* Inorganic Chemistry Transition Metals Appendix 1 Hydrated salts, metal-aqua complex ions and their relative acidity Doc B

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 Doc Brown's Chemistry  Periodic Table Revision Notes - Part 10. 3d block - Transition Metals

 Appendix 1 Hydrated salts, metal-aqua complex ions and their relative acidity

Revision notes for GCE Advanced Subsidiary Level AS Advanced Level A2 IB Revise AQA GCE Chemistry OCR GCE Chemistry Edexcel GCE Chemistry Salters Chemistry CIE Chemistry revising courses for pre-university students (equal to US grade 11 and grade 12 and Honours/honors level courses)  GCSE Periodic Table * GCSE notes Transition Metals * EMAIL comment

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 1. Hydrated salts, metal-aqua complex ions and their relative acidity

(not necessarily just transition metal ions)

  • All metal ions in solution are 'associated' with water. The water molecules can also be weakly bonded or more strongly as a ligand to form a complex ion, and these can also present in solid 'hydrated' salts on crystallisation e.g.

    • FeSO4.7H2O(s), CoCl2.6H2O(s), CuSO4.5H2O(s) etc.

      • Iron(II) sulphate heptahydrate, cobalt(II) chloride hexahydrate and copper(II) sulphate pentahydrate.

      • The above crystals contain 7, 6 and 5 molecules of water of crystallisation respectively.

      • A hexa-aqua ion is present in the first two, [M(H2O6)]2+ (M = Fe, Co)

    • In the case of copper(II) sulphate, 4 water molecules are covalently bonded to form a square planar complex ion, [Cu(H2O)4]2+ and the 5th water molecule is hydrogen bonded to this ion and a neighbouring sulphate ion helping to hold the crystal lattice together.

      • However, this blue crystal lattice is readily broken down on heating, a classical demonstration of a reversible reaction, since the white anhydrous solid turns blue on adding water (a simple test for water.

      • CuSO4.5H2O(s) CuSO4(s) + 5H2O(g/l)

  • Lewis acid-base theory:

    • A base is an electron pair donor and an acid is an electron pair acceptor.

    • Ligands like water, can donate a pair of non-bonding electrons (lone pair) into a vacant orbital of a central metal ion and so dative covalent (co-ordinate) bonds hold a complex together.

    • The central metal ion with vacant bonding orbitals can act as a Lewis acid.

    • Ligands act as Lewis bases by electron pair donation to form the metal-ligand bond.

  • Bronsted-Lowry acid-base theory (essentially a sub-set of Lewis Theory)

    • For more details see Equilibria Part 5

    • A base is a proton acceptor.

      • This is via an electron lone pair on the base (a Lewis base is a lone pair donor).

      • e.g. NH3, HCO3-, OH- etc.

    • An acid is proton donor.

      • This involves a heterolytic breakage of an X-H bond (a Lewis acid is an electron pair acceptor).

      • e.g. HCl, HCO3-, H2SO4, CH3COOH etc.

  • Many hexa-aqa complex ions can undergo acid-base reactions with water to produce solutions of pH less than 7.

    • Usually group 2, 3 and transition metal ions.

    • The positive central metal ion polarises a water molecule, releasing a proton, H+.

    • In the deprotonation reaction the proton transfers to water and the overall charge on the complex falls by 1 unit since the H2O - H+ = OH-, i.e. one of the ligands is now a hydroxide ion.

    • In these reactions the hydrated ions act as Bronsted Lowry acids and water acts as a Bronsted-Lowry base.

  • e.g. [M(H2O)6]2+(aq) + H2O(l) (c) doc b [M(H2O)5(OH)]+(aq) + H3O+(aq) 

    • e.g. when M = Mn, Fe, Co, Ni, Cu, Mg etc. give very weak acid solutions with pH's just less than 7.

      • Ti(II), V(II) and Cr(II) M2+ ions are redox unstable in the presence of air, but theoretically their salts give very weakly  acid solutions, but, since they are usually prepared by zinc-acid reduction from higher oxidation states, its not a very relevant fact here.

  • e.g. [M(H2O)6]3+(aq) + H2O(l) (c) doc b [M(H2O)5(OH)]2+(aq) + H3O+(aq) 

    • e.g. when M = Ti, V, Cr, Fe, Al etc. give very weak acids solutions (but generally stronger than for M2+) of pH's in the 3-5 region.

    • In the presence of alkali, OH-, removing H3O+ ions, the  equilibrium moves more to the right and more protons are lost from the complex in stages until the hydroxide precipitate is formed e.g. for iron(III), chromium(III) or aluminium.

      • [M(H2O)6]3+(aq) + 3OH-(aq) [M(H2O)3(OH)3]0(aq) + 3H2O(l) 

    • Some of the M3+ hydroxides are amphoteric and dissolve in excess alkali (1.) or acid (2.) e.g. to eventually form for chromium(III) or aluminium, 1. the soluble hexa-hydroxo complex ion or 2. the original hexa-aqua ion.

      1. [M(H2O)3(OH)3]0(aq) + 3OH-(aq)  [M(OH)6]3-(aq)+ 3H2O(l) 

      2. [M(H2O)3(OH)3]0(aq) + 3H3O+(aq)  [M(H2O)6]3+(aq) + 3H2O(l)

  • As a general rule the greater the polarising power of the central metal ion, the lower the pH of the resulting aqueous solution, i.e. the acid-base equilibrium is shifted more to the right causing an increase in acidity of the solution.

    • This effect and process facilitated by the central metal ion on one water ligand molecule can be envisaged as

      • [M-O-H2]n+ ==> [M-O-H]n+-1 + H+ (proton transferred to a water molecule)

      • One of the O-H bond pairs is 'attracted' onto the oxygen atom by the electric field effect of the central metal ion of charge n+, allowing proton transfer to the base water.

    • Polarising power is a function of ionic charge/ionic radius ratio. therefore ...

      • the greater the metal ion charge (n+), the more acidic the solution,

      • and the smaller the radius of the metal ion, the more acidic the solution, so ...

    • both increasing charge, or decreasing the central cation radius intensify the electric field polarising effect on a water ligand.

  • The acidity of the hexaaquions M3+(aq) due to the polarising influence of the central highly charged M3+ ion accounts for the lack of stability/existence of e.g. aluminium carbonate, iron(III) carbonate or chromium(III) carbonate, whereas MgCO3 , ZnCO3 and FeCO3 etc. with the less polarising M2+ ion exist. It also accounts for why you see bubbles of carbon dioxide when carbonates/hydrogencarbonates are mixed with aluminium chloride, iron(III) chloride or chromium(III) chloride solutions e.g.

    • 2[Fe(H2O)6]3+(aq) + CO32-(aq) 2[Fe(H2O)5(OH)]2+(aq) + H2O(l) + CO2(g)   

    • or [M(H2O)6]3+(aq) + HCO3-(aq) [M(H2O)5(OH)]2+(aq) + H2O(l) + CO2(g)

    • There several legitimate permutations based on these equations.

  • Amphoteric nature of ions such as Al3+ and Fe3+

    • In the above chemistry the acidic nature of the hexa-aqua ions was emphasised, BUT as soon as one proton has been lost the resulting complex ion can then act as a base.

    • e.g. in solutions of the weakly acidic  Fe3+ or Al3+ species with excess strong acid the hexa-aqua ion would predominate

    • [M(H2O)4(OH)2]+(aq) + 2H+(aq) (c) doc b [M(H2O)6]3+(aq)

    • and hydroxide precipitates will readily dissolve on adding an acid which is stronger than the hexa-aqua ion

    • e.g. Fe(OH)3(s) + 3H3O+(aq) ==> [Fe(H2O)6]3+(aq)

    • or Al(OH)3(s) + 3H3O+(aq) ==> [Al(H2O)6]3+(aq)

Advanced Inorganic Chemistry Page Index and Links

 

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

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