Periodicity and transition metals flashcards

Periodicity describes the repeating trends in physical and chemical properties of elements across periods of the periodic table. Students must understand how atomic radius, ionisation energy, electronegativity and melting point vary across Period 2 and Period 3, and be able to explain these trends using nuclear charge, shielding and bonding. Transition metals occupy the d-block and are defined by having an incomplete d sub-shell in at least one stable ion. Key properties include variable oxidation states, the ability to form coloured ions, catalytic behaviour and complex ion formation. Students should be able to explain why transition metals show these properties in terms of their electronic configurations, and understand ligand substitution and the role of ligands in determining colour and shape of complex ions.

1. Card 1

   Why does atomic radius decrease across Period 3 from Na to Cl?

   Nuclear charge increases across the period while shielding stays roughly constant, so electrons are pulled closer to the nucleus, reducing atomic radius.

2. Card 2

   What is the general trend in first ionisation energy across Period 3?

   First ionisation energy generally increases across Period 3 due to greater nuclear charge attracting outer electrons more strongly, with small dips at Mg-Al and P-S.

3. Card 3

   Why is the first ionisation energy of Al slightly lower than that of Mg?

   Al's outer electron is in a 3p sub-shell, which is slightly higher in energy and better shielded than Mg's 3s electron, making it easier to remove.

4. Card 4

   Why is the first ionisation energy of S slightly lower than that of P?

   In S, two electrons share a 3p orbital, causing repulsion that makes one easier to remove. P has one electron in each 3p orbital with no such repulsion.

5. Card 5

   How does the melting point of Period 3 elements vary and why does it peak at Si?

   Melting point rises from Na to Si due to stronger metallic bonding and then giant covalent structure at Si, then drops sharply for the simple molecular elements P, S, Cl, Ar.

6. Card 6

   What defines a transition metal at A-Level?

   A transition metal is a d-block element that forms at least one stable ion with an incomplete d sub-shell. Sc and Zn are excluded by this definition.

7. Card 7

   Why can transition metals display variable oxidation states?

   The 3d and 4s sub-shells are very close in energy, so different numbers of d electrons can be lost without requiring a large jump in energy, allowing multiple stable oxidation states.

8. Card 8

   Why are transition metal ions coloured in solution?

   Ligands split the d orbitals into two energy levels; electrons absorb visible light to jump between them, and the complementary colour is transmitted or reflected.

9. Card 9

   What is a ligand?

   A ligand is a molecule or ion that donates a lone pair of electrons to a central metal ion to form a coordinate (dative covalent) bond in a complex ion.

10. Card 10

    What is the coordination number of a complex ion?

    The coordination number is the total number of coordinate bonds formed between the central metal ion and its surrounding ligands, commonly 4 or 6.

11. Card 11

    Why do transition metals act as heterogeneous catalysts?

    Reactants adsorb onto the metal surface, weakening bonds and lowering activation energy. Variable oxidation states also allow homogeneous catalysis by electron transfer.

12. Card 12

    What happens to the colour of a complex ion when one ligand is substituted for another?

    Different ligands cause different degrees of d orbital splitting, changing the energy of light absorbed and therefore altering the observed colour of the complex.

• Many students incorrectly state that all d-block elements are transition metals, forgetting that scandium and zinc are excluded because their stable ions do not have an incomplete d sub-shell.
• Students often confuse the colour of a complex ion with the colour of light it absorbs, when in fact the observed colour is the complementary colour to the one absorbed.
• A common error is to say that ionisation energy increases uniformly across a period, ignoring the dips at aluminium and sulphur which arise from sub-shell and electron pairing effects.

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