10 meq/L SO 42- 16 meq/L cations = 16 meq/L anions 16 meq/L cations - 16 meq/L anions = 0 The chemical behavior of dissolved and solid substances in water are often evaluated in terms of ionic strength I, a parameter related to ion charge concentration by m I = 1 M (z ) 2 2 ∑ i i i 2 for m cations and anions. The second order dependency of I on z demonstiates that multivalent ions have a greater influence on ionic strength than monovalent ions. Example 1.3 Consider a solution containing 10 mM NaCl and 2 mM MgSO . There are 4 cations and 2 anions – 10 mM Na , 2 mM Mg , 10 mM Cl, and 2 mM SO . 42- I = 0.5 x [10(+1) + 10(-1) + 2(2+) + (2(-2) ] = 0.5 x [10+10+8+8] = 18 mM If the MgSO is4replaced with MgCl 2 + 2+ - Then there are 2 cations and only 1 anion – 10 mM Na , 2 mM Mg , and 14 mM Cl 4 So now the ionic strength is I = 0.5 x [10(+1) + 14(-1) + 2(+2) ] = 0.5 x [ 10+14+8] = 16 mM 2- Note the greater influence of the multivalent ion SO 4 on ionic strength compared to monovalent Cl. - A general classification scheme for different types of water is based on concentrations of total dissolved solids (TDS); fresh water < (1500 mg/L TDS) brackish water (5000 mg/L TDS) < saline water The most accurate value of ionic strength is obtained from a total water analysis that includes all ionic species; however, partial analyses are more commonly encountered rather than complete ones. In such circumstances, the molal ionic strength can be estimated from the TDS (mg L or ppm) or specific conductance (SPC, μS cm ) -1 of the water: -5 I = 2.5 x 10 x TDS “average waters” I = 2.8 x 10 x TDS Ca(HCO ) waters 3 2 I = 1.9 x 10 x SPC Ca(HCO ) w3 2rs The specific conductance measurement usually provides a more accurate estimate for ionic strength than TDS. This is because conductance and ionic strength are both measures of the total concentrations of ionic species, whereas TDS includes concentrations of uncharged species that do not contribute to ionic strength. 1.3 Physical Properties of Water In the liquid state, water is characterized by a number of unique physical and chemical properties. These include a high boiling point and high heat of vaporization; o the maximum density of water is near 4 C, which means that water expands upon freezing and that ice floats. It has a very high surface tension, the greatest dielectric constant of any liquid, and makes an excellent solvent for ions and ionizable solutes. All of the remarkable characteristics of water are manifest principally from a very pronounced dipole moment that is conferred by a combination of molecular composition and shape. A water molecule consists of two hydrogen atoms that are bonded to one oxygen atom. Because oxygen is more electronegative than hydrogen, the shared electron pair that makes-up each covalent O – H bond is bias toward the oxygen atom. If the water 5 molecule was linear, the electric dipoles associated with the two O–H bonds would be exactly opposite from each other and would cancel the electric imbalance; however, the water molecule is bent at an angle of 104.5 and the polarities of the two O – H bonds reinforce each other, giving rise to a permanent dipole moment associated with the molecule as a whole. In other words, the electric dipole moment of water is a vector property that arises from an uneven distribution of unlike charges. Dipole-dipole interactions between two adjacent water molecules form a hydrogen bond . When only two molecules are present, the product of hydrogen bonding is referred to as a water dimer . If addition molecules are nearby, which is the case in liquid water, more bonds are possible because the oxygen of one water molecule has two lone pairs of electrons, each of which can form a hydrogen bond with hydrogen atoms on other water molecules. This can repeat so that every water molecule is H-bonded with up to four other molecules; two through its two lone pairs on oxygen, and two through its -1 two hydrogen atoms. The energy of each hydrogen bond is low (ca., 21 kJ mol ); however, the formation of multiple hydrogen bonds in liquid water is a major factor contributing to a high boiling point, high melting point and high surface tension. Figure 1.1: Schematic diagram of the dipole-dipole interaction and formation of a hydrogen bond between two water molecules. 1.4 Solutions For the dissolution of a solute in a solvent, the solvent molecules must be displaced to accommodate the solute. In the case of water, this means disrupting structural arrangements arising from hydrogen bonding between adjacent water molecules. This requires an input of energy that can be recovered only if a stronger attractive interaction can be established between the water molecules and the solute. This will happen if the solute itself is sufficiently charged to replace the disrupted hydrogen bonds between water molecules with ion-dipole interactions of comparable or greater strength. 2 3A hydrogen bond between molecules with dipole moments is essentially an electrostatic interaction. A water dimer is the smallest water cluster established by hydrogen bonding, and serves frequently as a model for investigations on the physicochemical properties of water. 4An ion-dipole interaction is an attractive force that results from the electrostatic attraction between an ion and a neutrally charged molecule that has a dipole moment. 6 The strength of ion-dipole interactions depends on the size of an ion and its electric charge. Higher charge and smaller ionic diameters favor the establishment of ion-dipole bonds, which are ultimately r
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