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The production of coloured tissues, particularly insect-attracting petals, depends upon the synthesis of pigments. Plants are able to mix, modify and enhance pigments to produce a vast array of final petal colours. These colours are usually distributed across the flower in patterns, which vary in their degree of regularity and complexity between different species. While colour contrast is much more important than pattern for attracting pollinators from a distance, pattern becomes important at close range and allows animals to distinguish between flowers of different species and to learn to ‘handle’ flowers. This chapter considers the effects of mixing pigments together, the regulation of pigment distribution in the flower, and the use of metals, pH, and cell shape to modify the final colour of the flower.

The members of the Sc Group are Sc scandium, Y yttrium, La lanthanum, Ce cerium, Pr praseodymium, Nd neodymium, Pm promethium, Sm samarium, Eu europium, Gd gadolinium, Tb terbium, Dy dysprosium, Ho holmium, Er erbium, Tm thulium, Yb ytterbium, Lu lutetium, and Ac actinium. All these elements resemble each other greatly, especially in the series La–Lu (called the lanthanoids). Their slight differences may be assigned largely to size similarities, but a few oxidation state changes give rise to marked differences. The predominant oxidation state is III, but the IV state for Ce, and the II state for Eu are also important in their aqueous chemistries. The electron structures of these elements along with some other of their pertinent properties are shown in Table 12.1. Note the progression in the sizes of M+3 rising from Sc to Ac, but decreasing from La to Lu. This behavior causes Y+3 to fall in between Dy+3 and Ho+3, which results in yttrium’s chemistry usually resembling the latter lanthanoids. For this reason, Y will be treated as a lanthanoid in succeeding sections. The successive filling of the 4f electron level from La through Lu should also be noted, as well as the interesting 5d occupancy for Gd. The richest ore of Sc is the rare mineral thorveitite Sc2Si2O7, but it also occurs in very small quantities in some lanthanoid, uranium, and tungsten ores. Yttrium and the lanthanoids (abbreviated Ln), except for Pm, occur in monazite LnPO4 (mostly light lanthanoids), bastnaesite LnCO3F (mostly light lanthanoids), xenotime LnPO4 (mostly heavy lanthanoids), loparite (mostly light lanthanoids), and lateritic clays (some with mostly light lanthanoids, others with mostly heavy lanthanoids). All isotopes of Pm are radioactive and it does not occur with the lanthanoids. Exceedingly small amounts are present in uranium ores where it has been produced by the spontaneous fission of U-238. Its major source is artificial production, the longest lived isotope being Pm-145 (half life of 17.7 years). Ac is also without a stable isotope, the radioactive element resulting from the decay of naturally occurring Th and U. The longest lived Ac species is Ac-227 which has a half life of 21.77 years.

The elements making up the Actinoid Metals are those with atomic numbers from 89 through 103: Ac, Th, Pa, U, Np, Pu, Am, Cm, Bk, Cf, Es, Fm, Md, No, and Lr. The name is meant to parallel the lanthanoids. They are generally abbreviated as An. Their valence electron structures are 7s26d0−25f0−14. These elements resemble the lanthanoids somewhat, but they have a much wider variation in oxidation states. Nor do they resemble each other to the extent that the lanthanoids do, this being a result of the oxidation state variations. Ac resembles La greatly, but Th, Pa, and U resemble their vertical congeners (Hf, Ta, W) more than they resemble Ce, Pr, and Nd. From Np onwards, the resemblance to the lanthanoids increases such that by Am, the actinoid elements are behaving very similarly, showing a predominant oxidation state of III. All of this occurs because the 7s, 6d, and 5f levels are much closer in energy than the 6s, 5d, and 4f levels. Table 18.1 lists the actinoids with several of their pertinent characteristics. No stable isotopes of any of these elements exist, the last element in the Periodic Table with a stable isotope being Bi (Bi-209). However, some of the An elements have isotopes with very long half lives, which means that they are found in nature in relative abundance, most notably as Th-232 (1010.1 years), U-235 (108.8 years), and U-238 (109.7 years). Others are products of the decay of the above isotopes, so even though they are shorter lived, they persist in nature since they are continually being produced. The most important nuclides of this type are Ac-227 (21.8 years) and Pa-231 (104.5 years), both coming from U-235 decay. In U ores, very small amounts of Np-237 (106.3 years), Np-239 (2.4 days), and Pu-239(104.3 years) arise from the interaction of neutrons with U isotopes. Isotopes of the elements beyond U are produced artificially, Np and Pu by neutron capture by U, Am and Cm by multiple neutron capture by Pu, and elements beyond Cm by further neutron captures or bombardment of lower atomic number actinoids with ions of He, B, C, N, or O.

Soil forms a thin mantle over the Earth's surface and acts as the interface between the atmosphere and lithosphere, the outermost shell of the Earth. It is a multiphase system, consisting of mineral material, plant roots, water and gases, organic matter at various stages of decay, and a variety of live organisms. The first step towards understanding what controls the abundance and activities of these organisms, and also the factors that lead to spatial and temporal variability in soil biological communities, is to gain an understanding of the physical and chemical nature of the soil matrix in which they live. This chapter provides background on the factors responsible for regulating soil formation, and hence the variety of soils in the landscape. It also discusses the key properties of the soil environment that most influence soil biota, leading to variability in soil biological communities across different spatial and temporal scales.

In order to construct an E–pH diagram one needs to follow eight basic steps: (1) Select the species of the element involved which contain one or more of the following entities: the element, oxygen, and hydrogen. This is best done by reading the descriptive chemistry of the element in a good inorganic text and identifying the species, both soluble and insoluble, which persist, at least for several minutes, in aqueous solution. (2) Starting at the lower left-hand corner of an E–pH framework, arrange the selected species in vertical order of increasing oxidation number of the element. Then, if there are different species with the same oxidation number, arrange them in horizontal order of decreasing protonation (increasing hydroxylation). If there is only one species of a given oxidation number, this species extends across the entire pH range for the purposes of diagram construction. (3) Draw in border lines between the species, that is, the lines representing the transformation of a species to another species. You will not know exactly where these lines occur but the approximate regions are sufficient for the purposes of diagram construction. (4) Write equations for the transformations that have been indicated. Some of them will involve electrons and therefore will be half-reactions. Such equations must always be written as reductions, that is, with the electrons on the left. In addition, no reaction should contain the OH− ion; only the H+ and/or HOH instead. (5) From appropriate tabulations, obtain the standard free energy values (ΔG° in kJ/mole) of every species in the equations. These ΔG° values are to be employed in the following relationship which applies to each of the above equations. . . . ΔG° (reaction) = ∑ΔG° (products) − ∑ΔG° (reactants) (6) . . . (6) The ΔG° (reaction) values for each equation are to be converted into E° values for those equations containing electrons and into K values for those equations which do not. This is done by use of the following expressions: . . . E° = ΔG° /−96.49n log K = ΔG° /−5.7 (7/8) . . . where n represents the number of electrons in an equation.

In Chapter 2, a method for the construction of single-element E–pH diagrams has been presented. No agent which could produce an insoluble compound nor any agent which could complex with any of the simple ions in the single-element diagrams except OH− has been included. However, it is of interest in many cases to derive E–pH diagrams for systems which involve the precipitation or complexation of one or more of the simple ions present in the single-element diagram. One method of deriving such diagrams is to recognize that precipitation or complexation of a simple cation or anion reduces the concentration of the simple ion in solution considerably. If this reduced simple-ion concentration is calculated, it can then be used to construct a new E–pH diagram for the simple ion. Therefore, since the predominant species in the region labeled as the simple ion is no longer the simple cation or anion due to the precipitation or complexation, the region is re-labeled with the precipitated compound or the complex. E–pH diagrams which involve the precipitation or complexation of one or more of the simple ions present in the single-element diagram may also be obtained using one of the available computer programs. For complicated systems, the hand calculations become time consuming and it is often better to employ a computer program. In order to determine the changes in a single element E–pH diagram that occur due to the addition of a precipitating species and the resulting formation of an insoluble compound, the following five steps may be followed. (1) Select an element of interest showing a simple cation or anion and construct the E–pH diagram for the element at a soluble species equilibrium concentration of 10−4.0 M. (2) Select a precipitating agent which will form an insoluble compound with a simple cation or anion of the element of interest. In order to determine which species will form insoluble compounds with the simple ions of an element, it is often useful to consult a list of solubility rules.

The elements which constitute the Carbon Group of the Periodic Table are carbon C, silicon Si, germanium Ge, tin Sn, and lead Pb. All five of the elements have atoms characterized by an outer electron structure of ns2np2 with n representing the principal quantum number. This electron arrangement signals the possibility of oxidation states of IV and II. Such is the case with the II oxidation state becoming more stable from C to Pb. As one descends the group, there is a marked change from non-metallic (C) to metallic character (Pb). Reflecting very high ionization energies, C, Si, and Ge do not form a simple cation, they instead bond covalently. In line with the trends just mentioned, the inorganic aqueous chemistry moves from anionic (C) to cationic (Pb). The inorganic aqueous solution chemistry of C is represented by four acids and their anionic derivatives: carbonic acid H2CO3, oxalic acid H2C2O4, formic acid HOOCH, and acetic acid HOOCCH3. Note that in all of these the ionizing H+ ions are not attached to C but to O. The inorganic aqueous chemistry of Si is dominated by anions SiO(OH)3− and SiO2(OH)2−2 and their many polymeric forms and by the hexafluorosilicate anion SiF6−2. Ge is very similar to Si. Cationic species, largely absent in all three previous elements, are shown in both Sn and Pb. The covalent single bond radii of C, Si, and Ge are 77, 118, and 122 pm, and the ionic radii in pm of the other two elements are Sn+2(118), Sn+4 (83), Pb+2 (133), Pb+4 (92). a. E–pH diagrams. In order to understand the E–pH relationships of the aqueous species of C, it is important to consider both the thermodynamic and the kinetic relationships. Thermodynamics tells us whether a reaction will occur but it says nothing about how fast. The rate is a kinetic matter. When acetic acid HC2H3O2 is entered into a C species E–pH diagram, Figure 8.1 results. This figure shows that at equilibrium HC2H3O2 is not stable and disproportionates into H2CO3 and CH4. The same E–pH diagram results when formic acid HOOCH or when oxalic acid H2C2O4 is entered.

Water has an exceptional ability to dissolve minerals. It is safe and chemically stable, and it remains liquid over a wide temperature range. Thus, it is the best solvent and reaction medium for both laboratory and industrial purposes. Water is able to dissolve ionic and ionocovalent solids because of the high polarity of the molecule (dipole moment μ = 1.84 Debye) as well as the high dielectric constant of the liquid (ε = 78.5 at 25°C). This high polarity allows water to exhibit a strong solvating power: that is, the ability to fix onto ions as a result of electrical dipolar interactions. Water is also an ionizing liquid able to polarize an ionocovalent molecule. For example, the solvolysis phenomenon increases the polarization of the HCl molecule in aqueous solution. Finally, owing to the high dielectric constant of the liquid, water is a dissociating solvent that can decrease the electrostatic forces between solvated cations and anions, allowing their dispersion as H+solvated and Cl−solvated through the liquid. (The attractive force F between charges q and q′ separated by the distance r is given by Coulomb’s law, F = qq′/εr2.) These characteristics are rarely found together in common liquids. The dipole moment of the ethanol molecule (μ = 1.69 Debye) is close to that of water, but the dielectric constant of ethanol is much lower (ε = 24.3). Ethanol is a good solvating liquid, but a poor dissociating one; consequently, it is considered a bad solvent of ionic compounds. Dissolution in water of an ionic solid such as sodium chloride is limited to dipolar interactions with Na+ and Cl− ions and their dispersion in the liquid as solvated ions, regardless of the pH of the solution. Cations with higher charge, especially cations of transition metals, retain a fixed number of water molecules, thereby forming a true coordination complex [M(OH2)N]z+ with a well-defined geometry. In addition to the dipolar interactions, water molecules behave as true ligands because they are Lewis bases exerting an electron σ-donor effect on the empty orbitals of the cation.

Condensation of metal complexes in solution forms entities in which the cations are linked by hydroxo (HO−) or oxo (O2−) bridges. The reaction is initiated by the addition of a base to an aquocomplex: . . . 2[Cr(OH2)6]3++ 2HO- → [Cr2(OH)2(OH2)8]4+ + 2 H2O . . . or by the addition of an acid to an anionic complex: . . . 2 [CrO4]2- + 2H+ → [Cr2O7]2- + H2O . . . Thus, purely aquo- and purely oxocomplexes are stable in solution, and the condensation of cations is initiated by hydroxylation. With regard to electrically charged hydroxylated complexes, the reaction forms discrete and soluble entities—polycations and polyanions with a molecular complexity which depends on acidity conditions. This chapter presents a detailed study of their formation and structure. With regard to noncharged hydroxylated complexes, the condensation reaction is no longer limited and leads to the formation of a solid (a subject that is examined in the following chapters). The hydroxylation reaction is the key stage to initiate the condensation of cations in solution. It is thus important to precise the mechanism of the successive steps of the process, in order to understand why the behavior of a cation is closely related to its oxidation state, and why the reaction product may be a discrete molecular species or a solid. As a cation generally exhibits its maximum coordination number in the initial monomeric complex and in condensed species, the condensation reaction is a substitution that proceeds according to one of three basic mechanisms: dissocia­tion, association, and interchange or direct displacement [1, 2]. Dissociative substitution is a two-step process involving the formation of a reduced-coordination intermediate: In the first step, a labile ligand, the leaving group, breaks its bond in the starting complex before a nucleophilic entering group completes, in the second step, the cation coordination (Fig. 3.1 a). Associative substitution is also a two-step process in which the intermediate temporarily has increased coordination. The bond with the nucleophilic entering group (first step) occurs prior to the release of the leaving group (second step) (Fig. 3.1 b).

Cardiovascular disease knows no ethnic, national, or geographic boundaries. Men and women throughout the world can become affected by the obstruction of their arteries by cholesterol and the mineral hydroxylapatite (HA).This complex process leads to dysfunction of the arterial system and, because of the necessity of circulation of oxygenated blood, it also affects many tissues and organs. The whole process of occlusion (mineralization of the blood vessels), including precipitation and inorganic crystal formation, takes place in stages. The first stages are thought to involve cholesterol deposits (atherosclerotic plaque formation) in the interior of the vessel walls, or “intima,” as it is known. The formation of hydroylapatite, or “calcification,” begins with the attraction and localization of ions, mainly Ca2+ and PO43+, within the arteries. The vessels become altered and lose their suppleness, effectively interfering with their function as conduits for the blood (Pawlikowski 1986, 1991a,b, 1993, Pawlikowski et al. 1994). The initial stages of deposition can be detected with sensitive physical and chemical methods in vivo and with traditional laboratory methods and techniques on excised samples. In the mineralization stage, grains and crystals may become visible on heart valves as well as in the aortic tissue. (Pawlikowski and Pfitzner 1995,1999), and the new compounds can be identified using scanning electron microscopy and X-ray diffraction. Reasons for the destruction of tissues and the nucleation of minerals can be attributed to allodefects and autodefects. Autodefects in vessels are those attributable to abnormalities in the component tissues in the wall or pre-existing physical conditions. For example, at arterial bifurcations, intense local trauma from the flowing blood fluid might cause changes in those regions and attract cirulating ions. Autodefects are the result of reactions between biological tissues and foreign materials, such as small particles (dust) of all sorts, including bacteria or minerals that have been inhaled and travelled from the lungs via the blood into vessels throughout the body. Alternatively, allodefects may arise from poisons produced by bacteria and viruses during infections, or by other and various chemical products, such as food preservatives, that might be part of the circulating blood.

Disturbances disrupt ecosystem processes, thereby affecting the distribution and abundance of biota. Disturbances alter light and temperature regimes, carbon dioxide and nutrient fluxes, and productivity. The disturbance of one ecosystem factor is likely to affect most others, as biogeochemical cycles are coupled — to each other and also to landscape and anthropogenic factors. This chapter addresses how organisms respond to alterations of ecosystem processes in the air, soil, and water. Specifically, how do disturbances alter terrestrial light levels, air temperature, carbon dioxide levels, soil nutrients, soil pH, and soil organisms? In aquatic habitats, how do disturbances alter nutrient fluxes, enrichment, and acidification? In addition, transfers of energy and matter across interfaces of aerial, terrestrial, and aqueous parts of an ecosystem are covered. The responses of productivity to disturbance are also examined.

The determination of dissociation constants in non-aqueous and mixed-aqueous solvents is described. pH-scales, upon which the dissociation constants are based, are defined and compared with the corresponding aqueous scales. Experimental measurements, mostly based on potentiometric titrations using a glass electrode or spectrophotometric measurements of acid-base equilibria, are detailed. Absolute pKa-values are typically anchored to suitable indicator acids, such as picric acid, whose dissociation constants can be conveniently determined by several independent methods. Solute–solute interactions, especially specific association reactions, such as ion-pair and homohydrogen-bond formation, invariably accompany, and often dominate, acid–base equilibria in non-aqueous media. In mixed-aqueous solvents, simultaneous electrode calibration and pKa-determination may be achieved in a single, simple potentiometric titration. The measurement of autoionization constants, which quantify the tendency of a solvent to self-ionize, is also described.

This chapter draws up the abiotic frame for organisms set by the physical and chemical properties of a specific ecosystem. The abiotic frame is a combination of several features, including wind, turbulence, temperature and light, but also by nutrient status, pH and oxygen supply. Based on this abiotic frame, large-scale movements, as well as stratification phenomena of lakes are discussed. The importance of the surrounding land, that is, the catchment area, is stressed; specifically, how the catchment area may strongly affect the physical and chemical features of the lake or pond. In addition, this chapter explains how lakes and ponds have been, and still are, formed in the landscape and how organisms handle the abiotic frame.

Carnivorous Nepenthes pitcher plants contain aquatic ecosystems within each fluid-filled pitcher. Communities of arthropods and microbes colonize pitcher pools, and some organisms are endemic to the pitcher habitat. Flies and mites are the most apparent colonizers, and together with numerous protists, fungi, and bacteria, they form a food web of predators, decomposers, and primary producers. Bacterial diversity and composition are correlated strongly with fluid pH. Closely related organisms co-occur within pitchers, suggesting that competition is not the primary structuring force of pitcher communities. Pitchers are ephemeral habitats when compared with surrounding soil, and the former communities have fewer organisms and are less predictable than the latter. It is still unknown to what extent pitcher plants and their inhabitants influence one another’s fitness.

This chapter summarizes the far-reaching consequences and ecological implications of hypoxia and anoxia (low/no oxygen), which have become a global key stressor to marine ecosystems, with over 500 dead zones recorded worldwide. The focus is on coastal and estuarine ecosystems, which are increasingly at threat due to human impact and climate change. The contribution begins by discussing occurrences of and triggers for oxygen depletions, various definitions, thresholds, and types. It then presents responses to hypoxia/anoxia at various levels, i.e. the individual, population, community, and ecosystem levels. Examples illustrate the cascading effects, direct and indirect, of low dissolved oxygen triggering changes in behaviour, growth, recruitment, species diversity, biological interactions, trophic dynamics, community structure, and habitat complexity, ultimately threatening biodiversity and altering ecosystem structure and function. The chapter closes with a reflection on the interplay and synergy of multiple additional stressors increasingly being studied in combination with hypoxia such as hydrogen sulphide (H₂S), higher temperatures, and/or lowered pH.

Chemical equilibria in surface waters stem from complex interactions between physical background and living components of ecosystems. Catchments differ in geological background, climate, and land use; their run-off bears a distinctive chemical ‘fingerprint’. This chapter illustrates how the monitoring of standard parameters, such as oxygen, pH, conductivity, major ions, nutrients, and carbon, can lead to an interpretation of key aspects of the functioning of major ecosystem processes and how chemical constituents may affect the distribution of aquatic organisms. This requires understanding principles that underlie available measurement techniques and it demands a certain familiarity with the intrinsic variability of parameter values and of their chemical interaction. It is not required that field scientists be able to conduct detailed chemical assessments, but all should be able to collect samples yielding high-quality data. Therefore, detailed advice on chemical monitoring practice is provided, including sample collection, filtering, sample processing, and is discussed with the context of several case studies.

The production of coloured tissues, particularly insect-attracting petals, depends upon the synthesis of pigments. However, very few flowers are coloured simply by the synthesis of one single pigment, and equally few petals are composed of a single block of unchanging colour. Plants are able to mix, modify, and enhance pigments to produce a vast array of final petal colours. These colours are usually distributed across the flower in patterns, which vary in their degree of regularity and complexity between different species. While colour contrast is much more important than pattern for attracting pollinators from a distance, pattern becomes important at close range and allows animals to distinguish between flowers of different species and to learn to ‘handle’ flowers (extract the reward) with the minimum possible expenditure of energy. This chapter considers the effects of mixing pigments together, the regulation of pigment distribution in the flower and the use of metal ions and pH to alter petal colour. It concludes with a discussion of the role of petal surface structure in the production of colour and the modification of pigment colour.