Editing Coordination complex (section)

==Stability constant==
{{Main|Stability constants of complexes}}

The affinity of metal ions for ligands is described by a stability constant, also called the formation constant, and is represented by the symbol K<sub>f</sub>. It is the [[equilibrium constant]] for its assembly from the constituent metal and ligands, and can be calculated accordingly, as in the following example for a simple case:

:xM <sub>(aq)</sub> + yL <sub>(aq) </sub> {{eqm}} zZ <sub>(aq)</sub>

:<math>K_f = \frac{[\text{Z}]^z}{[\text{M}]^x[\text{L}]^y}</math>

where : x, y, and z are the [[stoichiometric]] coefficients of each species. M stands for metal / metal ion , the L for Lewis bases , and finally Z for complex ions. Formation constants vary widely.  Large values indicate that the metal has high affinity for the ligand, provided the system is at equilibrium.<ref>{{Cite web|url = http://www2.ucdsb.on.ca/tiss/stretton/CHEM2/solubil8.htm|title = Complex Ion Equilibria|access-date = 2015-04-22|archive-date = 2015-03-08|archive-url = https://web.archive.org/web/20150308110908/http://www2.ucdsb.on.ca/tiss/stretton/CHEM2/solubil8.htm|url-status = dead}}</ref>

Sometimes the stability constant will be in a different form known as the constant of destability. This constant is expressed as the inverse of the constant of formation and is denoted as K<sub>d</sub> = 1/K<sub>f</sub> .<ref>{{Cite web|url = http://www2.bakersfieldcollege.edu/wcooper/Chem%20B1B%20Notes%20Fall_09/Chap_17_Solubility_Equilibrium_Notes.pdf|title = Solubility and Complex-ion Equilibria|last = Stretton|first = Tom|access-date = 2015-04-22|archive-date = 2018-09-20|archive-url = https://web.archive.org/web/20180920175811/http://www2.bakersfieldcollege.edu/wcooper/Chem%20B1B%20Notes%20Fall_09/Chap_17_Solubility_Equilibrium_Notes.pdf|url-status = dead}}</ref> This constant represents the reverse reaction for the decomposition of a complex ion into its individual metal and ligand components. When comparing the values for K<sub>d</sub>, the larger the value, the more unstable the complex ion is.

As a result of these complex ions forming in solutions they also can play a key role in solubility of other compounds. When a complex ion is formed it can alter the concentrations of its components in the solution. For example:

:Ag{{su|p=+|b=(aq)}} + 2{{nbsp}}NH<sub>3</sub> {{eqm}} Ag(NH<sub>3</sub>){{su|b=2|p=+}}

:AgCl<sub>(s)</sub> + H<sub>2</sub>O<sub>(l)</sub> {{eqm}} Ag{{su|b=(aq)|p=+}} + Cl{{su|b=(aq)|p=−}}

If these reactions both occurred in the same reaction vessel, the solubility of the silver chloride would be increased by the presence of NH<sub>4</sub>OH because formation of the Diammine argentum(I) complex consumes a significant portion of the free silver ions from the solution. By [[Le Chatelier's principle]], this causes the equilibrium reaction for the dissolving of the silver chloride, which has silver ion as a product, to shift to the right.

This new solubility can be calculated given the values of K<sub>f</sub> and K<sub>sp</sub> for the original reactions. The solubility is found essentially by combining the two separate equilibria into one combined equilibrium reaction and this combined reaction is the one that determines the new solubility. So K<sub>c</sub>, the new solubility constant, is denoted by:
:<math>K_c = K_{sp} K_f</math>
<!--==Specific metal complexes==

===Mercury===

The speciation, solubility, mobility, and toxicity of [[Mercury (element)|mercury]] within aquatic environments are strongly influenced by its [[Chemical Reaction|complexation]] with inorganic and organic [[ligands]]; most notable is mercury’s interaction with [[dissolved organic matter]] (DOM).  As a result, the speciation of mercury depends on the concentration of each ligand and the stability constants of mercury complexes they form.  For mercury, important inorganic ligands include hydroxide, chloride, and sulfide.  However, complexation by natural organic compounds often controls the [[aqueous geochemistry|biogeochemical]] cycling of mercury. For example, complexation with DOM may limit the availability of Hg<sup>2+</sup> for conversion to methylmercury or enhance the formation of elemental mercury (Hg<sup>0</sup>) from Hg<sup>2+</sup>, further lowering the availability of Hg<sup>2+</sup> for conversion to methylmercury.<ref name="Lamborg et al">Lamborg, C. H., C. Tseng, W. F. Fitzgerald, P. H. Balcom, and C. R. Hammerschmidt, 2003, Determination of the Mercury Complexation Characteristics of Dissolved Organic Matter in Natural Waters with “Reducible Hg” Titrations: Environmental Science and Technology, 37, 3316-3322.</ref><ref name ="Ravichandran et al">Ravichandran, M., 2004, Interactions between mercury and dissolved organic matter—a review: Chemosphere, 55, 319-331.</ref>

Mercury binding to dissolved organic matter is evaluated in terms of the [[stability constants of complexes|stability constants]] of Hg-DOM complexes.<ref name="Ravichandran et al" />  The complexation reaction between Hg<sup>2+</sup> and an organic ligand is assumed to take the following form:

Hg<sup>2+</sup> + L<sup>n−</sup> → 2 HgL<sup>(2-n)+</sup>

where K, the stability constant, is equal to:

K= [HgL<sup>(2-n)</sup>]/[Hg<sup>2+</sup>][L<sup>n−</sup>].

<ref name="Lamborg et al" /> Metal cations bind to the acid sites in organic matter, the most common of which include carboxylic acids, phenols, amines, alcohols, and thiols.  Mercury is classified as a type B metal and shows a preference for ligands of sulfur, the less electronegative halides, and nitrogen, over ligands containing oxygen.  Therefore, mercury is expected to preferentially bind with thiol and other sulfur-containing functional groups which, despite their relatively low abundance in DOM, far exceed the amount of mercury available in aquatic environments.<ref name="Ravichandran et al" /> Previous research has shown that binding of Hg<sup>2+</sup> by organic ligands does not occur immediately when both are present; instead, complexation tends to follow a pseudo-first-order rate constant with exponential decay.<ref name="Lamborg et al" /> 
 
Predicting mercury speciation due to Hg-DOM complexes is difficult due to organic matter heterogeneity, electrostatic effects, and the variation in stability constants; however, recent studies have shown conditional stability constants for mercury binding to be in the range of 10<sup>22</sup> – 10<sup>28</sup>.<ref name="Ravichandran et al" /> 
    
Competitive ligand exchange (CLE) experiments, in which Hg binding to natural organic ligands is measured in the presence of an added ligand whose complexation with Hg<sup>2+</sup> is well characterized, is used to determine the concentrations and apparent stability constants of natural mercury complexes.  One such experiment used both chloride and thiosalicyclic acid (TSA) as competing ligands; estimated conditional stability constants ranged from 10<sup>26.1</sup> to 10<sup>26.9</sup> for chloride competition and 10<sup>27.3</sup> to 10<sup>29.2</sup> for TSA competition.<ref name="Han et al">Han, S. and G. A. Gill, 2005, Determination of Mercury Complexation in Coastal and Estuarine Waters Using Competitive Ligand Exchange Method: Environmental Science and Technology, 39, 6607-6615.</ref>  This study showed that lower concentrations of binding sites with higher stability constants are more important for Hg speciation than higher concentrations of weaker binding sites.

Conditional stability constants vary with pH due to competition with protons for binding sites.<ref name="Han et al" />  They also vary with salinity and, contrary to what may be expected, are generally lower in freshwater than seawater.<ref name="Lamborg et al" /> In seawater, organic complexes must have very high K values in order to compete with chloride for mercury complexation.  In freshwater systems, hydroxide is the most common inorganic ligand, yet its concentration is low enough that organic ligands do not need to have a high K value to compete for mercury complexation.  In both environments, organic compounds dominate mercury complexation.<ref name="Lamborg et al" /><ref name="Han et al" />

In examining the role of DOM in mercury speciation under sulfidic environments, we see that inorganic sulfide also plays an important role within anoxic environments due to the very strong binding of sulfide with mercury.  HgS{{su|p=0|b=(aq)}}, Hg(S<sub>2</sub>H)<sup>−</sup>, Hg(SH){{su|b=2|p=0}}, and [[mercury sulfide|HgS<sub>(s)</sub>]] are likely to be the most important species.  Recent studies however, have shown that stability constants of Hg-DOM complexation are higher than those for mercury sulfide complexation.  These results then imply that organic matter can out complete sulfide for the complexation of mercury within anoxic environments.<ref name="Ravichandran et al" />

DOM has been shown to affect the bioaccumulation of mercury via photochemical reduction and methylation.  Reduction of Hg<sup>2+</sup> to Hg<sup>0</sup> by sunlight is a commonly occurring process, yet is enhanced in the presence of DOM.  Conversely, photolysis of dissolved organic carbon (DOC) can produce radical oxygen species including hydroxyl radicals (<sup>−</sup>OH), which in turn have been shown to oxidize Hg<sup>0</sup> to Hg<sup>2+</sup>.<ref name="Ravichandran et al" />   Conversion of ionic mercury (especially Hg<sup>2+</sup>) to methyl mercury is an important process, as methyl mercury is a neurotoxin and has been show to bioaccumulate within the food chain.  Mercury methylation is a microbially mediated process wherein bacteria assimilate neutrally charged mercury species through passive diffusion.  Complexation with DOC limits this uptake mechanism as the DOC molecules are too large to pass through cell membranes; however, at low pH, DOC is less negatively charged and therefore less likely to complex mercury, thereby making mercury more available for methylation.  Additionally, DOM-mediated reduction of Hg<sup>2+</sup> to Hg<sup>0</sup> would further limit the availability of mercury for methylation by leading to the mercury volatilization.    Direct methylation of mercury can also occur by reaction with humic and fulvic acids in DOM.<ref name="Ravichandran et al" />-->