(2) Reminders and more examples of complexes using the technical terms described in (1)

A complex is formed by the combination of a central metal ion surrounded by, and bonded to, neutral molecules or ions acting as 'ligands' (bits stuck on or appendages).

• If you have already read Appendix 1. you should note that it is riddled with hydrated aqueous complex ions and the central metal ion does NOT have to be a transition element.

• The two ligands involved were H₂O and OH^–.

A ligand is an atom, ion or molecule which can act as an electron pair donor (Lewis base) and usually forms a dative covalent or 'co–ordinate' bond with the central metal ion.

• The lone pair donation is usually from an O, N or halogen atom of the ligand in this covalent co–ordinate bonding.

• The central metal atom or ion acts as a Lewis Acid, that is, an electron pair acceptor from the ligand by way of vacant 3d, 4s, 4p  and even 4d orbitals for the 3d–block transition elements.

• The ligand acts as a Lewis Base, that is, an electron pair donor e.g.

  • neutral ligands like H₂O: (water molecule, aqua in complex name)

    e.g. the familiar blue hexaaquacopper(II) [Cu(H₂O)₆]²⁺

    it has an octahedral shape, co-ordination number 6, charge 2+

    Since the ligand is neutral, the overall charge is the same as that of the central transition metal ion.

• or :NH₃ (ammonia molecule, ammine in complex name, neutral ligand)

• • e.g. the hexaamminenickel(II) ion [Ni(NH₃)₆]²⁺

  • it has an octahedral shape, co-ordination number 6, charge 2+

  • or the diamminesilver(I) ion [Ag(NH₃)₂]⁺  or   [H₃NAgNH₃]⁺

  • which is a linear shape with a co-ordination number of 2, charge +

  • Note that and indicate the two co-ordinate bonds and the specific direction of electron pair donation to form the dative covalent bonds between the ligands and the central metal atom or ion.

  • The uncharged water and ammonia are similar sized ligand molecules and smaller than the chloride ion below.

  • The size of ligands can have a bearing on the resulting shape of the complex ion.

  • Again, these two other examples of neutral ligands mean the overall charge on the complex is the same as the central metal ion.

• The carbon monoxide molecule can act as a neutral ligand electron pair donor.

• • e.g. in the neutral complex, nickel carbonyl Ni(CO)₄

  • it has a tetrahedral shape, co-ordination number 4, no overall electrical charge

  • Here you have a neutral central atom (not an ion), a neutral ligand, so overall an electrically neutral complex.

• and negatively charged ligands like :OH^– (hydroxide, hydroxo or hydroxy in complex name)

• • e.g. the hexahydroxochromate(III) ion [Cr(OH)₆]^3–

  • it has an octahedral shape, co-ordination number 6, overall charge 3- (+3 -6)

• or Cl^– (chloride ion, chloro in complex name)

• e.g. the tetrachloronickelate(II) ion, [NiCl₄]^2–

• it has a tetrahedral shape, co-ordination number 4, overall charge 2- (+2 -4 = -2)

• or the neutral diamminedichloroplatinum(II), [Pt(NH₃)₂Cl₂]

• which has co-ordination number of 4 and a square planar shape

• 

• no overall charge, 0 (+2 -2)

• Note the presence of two different monodentate ligands and two E/Z isomers!

• and :CN^– (cyanide ion, cyano in complex name).

  • e.g. the hexacyanoferrate(II) ion [Fe(CN)₆]^4–

  • it has an octahedral shape, co-ordination number 6, overall electrical charge of 4- (+2 -6)

• All six ligands mentioned above are monodentate ligands, forming one bond each with the central metal atom or ion.

• Complex ions undergo ligand exchange reactions (ligand displacement or ligand substitution reactions) e.g.

• [Cu(H₂O)₄(OH)₂]_(s) +  4NH_3(aq)   [Cu(NH₃)₄(H₂O)₂]²⁺_(aq) +  2OH^–_(aq)  +  4H₂O_(l)

• [Co(H₂O)₆]²⁺_(aq)  +  4Cl^–_(aq)   [CoCl₄]^2–_(aq)  +  6H₂O_(l) 

  • There are lots more described on the separate transition metal pages

  • e.g. Iron, Cobalt, Nickel & Copper

• I've deliberately included in the examples above, the most typical monodentate ligands you will come across and the shapes and co-ordination numbers you are also most likely to encounter.

(3) A simplified view of the bonding in an octahedral complex ion

• More on these examples and others below. (More details on molecule/ion shapes)

• 

• A an example of the bonding in a complex ion is shown in the above diagram. The negative cyanide ion is a monodentate ligand (forms one bond per ligand) and donates an electron pair into a vacant 3d, 4s or 4p orbital in the iron(III) ion to form six dative covalent bonds.

• The resulting ion has the formula [Fe(CN)₆]^3–, the overall charge of 3– is the aggregate of 6– (cyanide ions) plus 3+ (iron ion)

• The co–ordination number of 6, which means there are 6 central metal ion – ligand bonds. It doesn't necessarily mean six ligands, you can get a co–ordination number of 6 from three co–ordinated bidentate ligands (2 bonds per ligand), two tridentate ligands and from EDTA just one ligand can form 6 dative covalent bonds with a central metal ion.

• More on this below.

  • The most common complex ion you will come across is the hexaaqua cation of many metals.

  • It has the general formula [M(H₂O)₆]ⁿ⁺

  • n, the charge on the central metal ion and hence the overall charge on the complex ion n is usually 2 or 3

  •  e.g. n = 2 for titanium(II), vanadium(II), iron(II), cobalt(II), nickel(II), copper(II) and also the Group 2 alkaline Earth metals magnesium, calcium etc.

  • and n is 3 for scandium, titanium(III), vanadium(III), chromium(III), iron(III), cobalt(III) and also aluminium from Group 3.

  • The six neutral water ligands form 6 dative covalent bonds with the central metal ion because the bonding pair of electrons comes from donation of a lone pair from the oxygen atom of the water molecule.

  • Therefore the co–ordination number is 6 and it has a symmetrical octahedral shape.

  • The O–M–O bond angles are all 90^o or 180^o.

The ligand may attach itself by one or more bonds. The suffix '...dentate', prefixed by mono/uni/bi/ploy/multi e.g. monodentate (unidentate), bidentate, or polydentate (multidentate) is used to denote the number of bonds each ligand makes with the central metal ion.

The total number of ligand bonds to the central metal ion is called the co–ordination number.

• It is not the number of ligands, unless it is a monodentate ligand.

• There is no firm rules relating shape to a particular ligand.

• The six ligands don't have to be the same

•  e.g. the cis/Z isomer of [CrCl₂(H₂O)₄]⁺ complex ion.

• ... which is the dichlorotetraaquachromium(III) ion. This octahedral complex with a co–ordination number of 6, and note this has an overall ion charge of (2 x – from 2Cl^–) + (3+ from Cr³⁺) = +, water is an electrically neutral ligand ...

• ... and in equations the complex ion would be written as [Cr(H₂O)₄Cl₂]⁺

(4) More examples of unidentate/monodentate ligands:

•   and 

• e.g. above are shown two complexes with electrically neutral ligands: water H₂O:, ammonia :NH₃ and primary aliphatic amines like butylamine CH₃CH₂CH₂CH₂NH₂, 

• These ligands often form octahedral shaped complexes with a co–ordination number of 6.

• e.g. negative ligands: chloride Cl^– and cyanide CN^– ions (lone pair donors)..

• Above is the structure of the complex ion hexacyanocobaltate(III), [Co(CN)₆]^3-.

• The chloride ion Cl^{– can} form the tetrahedral e.g. the tetrachlorocuprate(II) complex ion ...

• [CuCl₄]^2–, note the overall charge is (2+) + (4 x –) = 2– and the co–ordination number of is 4.

• The chloride ion can be too bulky to form an octahedral complex or a square planar complex, though there is no firm rules relating complex shape to ligand.

• and CN^– square planar e.g. the tetracyanonickelate(II) complex ion ...

• [Ni(CN)₄]^2–, note the overall charge is (2+) + (4 x –) = 2– and the co–ordination number is 4.

  • Note that [Cu(H₂O)₄]²⁺, is in the hydrated salt CuSO₄.5H₂O, the tetraaquacopper(II) ion, with the less bulky water molecule ligand, forms a blue square planar complex, whereas with the larger chloride ion, a tetrahedral complex is formed.

  • The 5th water molecule hydrogen bonds between the copper complex ions in the blue crystals.

• A linearshaped complex is formed between a silver ion the ligands ammonia or cyanide.

  • cationic [H₃N–Ag–NH₃]⁺  and anionic [NC–Ag–CN]^–

• [Ag(NH₃)₂]⁺ is formed in 'ammoniacal' silver nitrate solution used in the test for aldehydes.

  • The diamminesilver(I) ion has co–ordination number of 2 and an overall charge of a single + because the ammonia molecule is an electrically neutral ligand.

(5) Examples of bidentate ligand complexes - chelation process of complex formation

• Examples of bidentate ('two toothed') ligands:

  • Neutral bidentate ligands: diamines like 1,2–diaminoethane (ethane–1,2–diamine) H₂NCH₂CH₂NH₂ (co-ordinate bonds via lone pair on the nitrogen atom :N). Note that this is an electrically neutral ligand like ammonia.

  • Negative bidentate ligands: ethanedioate ion C₂O₄^2–, (bonds via lone pair on the :O^–). The L represents where the dative covalent bond forms, L-L represents the bidentate ligand.

    • is derived from ethanedioic acid (oxalic acid), , so two co-ordinate bonds are formed by lone pair donation of electrons from two oxygen atoms.

  • shows three bidentate ligands co–ordinated to a central metal ion (co–ordination number 6, 'octahedral' in bond arrangement).

    • The L-M-L triangular bond system is known as a chelate ring.

  • Examples: [Cr(H₂NCH₂CH₂NH₂)₃]³⁺, H₂NCH₂CH₂NH₂ is often represented in shorthand by en,

    • and the chromium(III) complex ion is then simply written as [Cr(en)₃]³⁺.

    • All aliphatic linear diamines can act as bidentate ligands.

  • Bidentate ligands are the first of what are called polydentate ligands and such complexes are sometimes called chelates from the Greek for 'crab's claw' and the complex formation described as a chelation process.

(6) Examples of multidentate (>2) or polydentate ligands - chelation complex formation

• More examples of multidentate or polydentate ligands:

  • EDTA is an acronym abbreviation for the old name of EthyleneDiamineTetraAcetic acid a

  • It used in the form of its disodium hydrated salt which is soluble in water (structure below).

  • (Na^+-OOCCH₂)(HOOCCH₂)NCH₂CH₂N(CH₂COOH)(CH₂COO^-Na⁺).2H₂O

  • In solution it forms the chelating agent which can be considered to have the following structure

  • (^-OOCCH₂)₂NCH₂CH₂N(CH₂COO^-)₂  (diagram below of the 6 electron pair donor sites)

• The anion from the sodium salt of EDTA is often shown as EDTA^4– (for simplicity) and forms six co-ordinate bonds with a central metal ion and tends to displace most other ligands, mainly due to the increase in entropy

  • e.g. for these nickel(II) complex ions

  • [Ni(H₂O)₆]²⁺_(aq) + EDTA^4–_(aq) [Ni(EDTA)]^2–_(aq) + 6H₂O_(l)

  • [Ni(NH₃)₆]²⁺_(aq) + EDTA^4–_(aq) [Ni(EDTA)]^2–_(aq) + 6NH_3(aq)

  • The EDTA anion will displace most ligands from most transition metal ions.

• The haemoglobin (haem) molecule acts as a multi/polydentate ligand with iron(II) ions in blood chemistry.

  • in an extremely simplified form the structure is and iron(II) complex: [protein–Fe^II–O₂].

  • This would be referred to as oxyhaemoglobin.

  • The chlorophyll molecules also involves a polydentate ligand complex.

  • See The uv-visible absorption spectrum of chlorophyll, photosynthesis (porphyrin pigment)

  • and The uv-visible absorption spectrum of haemoglobin (porphyrin pigment)

  • for their structure

• Ligand displacement reactions

• One ligand can replace another depending on the relative bond strengths in a reaction called a ligand exchange reaction.

• When a bidentate or polydentate ligand is added to a pre–existing complex of monodentate ligands, it is highly likely a more stable complex will be formed.

  • This called the chelate effect, and the process is called chelation.

  • The overall enthalpy changes in breaking and making these co-ordinate bonds is not that great, so why the overpowering effect of bidentate and polydentate ligand molecules?

  • The principal reason for this, (ignoring bond strengths/energies), is the positive entropy change accompanying the 'release' of 4 or 6 small molecules which offer a greater variation of ways of arranging the particles or energy distribution.

• If the ligands are easily exchanged, the complex is described as 'unstable' and if the ligands are more strongly bound, the complex would be described as stable.

• Complex ion stability is also related to the oxidation state of the transition metal in the presence of a particular ligand.

• See Appendix 3. for more on complex ion shape and isomerism.

• See Appendix 5. for more on electrode potentials, oxidation state and complex ion stability.

• See Appendix 8. for more on complex ion stability, entropy changes and stability equilibrium constants

• See the individual transition metal pages below for examples of ligand exchange reactions.