COORDINATION CHEMISTRY

Introduction:

• Complex formation is a fundamental chemical process in natural waters - complexes modify metal species in solution, reducing activity of free ion, changing solubility, mobility, and toxicity.
• Surface chemistry can also be interpreted from a coordination chemistry approach, and processes such as ion exchange, sorption, and weathering are really ligand exchange reactions.

Fundamentals:

1. Coordination compounds are complexes that consist of one or more central atoms or ions with one or more attached molecules
2. The central atom is called a metal, and is a Lewis Acid;
3. The attached ions are ligands, and are Lewis Bases;
4. The total number of attachments between metal and ligands is the coordination number
5. Ligands attach to the metal via coordinate covalent bonds, i.e., the electrons of the bond is entirely derived from the ligand.

• Must involve electron pair not otherwise utilized in bonding withing the ligand.   The electron pair is equally shared, but controlled by the ligand.
• The ligand is a lewis base, and acts as an electron pair donor.  The metal is the electron acceptor and a lewis acid.

6. Most metals form cationic ions in solutions (electron acceptors), and participate in coordination complexes.
7. The strength of the coordination interaction is a function of the ability of a ligand to donate or share their available pair(s) of electrons.
8. Coordination number of a metal ion may vary from one complex to another.

Fe(H₂O)₆⁺³ ; Fe(CN)₆^-3 ;' Fe(Cl)₄^-

9. Water is also a (weak) Lewis Base, and can act as a ligand.   Aqueous ions are almost always initially coordinated with water, and any other complex actually forms by ligand exchange between water and the new ligand, i.e.:

Fe(H₂O)₆²⁺ + Cl^- <----> Fe(H₂O)₅Cl⁺ + H₂O

10. . Ligands that attach at a single point are monodentate. Examples are Cl^-, CN^-, OH^-

• Multidentate ligands chelate the metal, Ex. SO₄^2-

EDTA, for example,. forms 6-dentate complexes:

11. Polynuclear complexes form when there is more than one central ion or molecule:

Ion Pairs

• Temporary electrostatic complexes where each ion retains their coordinated waters, but they are still close enough to effectively form a complex between the metal and the ligand. 
• Usually form beween Hard cations and anions
• Another term is outer sphere complex.  These complexes are difficult to analyze directly, and are normally estimated from coulombic attraction and Bejerrum's ion association model;
• The log Kstab of ion pair complexes range from ~0 to 4, depending on charge density
• There can be ion pair or coordinative complexes for the same pair of ions, with an equilibrium established between them.  The major ion pairs in water are:

CaHCO₃⁺, ; CaSO₄°; CaOH⁺; MgCO₃°; MgSO₄°

• The degree of interaction is only slightly greater than the electrostatic interactions accounted for by the DeBye-Huckel theory, but this is enough to change activities.
• Ion pairs become more important in saline waters and brines, where the interactions must be specifically accounted for using the Pitzer Equations.

Reaction Rate:

• The rates of coordination reactions are characterized by the terms labile (very fast reactions) and inert (very slow reactions).
• These terms do not necessarily indicate the stability of the complex, only the kinetics.
• Ligand exchange involving water is usually fast, on the order of seconds

Complex Stability

• Equilibrium constants for complexes are usually stated for reactions written in the direction of complex formation, that is a stability constant:

Ligand + central metal ion -----> Complex

• The K_eq of the reaction is the stability or formation constant.    Large values of stability constants indicate stable complexes.  If the equilibrium constant is stated for the dissociation of the complex, it is called a dissociation or and instability constant.

HSAB

• Complex stability can be predicted in a general sense by the HSAB appraoch, or Hard soft acid bases rules.  The basic premise is that an individual cation "prefers" to to bond with one type of ligand over another

1. Type A metals tend to be sperically symetrical, with intert gas electron configurations, and are known as hard spheres, or hard acids. Examples are:

Alkalis; Na⁺, K⁺ (Weakest interaction, rarely considered)

Alkaline earths; Ca⁺², Mg²⁺;

Al⁺³; Th⁺⁴; Si⁺⁴

These cations tend to complex F^- and ligands containing oxygen;

CO₃^2-; OH^-; Borate

2. B-type cations have more easily polarized electron shells, form covalent complexes, and are termed soft acids, e.g.:

Ag⁺; Zn⁺²; Hg⁺²; Pb⁺², Sn⁺².

These metals tend to form complexes with soft ligands containing I, S, P, and N.

NH₄⁺; S^-; PO₃^-3 (Phosphite)

• Within these groups of metal ions complex stability is proportional to the charge of the metal ion - alkali metals are the weakest, and only complex with water, while transition metals complex with a variety of ligands.

  Irving Williams Series

  Mn2+ < Fe2+ < Co2+ < Ni2+ < Cu2+

1. ClO₄^- and NO₃^- rarely complex metals, while phosphate, hydroxide, and carbonate are potent complex formers.
2. Complexes containing 5 and 6 membered rings are the most stable, while four or 7+ membered rings are less stable.
3. Stability increases with the number of attachment points, so multiligand complexes are more stable than monodentate, called the chelate effect.

• When a ligand forms complexes with two metal ions at equal concentration in a solution, the ligand will form complexes preferentially with the metal ion that produces the complex with the larger stability constant.  Example: EDTA titration of hardness

  Ca²⁺ + EDTA^4- <----> (EDTA-Ca)^2-:    K = 10^+10.7

  Mg²⁺ + EDTA^-2 <----> (EDTA-Mg)^-2:    K = 10^+8.7

Since the stability constant of the calcium complex is significantly larger than the Mg complex, the reaction with calcium proceeds to completion before any reaction with magnesium:

!!remember, as metal is complexed its concentration decreases, changing the releation to the equilibrium constant.

• When a ligand that forms a very stable complex with a metal ion is added to a solution of the metal ion complexed with another ligand in a less stable complex, the stronger complexing agent (ligand) will tend to decompose the existing complex and withdraw the metal ion into a complex with itself (assuming the weaker complex is labile):   Example, EDTA Hardness titration:



A solution containing both Mg and Ca is dosed with the complexing agent Erichrome Black T (EBT), which forms a complex with Mg²⁺.  The EBT-Mg complex (K = 10⁺⁷) is weaker than the EDTA-Mg complex (K=10^+8.7)

The color of the EBT-Mg complex is red, and the color of uncomplexedd EBT is blue. A solution with excess EBT is therefore purple. Adding EDTA, the color of the solution will change from Purple to pure blue as the Red EBT-Mg complex is consumed.

Stability Constant Expressions

• For complexes that contain more than one ligand or central metal ion, there are two ways of writing stability constants.
1. Stepwise formation constants are equilibrium constants for the reaction in which the central ion consecutively adds one ligand.
   

Hg²⁺ + Cl^- <--> HgCl⁺;      log K₁ = 7.15

HgCl⁺ + Cl^- <---> HgCl₂°:      log K₂ = 6.9

HgCl₂° + Cl^- <---> HgCl₃^-      log K₃ = 2.0

where K_i is the equilibrium constant for the reaction in which the complex containing i ligands is formed from the complex containing i-1 ligands

2. An Overall formation constant is the equilibrium constant for the reaction in which the central metal ion combines with all of the ligands necessary to form a specific complex:
   

Hg²⁺ + Cl^- <---> HgCl      log b₁ = 7.15

Hg²⁺ + 2Cl^- <---> HgCl₂      log b₁ = 14.05

Hg²⁺ + 3Cl^- <---> HgCl₃      log b₁ = 16.05

Hg²⁺ + 4Cl^- <---> HgCl₄      log b₁ = 16.75

where b_i is the equilibrium constant for the reaction in which the complex containing i ligands is formed from the metal ion and i ligands.

• Either type of constant can be used, but not together.  It is probably safer to use the stepwise formation constants, but the books will usually have the overall.
• The general equation for these two then are:
  



pH Effect

• The addition of a protonated ligand introduces the complexity of what to do with the proton, as the ligand usually loses the proton during complexation (the proton often occupies the site of metal chelation...). This produces a more complex form of the equation:
  



**Remember that this reaction is reversible, so if the pH decreases (Proton activity increases), this type of complex will decompose.

· General Rule 1: The tendency of metal-ion solutions to protolyze (hydrolyze) increases with dilution and with decreasing pH

· General Rule 2: The fraction of polynuclear complexes in a solution decreases with dilution.

Organic Complexes

• These complexes are very difficult to predict, because of the complexity of the natural organic compounds.
• These complexes can act as independent ligands, forming multi-dentate complexes, or they can form "ternary complexes, especially on surfaces
• Iron, in particular, can be associated with organics as Fe(III), with the apparent solubility of iron exceeding the equilibrium by 3 orders of magnitude. Aluminum as well.

 

Hydrolysis of metal ions:

• In aqueous solution, free metal ions are complexed with water (aquated complexes), forming hydration shells. The metal carries the original charge however, since water is a neutral ligand that does not have to deprotonate to chelate.
• The interaction of these hydrated ions with acids and bases is a ligand exchange reaction called hydrolysis

  Al(H₂O)₆⁺³ + H₂O- <----> Al(H₂O)₅OH²⁺ + H₃O⁺

• So in a stepwise replacement of water molecules with hydroxyls, the loss of a proton into solution means that aluminum acts as both a Lewis and Bronsted acid.
• Hydrolysis is when water as a reactant is split up into proton and hydroxide, vs hydration, where water maintains itself.  A similar reaction would be:

  Al(H₂O)₆⁺³ + OH^- <----> Al(H₂O)₅OH²⁺ + H₂O

• So the complete reaction sequence of aluminum in water would be:

  Al(H₂O)₆⁺³ + H₂O <----> Al(H₂O)₅OH²⁺ + H₃O⁺              K_1,Al =

  Al(H₂O)₅OH²⁺ +H₂O <---->Al(H₂O)₄(OH)₂⁺ +H₃O⁺ K_2,Al =

  Al(H₂O)₄(OH)₂⁺ + H₂O <---->Al(H₂O)₃(OH)₃ +H₃O⁺ K_3,Al =

  Al(H₂O)₃(OH)₃^- + H₂O   <---->*Al(OH)₄^- +H₃O⁺            K_4,Al =

  *There is some evidence that Al(OH)₄^- assumes a tetrahedral geometry, without additional waters of hydration.