Doc Brown's Chemistry  Advanced Level Inorganic Chemistry Periodic Table Revision Notes – Transition Metals

 

 Appendix 2 Complexes – introduction: ligands, bonding, co–ordination number and charge on complex ions

What is a complex ion? What is a ligand? What do the terms monodentate ligand, bidentate ligand and polydentate ligand mean? What is the co–ordination number of a complex ion?  The structure of transition metal (3d–block) complexes is described with displayed formula diagrams and explainations include the formation of central metal ion – ligand dative covalent bonds. What shapes can complexes be? e.g. octahedral, tetrahedral, square planar and linear examples are presented.

(c) doc b GCSE/IGCSE Periodic Table Revision Notes * (c) doc b GCSE/IGCSE Transition Metals Revision Notes

INORGANIC Part 10 3d block TRANSITION METALS sub–index: 10.1–10.2 Introduction 3d–block Transition Metals * 10.3 Scandium * 10.4 Titanium * 10.5 Vanadium * 10.6 Chromium * 10.7 Manganese * 10.8 Iron * 10.9  Cobalt * 10.10 Nickel * 10.11 Copper * 10.12 Zinc * 10.13 Other Transition Metals e.g. Ag and Pt * Appendix 1. Hydrated salts, acidity of hexa–aqua ions * Appendix 2. Complexes & ligands * Appendix 3. Complexes and isomerism * Appendix 4. Electron configuration & colour theory * Appendix 5. Redox equations, feasibility, Eø * Appendix 6. Catalysis * Appendix 7. Redox equations * Appendix 8. Stability Constants and entropy changes * Appendix 9. Colorimetric analysis and complex ion formula * Appendix 10 3d block – extended data * Appendix 11 Some 3d–block compounds, complexes, oxidation states & electrode potentials * Appendix 12 Hydroxide complex precipitate 'pictures', formulae and equations

Advanced Level Inorganic Chemistry Periodic Table Index * Part 1 Periodic Table history * Part 2 Electron configurations, spectroscopy, hydrogen spectrum, ionisation energies * Part 3 Period 1 survey H to He * Part 4 Period 2 survey Li to Ne * Part 5 Period 3 survey Na to Ar * Part 6 Period 4 survey K to Kr and important trends down a group * Part 7 s–block Groups 1/2 Alkali Metals/Alkaline Earth Metals * Part 8  p–block Groups 3/13 to 0/18 * Part 9 Group 7/17 The Halogens * Part 10 3d block elements & Transition Metal Series * Part 11 Group & Series data & periodicity plots * All 11 Parts have their own sub–indexes near the top of the pages

Appendix 2. Complexes – introduction: ligands, bonding, co–ordination number and charge on complex ions

  • 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 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 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, aqua in complex name) or :NH3 (ammonia, ammine in complex name) and negatively charged ligands like :OH (hydroxide, hydroxo in complex name), Cl (chloride ion, chloro in complex name) and :CN (cyanide ion, cyano in complex name).

    • ...

      • Advanced Inorganic Chemistry Page Index and LinksA 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 90o or 180o.

  • 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.

    • Advanced Inorganic Chemistry Page Index and LinksIt 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. ...

      • ... 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]+

  • Examples of unidentate/monodentate ligands:

    • 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 octahedralshaped complexes with a co–ordination number of 6.

    • e.g. negative ligands: chloride Cl, cyanide CN,

    • The chloride ion Clforms the tetrahedrale.g. the tetrachlorocuprate(II) complex ion ...

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

    • 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.

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

    • neutral ligands: diamines like 1,2–diaminoethane (ethane–1,2–diamine) H2NCH2CH2NH2 (bonds via lone pair :N).

    • Advanced Inorganic Chemistry Page Index and Linksnegative ligands: ethanedioate ion C2O42–, (bonds via lone pair on the :O). The L represents where the dative covalent bond forms.

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

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

      • and the complex simply written as [Cr(en)3]3+.

    • 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.

  • More examples of multi/polydentate ligands:

    • EDTA4– (old name 'EthyleneDiamineTetraAcetic acid') forms six bonds with a central metal ion and tends to displace all other ligands.

      • [Ni(NH3)6]2+(aq) + EDTA4–(aq) [Ni(EDTA)]2–(aq) + 6NH3(aq)

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

      • in an extremely simplified form the structure is: [protein–FeII–O2]

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

    • The principal reason for this, (ignoring bond strengths), 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 (Kstab).


Scandium * Titanium * Vanadium * Chromium * Manganese * Iron * Cobalt * Nickel * Copper * Zinc * Silver & Platinum

Introduction 3d–block Transition Metals * Appendix 1. Hydrated salts, acidity of hexa–aqua ions * Appendix 2. Complexes & ligands * Appendix 3. Complexes and isomerism * Appendix 4. Electron configuration & colour theory * Appendix 5. Redox equations, feasibility, Eø * Appendix 6. Catalysis * Appendix 7. Redox equations * Appendix 8. Stability Constants and entropy changes * Appendix 9. Colorimetric analysis and complex ion formula * Appendix 10 3d block – extended data * Appendix 11 Some 3d–block compounds, complexes, oxidation states & electrode potentials * Appendix 12 Hydroxide complex precipitate 'pictures', formulae and equations

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