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A level inorganic chemistry: 3d block-transition metals - explaining complex ions - dative bonding. ligands, coordination number

Periodic Table - Transition Metal Chemistry - Doc Brown's Chemistry  Revising Advanced Level Inorganic Chemistry Periodic Table Revision Notes

 

 Appendix 2 Complexes – introduction: ligands, bonding, co–ordination number and the charge and shape of 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 explanations 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



Three crucial definitions to learn in connection with complex formation:

A ligand is a molecule or ion that forms a co-ordinate (dative covalent) bond with a central transition metal atom or ion by donation of a pair of electrons.

A unidentate (monodentate) ligand forms one co-ordinate bond per ligand with the central metal atom or ion.

A bidentate ligand forms two co-ordinate bonds per ligand with the central metal atom or ion.

A polydentate ligand forms over two co-ordinate bonds per ligand with the central atom or ion.

A complex is the resulting central metal atom or ion (often a transition metal) surrounded by, and bonded to, a number of ligands e.g. often 2, 3, 4 or 6 (can be even 1).

The co-ordination number is the number of co-ordinate bonds to the central metal atom or ion of the specific complex e.g. most often 2, 4 or 6.

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 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)

        • 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 covalent bond between the ligand and the central metal atom or ion.

        • The uncharged water and ammonia are similar sized ligand molecules and bigger 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)

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

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

        • cisplatin, no overall charge, 0 (+2 -2)

        • Note the presence of two different monodentate ligands and two 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) + 4NH3(aq) rev [Cu(NH3)4(H2O)2]2+(aq) + 2OH(aq) + 4H2O(l)

      • [Co(H2O)6]2+(aq) + 4Cl(aq) rev [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.

    • 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 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.

    • 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. ...

      • ... 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+, 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.

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

    • negative bidentate ligands: ethanedioate ion C2O42–, (bonds via lone pair on the :O). The L represents where the dative covalent bond forms.

      • 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).

    • 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 co-ordinate bonds with a central metal ion and tends to displace all other ligands, mainly due to the increase in entropy.

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

    • 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–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.

    • 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 (Kstab).

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


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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