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The general concept of speciation applies to many chemical systems.
It is most commonly used in the context of aqueous solutions and
the various gases and solids in contact with them. Typically, such
solutions include natural water systems such as sea, river, ground
and lake waters. However, they also extend to "biofluids" such as
blood plasma, intestinal secretions and plant xylem, as well as to
solutions used in laboratory and industrial processes.
 
The speciation of a particular chemical entity (like an element,
an ion, a molecule and so on) in a given system is defined here
as the identity and abundance of each and every physico-chemical
form in which that entity occurs.
 
The term 'chemical speciation' is sometimes used differently
elsewhere. See Bernhard M, Brinckman FE and Irgolic KJ in the
"Importance of Chemical Speciation in Environmental Processes"
(Springer-Verlag, Berlin, 1986) for a full discussion.
 
 

Chemical speciation is important because the properties of any chemical system often depend on the concentration of the individual species rather than simply on the total amounts present. A detailed knowledge of the chemical speciation of a system and, in particular, how it changes with conditions, is thus essential to understand fully the properties of the system and to predict how the system will behave under changed circumstances.
The chemical speciation of a system can sometimes be determined experimentally but most such analytical methods suffer from a number of problems. First, the measurement itself typically affects the chemistry of the system and disturbs the equilibria upon which the chemical speciation depends. Secondly, many experimental techniques are not sufficiently species specific, suffering from interferences of one kind or another. Thirdly, measurements are often not sensitive enough to determine the very low (and yet significant) concentrations at which some chemical species can occur. An alternative approach to determining chemical speciation is provided by calculations, based on the principles of thermodynamics. Models of the system are developed to predict how the system responds towards different equilibrium conditions. To fully appreciate this powerful approach it is necessary to understand some basic chemical principles concerning equilibria in solution. Consult the index searching for "equilibrium chemistry" if you need information in this respect.
Each chemical reaction occurring in a system involves 'reactants' which combine to form 'products'. In general, reactions proceed until equilibrium is attained. Equilibrium thus means the state in which there is no further tendency for the chemical change to take place. It is very important to appreciate that reactions often do not go 'to completion', i.e. at equilibrium, the reactants remain to some extent. The equilibrium position therefore represents a balance between the concentrations of reactants and products which is chemically stable. Increasing the concentrations of the products in a reaction mixture tends to drive the reaction backwards, i.e. equilibria are dynamic in nature, with the products capable of being transformed back into reactants if conditions change. For systems to achieve equilibrium it is necessary that the rates of reaction are sufficiently rapid, relative to the time under consideration. Most reactions in aqueous solution tend to reach equilibrium rather quickly; however, sometimes the timescale can be much longer. Precipitation processes in particular tend to be relatively slow. Note that the mechanism by which species react is totally irrelevant to the chemical speciation determined by thermodynamic calculations. Also note that, provided the equilibrium state is attained, no consideration is given to how long this takes to achieve. With equilibrium modelling, then, we are only concerned with the characteristics of the final (equilibrium) state itself. The issue of reaction kinetics, and the associated problem of knowing which chemical species are in equilibrium with one another, are profound matters which have wide-ranging repercussions for the modelling that you can do using JESS. Consult the index searching for 'quasi-equilibrium' (without quotes) to obtain more information.
The extent to which reactants and products exist at equilibrium depends directly on the dynamic nature of the equilibria involved. Changing the prevailing conditions will, in general, tend to make the equilibrium 'shift' in favour of either more or less formation of the products. Such shifts may be caused by many different factors. Changing the temperature, adding substances (either new ones or more of those already present), diluting the solution or concentrating it are good examples of the kinds of change which affect equilibria. Even adding substances that are not specifically part of a reaction may cause the equilibrium position to adjust itself. This typically occurs because reactions involving ions in solution are naturally sensitive to the presence of any other dissolved ions. This phenomenon can be related to a quantitative measure of the overall concentration of ions in the solution, known as the 'ionic strength'.
Species are often classified into one of two kinds - 'free' or 'complex'. For example, it is said that when a free metal ion becomes bound to a free ligand, a complex species is formed. A 'complex' is any species that is made up of two or more entities capable of existing independently. 'Free' species cannot be separated in this way. Accordingly, the 'free concentration' of an ion is that concentration of the ion not bound to any species (other than solvent molecules, often water). Likewise, the 'total concentration' of the ion is the overall concentration of the ion in all its forms (i.e. the 'free' ion plus all the complex species in which that ion occurs). In equilibrium calculations, you often need to specify the amounts of substances either as 'free' or as 'total' concentrations. An essential concept of chemical speciation is that different species exhibit different properties. For instance, complexation tends to reduce the overall electrical charge of the species in solution. Similarly, 'free species' tend to have a kinetic advantage over complexes - they are frequently more mobile in diffusion, for example. It is this kind of difference that makes the observed characteristics of chemical systems dependent on the chemical speciation. A dramatic example is the change in toxicity which occurs when cyanide is complexed by iron(II). The complex, ferrocyanide, is so safe that it can be added routinely to table salt as an anti-caking agent. Note that 'free' species are often said to be 'hydrated' to emphasise that only water is associated with them. This is a little confusing since, in equilibrium calculations, it is usual to consider all species to be fully 'solvated' (i.e. surrounded by, and interacting with, an appropriate number of solvent molecules).
Many different kinds of chemical reaction may occur in solution. A brief outline of some elementary, but important concepts, follows for those readers without sufficient chemical background. Acids and Bases One of the most important species in any aqueous solution is the hydrogen ion or proton (H+). It is the free concentration of this species that ultimately gives rise to the pH of the solution. The higher its concentration, the lower the pH. At high pH, the hydrogen ion concentration becomes very small and the hydroxide ion then predominates. The simplest way of thinking about acid-base reactions is based on the definition of acids as compounds that lose protons (e.g. by dissociation) and of bases as compounds that bind protons (either by direct association or, indirectly, by producing hydroxide for example). A strong acid loses its protons almost entirely, a weak acid loses them to a much lesser extent. Sulphuric acid is a strong acid; acetic and salicylic acids are weak acids. Conversely, a strong base such as hydroxide (OH-) binds protons strongly, whereas a weak base binds protons only weakly. Metal-ligand Equilibria Most metals in solution occur as positively-charged ions (called cations). These usually bind with negatively-charged ligands (called anions). A ligand is simply a species that binds to another species, frequently a metal ion. Ligands are often classified into inorganic (like nitrate, sulphate, chloride, etc.) and organic (like the amino acids and acetic acid). They are not always negatively charged - indeed neutral ligands (such as ammonia, NH3) are fairly common and even some positively charged ligands are known. Metal ions generally bind to more than one ligand to form complexes such as Cu(NH3)2. How many ligands can bind to a metal ion depends on how many electron-donating atoms (such as N, O, S and P) it has and on the chemical characteristics, particularly the size, of the metal ion. Precipitation Precipitates are formed when (and only when) the product of the concentrations of those species forming the solid exceeds a characteristic limit, known as the 'solubility product'. Whenever solids in contact with solution are in chemical equilibrium with it, the product of the concentrations of the species forming the solid is held at this constant value. This is so irrespective of how much solid is formed or how much more solid might be added from external sources (either as the solid or as its component species). Typically, the value of the solubility product varies with temperature and ionic strength (like equilibrium constants generally). However, under given conditions, the solubility product represents an invariant upper limit for the component species in solution or, equivalently, the lower limit of thermodynamic stability of the solid (below which the it cannot exist at equilibrium). Although this product of the concentrations of the species forming the solid remains constant in the presence of the solid, the total amount of individual component species (in terms of either their free or their total concentrations) may vary considerably. This is because (i) the concentrations of other component species may be adjusted to compensate and/or (ii) other interactions in solution may lead to some other binding of the solid's component species, increasing the total concentration of that component whilst nevertheless decreasing its contribution to the solubility product. Redox Equilibria A redox reaction is one in which one species is oxidized whilst another is reduced. Originally, oxidation was simply a process involving chemical reaction with oxygen. Nowadays, however, oxidation is regarded as a general type of reaction in which there is a loss of electrons; conversely, reduction occurs when any species gains electrons. Species vary in their tendency to be oxidized or reduced. For example, metal ions like those of calcium(II) are not easily oxidized nor reduced. On the other hand, ions like Fe+2 and Fe+3 are readily interconverted, depending only on whether the solution occurs in an oxidizing or reducing environment. For solutions in contact with air (i.e. under 'aerobic conditions') most of the iron would be in the oxidized form, Fe+3 whereas in the reducing environment of a peat bog (i.e. under 'anaerobic conditions') it would tend to be in the reduced form, Fe+2. Since the strength with which metal ions bind to ligands very much depends on their oxidation state, it is important to know in what form the metal predominates. This can, in general, be determined by equilibrium calculations.
In all but the most basic of solutions, chemical species compete amongst themselves for predominance. In general, species can exist only at the expense of other species that have one or more elements in common. For example, ligands compete with each other for access to the metal ions in solution. Which ligand is most successful depends on various factors such as the relative strengths of binding and the relative concentrations. Obviously, the higher the concentration of a particular ligand and the more strongly it tends to bind to a particular metal ion the more likely are those metal-ligand complexes to be formed. Other species that could be formed by the metal ion and the ligand in question can thus be precluded. Competition in solution also encompasses the formation of those species favoured by the pH of the solution, such as the appropriately protonated form of a ligand or of the appropriate hydroxy-complex formed by the metal ion. In this way, the pH generally determines the relative availabilty of metal ions and ligands to interact with one another, and hence also determines how successfully their complexes can compete for existence in solution. Since there are so many factors that influence the outcome of this competition between species it is generally necessary to perform the equilibrium calculation using a suitable computer program to predict which species will predominate. Attempts to guess the outcome of the competition between species in solution without recouse to such calculations are notoriously unreliable.
Chemical 'buffering' is the phenomenon exhibited by certain systems in which the effect of certain imposed changes (such as the addition acid) has a smaller impact than might be anticipated because the system responds in a way that opposes the change. The most frequently cited buffering action concerns pH. In such cases the addition of either acid or base produces only a relatively minor shift of pH because a component in solution (called the 'buffer') offsets the change by gaining or releasing protons respectively. However, the competitive nature of the equilibria established in many solutions means that components other than H+ are also often well buffered. Indeed, this happens whenever the imposed change acts on any equilibrium that is poised to counteract it. For example, a free metal ion concentration can be buffered by an equilibrium in which the metal is bound to a significant extent - addition of a complexing agent may then have little effect because the equilibrium will tend to replenish the free metal ion concentration by dissociation.
Equilibrium modelling is the chief means by which JESS provides information about chemical speciation in solution. However, it is essential to appreciate that this information is rarely achieved in the form of a single, quantitative answer; a less specific understanding of the chemical relationships involved in your system and of the uncertainties which exist (e.g. because of deficiencies in the scientific literature) is a much more frequent outcome. A more detailed discussion can be found in the section on "Modelling Techniques" in the JESS Primer. The Aims of Equilibrium Modelling Detailed information on the identities of the predominant species in solutions under given conditions is, as indicated above, an ideal objective that can truly be accomplished only for rather simple chemical systems. Nevertheless, the great predictive power of thermodynamics will often allow you to develop a sound perspective of the speciation, even in very complicated mixtures. In particular, you will regularly be able to see how the predominant species are affected by changing conditions. Predictions about the behaviour and properties of the solution should be your ultimate goal and can be often be accomplished.
The following functions are available in the estimation of UCCs for the correction of activity coefficients when equilibrium constants are not available at an ionic strength required. Linear D-H Limiting Law Guntelberg Davies Helgeson Morel All these functions have the advantage that they can be calculated without species-specific parameters (i.e. they can be used as 'stand-alone' models). For this reason, they can be used for 'global' corrections when there is no better information available. Other functions can be employed for activity coefficient corrections provided the necessary parameters have been entered into the JTH database for the species in question. For example, the Murray-Linder function has been evaluated for a number of simple species in the JESS Parent database, JPD. More information can be obtained by searching the index for "species-specific + functions" and for "ionic strength correction".
A number of extended Debye-Hueckel functions permitting species-specific correction of activity coefficients can be used in the GEM stage calculations. This obviously applies only to those species where the necessary parameters have been evaluated and entered into the source thermodynamic database. The options available include the Murray-Linder and the Guggenheim functions in addition to the stand-alone extended Debye-Hueckel formulae used as global correction methods. Details about the Murray-Linder function can be found in Talanta, 1985, 32, 483-489.
The redox state of a system at equilibrium is often characterised in terms of the quantity 'Eh' or its equivalent 'pe'. 'Eh' is the so-called redox potential of the system (in mV) and 'pe' is the negative log of the (hypothetical) concentration of the aquated electron. 'pe' is related to 'Eh' by the Nernst equation and is, in many ways, analogous to 'pH'. In practice, 'Eh' and 'pe' are determined by the predominant 'redox couple' in the chemical system, i.e. the ratio of the concentrations of the two most important species in redox equilibrium, in combination with the standard reduction potential of the corresponding redox couple. For example, at pH=7 and pe=5 (or Eh=296mV), the concentrations of species would be such that [Fe+2]=2.7e-4 and [Fe+3]=2.5e-12 given the the standard potential of the reaction 'Fe+3 + e-1 = Fe+2' is 0.771V (which is equivalent to lgK=13.0). Experimental measurement of Eh is generally highly problematic. This is mainly because of the irreproducibility engendered by physical and chemical fouling of the electrodes. Solutions containing sulfides (including -SH groups on proteins) are especially difficult. You also need to correct the raw potential readings for the effect of the reference electrode (SUBTRACT 222mV for Ag+/AgCl(s) and 267mV for calomel).