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 The Periodic Table Part 10

Part 10d "3d block - Transition Metals" revision notes

10d: Extra Data and extra hydroxide precipitate 'pictures'

 Advanced Level Inorganic Chemistry Revision notes

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 GCSE Chemistry revision notes * GCSE notes on Transition Metals * EMAIL query?comment

Part 10 3d block sub-index: 1. Introduction * 2. 3d-block data, general trends and character of Transition Metals * 3. Scandium * 4. Titanium * 5. Vanadium * 6. Chromium * 7. Manganese * 8. Iron * 9. Cobalt 10. Nickel * 11. Copper * 12. Zinc * 13. Other Transition Metals e.g. silver Ag or platinum Pt * Appendix 1. Acidity of hexa-aqua ions * Appendix 2. Complexes & ligands - the basics * Appendix 3. The shapes of complexes and isomerism * Appendix 4. Electron configuration and colour theory * Appendix 5. Redox equations, feasibility, calculating E^ø_reaction * Appendix 6. Catalysis - examples and theory * Appendix 7. Balancing redox equations * Appendix 8. Stability Constants of complex ions * Appendix 9. Colorimetry - quantitative analysis and determining the formula of a complex ion * Appendix 10. Preparation of complexes * Extra 3d block - Transition Metals data * Extra Hydroxide precipitate 'pictures' * Extra Electrode Potential Chart for 3d-block * Extra comparison of 3d-block formulae and oxidation states

Advanced Periodic Table Index * Part 1 A brief Periodic Table history * the modern Periodic Table * Part 2 Electronic structure of atoms : Spectroscopy and the H spectrum : Ionisation energies * Part 3 Period 1 survey : 1. Hydrogen : 2. Helium : Summary of  Period 1 : heavier element formation-stellar nuclear fusion * Part 7 s-block metals Gps 1/2 Alkali/Alkaline Earth Metals * Part 11 Group and Series data summaries and links to periodicity plots

Quick click to Introduction * Sc * Ti * V * Cr * Mn * Fe * Co * Ni * Cu * Zn * Ag/Pt etc.

Part 10d-1 DATA TABLE 2 - Extended SUMMARY FOR 3d BLOCK METALS

• 3d block data notes:
  1. Atomic and ionic radii are quoted in pm (1 picometre, 10^-12 m), 1000 pm = 1 nm (1 nanometre = 10^-9 m)
  2. na means 'not applicable' or 'not available'.
  3. The electronegativity values are from the Pauling scale.
  4. Ionic radii relate to the 'isolated' theoretical ion, NOT hydrated aqueous ions.
  5. Polarizing power is a measure of the ions charge density which has important chemical consequences e.g.
     • the bonding nature of metals with non-metals e.g. an ionic or a covalent MCl_n
     • and the acidity of the hydrated aqueous ion [M(H₂O)₆]ⁿ⁺_(aq).
     • The ion charge divided by the ionic radius of the 'isolated ion' is a 'reasonable' number scale for easy comparison of polarising power, and in the tables I've multiplied the charge/radius by 100 to make a suitable scale.
     • Obviously, the larger the charge, or the smaller the volume or radius, the greater the charge density or polarising power.
     • Note that as the oxidation state of the transition metal increases,
       • i.e. increase in charge if an ionic compound, the greater the polarising power of the cation, which increases the covalent character of the compound, exemplified by comparing iron(II) and iron(III) compounds or complex ions
         • e.g. FeCl₂ is essentially an ionic compound and FeCl₃ is covalent in character,
       • the greater the polarising power of the central metal ion, the greater the acidity of the hexaaqua ion
         • e.g. [Fe(H₂O)₆]³⁺ is more acidic than [Fe(H₂O)₆]²⁺
           • Going from iron(II) to iron(III) involves an increases in cationic positive charge and decrease in radius of the 'isolated' central metal ion. The decrease in radius is bound to result from the same nuclear charge of 26+ 'pulling in' 24 and 23 electrons respectively, i.e. less electron density in the same quantum level, less space occupied.
           • For more details see Transition Metals Appendix 1 Acidity of hexaaqua-ions
     • Relative polarising power of Groups 1-3 ions for comparison with the 3d block ions above:

     • Group of the Periodic Table	Metal ion and ionic charge	ionic radius/pm	relative polarizing power = 100 x charge / radius
       1	Na+	98	1.0
       2	Mg2+	78	2.6
       3	Al3+	60	5.0
       1	K+	133	0.75
       2	Ca2+	106	1.9

     • Note the substantial increase in polarizing power of the cations across Period 3 from sodium to aluminium as the ion charge increases and the ionic radius decreases. From the data from Groups and 1 and 2 you can see the polarising power of similarly charged cation decreases down a group as the ionic radius increases.
  6. -

Quick click to Introduction * Sc * Ti * V * Cr * Mn * Fe * Co * Ni * Cu * Zn * Ag/Pt etc.

Part 10d-2 A summary of some 3d-block compounds, complexes and oxidation states

Most are mentioned in the detailed individual element notes, but some have been added to illustrate other oxidation states you may not encounter on your course - but some good oxidation number practice!

Quick click to Introduction * Sc * Ti * V * Cr * Mn * Fe * Co * Ni * Cu * Zn * Ag/Pt etc.

Part 10d-3 Standard Electrode Potential Chart Diagram for the 3d-block elements

• Redox potential chart comments:

• All data quoted is for standard conditions i.e. 298K, 1 atm. pressure and 1 mol dm^-3 solutions of ions.

• Other than the solid metals, MnO₂ and FeO₄^2-, hydrogen gas, you can assume all ions are in aqueous media.

• Unless an oxyanion, oxocation or another ligand in a complex is indicated, you assume you are dealing with hexaaqua-metal ions (H₂O ligand only).

• Further comments below draw out some general patterns and other points of interest.

  • All except scandium (Sc³⁺), which is not that reactive to acids despite the relatively negative M/M³⁺ potential, form a hydrated M²⁺ ion either by reaction of the metal with acid or reduction of a higher oxidation state complex-compound.

  • The stable oxidation states in aqueous solution containing dissolved oxygen from air are for the hydrated ions ...

    • (only Sc³⁺), [TiO]²⁺, VO²⁺, Cr³⁺, Mn²⁺, Fe³⁺, Co²⁺ and Ni²⁺ (only Zn²⁺).

    • On the basis of the electrode potential chart above, the argument is simple. In neutral or acid solution the oxidising potential of the oxygen-proton-water system is +1.23V. Therefore any e.g. M³⁺/M²⁺ potential less positive than +1.23V will result in the oxidation of the lower oxidation state species to the higher oxidation state species in the presence of dissolved oxygen which is reduced to water.

      • Oxidation states higher than the stable ones tend to oxidise water liberating oxygen and as mentioned above, lower oxidation states tend to be reducing and liberate hydrogen from water. So the Mn³⁺/Mn²⁺ and Co³⁺/Co²⁺ potentials lie above +1.23V so Mn³⁺ and Co³⁺ will oxidise water and cannot be stable in acid solution.

    • Note that the +4 oxidation states of Ti and V exist as hydrated oxo-cations because the high polarising power of the highly charged central metal ion causes deprotonation (see Appendix 1. Acidity of hexa-aqua ions).

    • The rest are [M(H₂O)_n]^2+/3+ where n is usually 6, can be 4 for Cu and Zn.

  • Apart from iron, there is a tendency for the lower oxidation state to become increasingly more stable with increasing atomic number.

  • Higher oxidation states which are normally oxidising in aqueous solution can be stabilised by complexing e.g. compare the Co(II)/Co(III) potential when complexed with water (+1.82V) and with the ligand ammonia (+0.10V).

  • There are classic examples of disproportionation where an intermediate oxidation state species spontaneously changes into a higher and lower' oxidation state species e.g. the disproportionation reactions

    • Cu(I) ==> Cu(0) + Cu(II) and Mn(VI) ==> Mn(II) + Mn(VII).

    • These are described in detail, complete with electrode potential arguments for thermodynamic feasibility, under the respective metal.

  • How do you work out what will oxidise what? or what will reduce what?

    • How to work out the feasibility of reaction from electrode potential data is described in Appendix 5.

    • Using an electrode potential chart like the one above or a list of redox potentials the following rules apply.

      • To facilitate an oxidation, the half-cell potential of the oxidising agent must be less negative or more positive than the redox potential of the 'system' you wish to oxidise.

        • So using at the redox potential chart for example:

        • Dissolved oxygen will oxidise Co²⁺ to Co³⁺ in presence of ammonia - forms the amine complexes, but the hexaaqua complex of Co²⁺ is stable in the presence of oxygen if no ammonia present.

          • Co³⁺/Co²⁺ (H₂O ligand, E^Ø = +1.82V), O₂/H₂O (E^Ø = +1.23, less than +1.82 but more than +0.10V), Co³⁺/Co²⁺ (NH₃ ligand, E^Ø = +0.10V)

          • So [Co(H₂O)₆]²⁺ is stable in the presence of oxygen, but [Co(NH₃)₆]²⁺ will be oxidised to [Co(NH₃)₆]³⁺.

          • You could then further predict that [Co(H₂O)₆]³⁺ will oxidise water to oxygen!

      • To facilitate a reduction, the half-cell potential of the reducing agent must be more negative or less positive than the redox potential of the 'system' you wish to reduce.

        • So using the redox potential chart and the half-cell redox potential for I₂/I^- of +0.54V:

          • hexaaquairon(III) ions will be reduced by iodide ions because E^Ø for Fe³⁺/Fe²⁺ (H₂O ligand) is +0.77V i.e. the Fe³⁺ will oxidise the iodide ions rather than iodine oxidising the Fe²⁺ ions. [Fe(H₂O)₆]³⁺ is reduced to [Fe(H₂O)₆]²⁺, iron(III) to iron(II).

          • However if the ligand is the cyanide ion, then iodide ions will not reduce the Fe³⁺ cyanide ion complex but iodine would oxidise [Fe(CN)₆]^4- to [Fe(CN)₆]^3-, iron(II) to iron(III).

          • Fe³⁺/Fe²⁺ (H₂O ligand, E^Ø = +0.77V), I₂/I^- (E^Ø = +0.54, less than +0.77 but more than +0.36V), Fe³⁺/Fe²⁺ (CN^- ligand, E^Ø = +0.36V)

Quick click to Introduction * Sc * Ti * V * Cr * Mn * Fe * Co * Ni * Cu * Zn * Ag/Pt etc.

Part 10d-4 PICTURES of PRECIPITATES and COMPLEX FORMATION

Prior to, and with, excess reagent for selected 3d block aqueous ions.

They can be used as simple tests for identifying transition metal ions.

Text and formula summary of these reactions of 3d block ions is given below but there are more details in most cases in the notes for each metal for ...

[M(H₂O)₆]ⁿ⁺ where M = metal, n = 2 or 3 and Al³⁺ for comparison (See Chemical Tests for more precipitates)

... and many are useful simple tests to identify metal ions (ppt. = precipitate, gel. = gelatinous)

The comparison equations for the aluminium ion (NOT a 3d block metal)

• The addition of limited amounts of the bases sodium hydroxide or ammonia solution to an aluminium salt solution.

  • [Al(H₂O)₆]³⁺_(aq) + 3OH^-_(aq) ==> [Al(H₂O)₃(OH)₃]_(s) + 3H₂O_(aq)

  • A white gelatinous precipitate of aluminium hydroxide is formed.

    • Simplified equation: Al³⁺_(aq) + 3OH^-_(aq) ==> Al(OH)_3(s)

• The further addition of excess sodium hydroxide or ammonia solution.

  • With excess ammonia there is no effect, but with excess sodium hydroxide the aluminium hydroxide dissolves to form a soluble aluminate complex anion - amphoteric behaviour.

  • [Al(H₂O)₃(OH)₃]_(s) + 3OH^-_(aq) ==> *[Al(OH)₆]^3-_(aq) + 3H₂O_(aq)

    • Simplified equation: Al(OH)_3(s) + 3OH^-_(aq) ==> *[Al(OH)₆]^3-_(aq)

    • *The products will be an equilibrium mixture including [Al(H₂O)₂(OH)₄]^-_(aq) and [Al(H₂O)(OH)₅]^2-_(aq) too. You could write the equation in terms of forming these species too and any of the three possibilities should get you the marks.

  • To complete the 'amphoteric' picture of aluminium hydroxide consider it dissolving in mineral acids to form typical salts e.g. aluminium chloride, aluminium nitrate and aluminium sulphate.

    • Al(OH)_3(s) + 3HCl_(aq) ==> AlCl_3(aq) + 3H₂O_(l)

    • Al(OH)_3(s) + 3HNO_3(aq) ==> Al(NO₃)_3(aq) + 3H₂O_(l)

    • 2Al(OH)_3(s) + 3H₂SO_4(aq) ==> Al₂(SO₄)_3(aq) + 6H₂O_(l)

• The addition of sodium carbonate solution to an aluminium salt solution.

  • Bubbles of carbon dioxide and a white gelatinous precipitate of aluminium hydroxide are formed.

    • 2[Al(H₂O)₆]³⁺_(aq) + 3CO₃^2-_(aq) ==> 2[Al(H₂O)₃(OH)₃]_(s) + 3CO_2(g) + 3H₂O_(aq)

    • There several equation 'permutations' to represent this quite complicated reaction, so I've just composed one that shows the formation of both observed products. Since sodium carbonate solution is alkaline you can legitimately write a hydroxide ppt. equation as for sodium hydroxide above but it doesn't show the formation of carbon dioxide.

      • You can write an equation to show the formation of carbon dioxide leaving a soluble cationic complex of aluminium in solution and this equation fits in well with the acid-base nature of this reaction.

        • [Al(H₂O)₆]³⁺_(aq) + CO₃^2-_(aq) ==> 2[Al(H₂O)₄(OH)₂]⁺_(aq) + CO_2(g) + 3H₂O_(aq)

        • This equation shows the hexaaquaaluminium ion acting as a Bronsted-Lowry acid donating two protons to the carbonate ion (B-L base) to form carbon dioxide and water.

    • This reaction shows why 'aluminium carbonate' 'Al₂(CO₃)₃' cannot exist. The hydrated highly charged central metal ion is too acidic to co-exist with a carbonate ion. The same situation applies to the chromium(III) Cr³⁺ and iron(III) Fe³⁺ ions i.e. no chromium(III) carbonate or iron(III) carbonate exists. However with a lesser charged, lesser acidic ion, carbonates can exist, so there is an iron(II) carbonate FeCO₃.

• The addition of excess sodium carbonate solution has no further effect. Sodium carbonate is too weak a base to effect the amphoteric nature of aluminium hydroxide.

Quick click to Introduction * Sc * Ti * V * Cr * Mn * Fe * Co * Ni * Cu * Zn * Ag/Pt etc.

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