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Structure 2.3 Flashcards

Active Recall for The Metallic Model

Flashcards: read all 15 questions and answers
  1. Define metallic bonding using the exact IB wording.
    The electrostatic attraction between a lattice of positive metal cations and a sea of delocalised electrons. You must say CATIONS (not atoms) and DELOCALISED (not free).
  2. Why are metals good electrical conductors?
    Delocalised electrons are free to move through the structure when a potential difference is applied, carrying charge. Cation vibrations also transfer kinetic energy (thermal).
  3. Why are metals malleable and ductile?
    Layers of cations can slide over each other without breaking bonds because the electron sea is non-directional and redistributes around the new positions.
  4. EXAM TRAP: What is the key vocabulary mistake in metallic bonding?
    Saying free electrons instead of DELOCALISED electrons, or saying atoms instead of CATIONS. Also never say metals have more metallic bonds - say stronger electrostatic attraction.
  5. What two factors increase metallic bond strength?
    (1) Higher cation charge = more delocalised electrons = stronger attraction. (2) Smaller ionic radius = electrons closer to nucleus = stronger attraction. E.g. Mg > Na.
  6. Compare Na and Mg melting points and explain why.
    Mg (650 C) >> Na (98 C). Mg2+ has higher charge (+2 vs +1), more delocalised electrons (2 vs 1), and smaller ionic radius (72 pm vs 102 pm) giving stronger metallic bonding.
  7. What is an alloy?
    A MIXTURE of a metal with one or more other elements. NOT a compound. The different-sized atoms disrupt the regular lattice, preventing layers from sliding, making it harder.
  8. What is a substitutional alloy?
    An alloy where the added atoms are SIMILAR in size to the host metal and REPLACE some host atoms in the lattice positions. Examples: bronze (Cu + Sn), brass (Cu + Zn).
  9. What is an interstitial alloy?
    An alloy where the added atoms are MUCH SMALLER and fit into the GAPS (interstices) between the host metal atoms. Example: steel (Fe + C).
  10. EXAM TRAP: Why are alloys harder than pure metals?
    Different-sized atoms prevent the LAYERS OF CATIONS FROM SLIDING over each other. Simply saying it disrupts the lattice is NOT enough for full marks.
  11. How does alloying affect electrical conductivity?
    Alloying typically DECREASES conductivity because the irregular lattice scatters the flow of delocalised electrons.
  12. Why do metals have lustre (shine)?
    Delocalised electrons absorb photons of visible light and re-emit them at many frequencies, producing a shiny reflective surface.
  13. (HL) Define transition element.
    By common exam convention: a d-block element that forms at least one stable ion with an incomplete (partially filled) d-sublevel. Note: IUPAC also includes neutral atoms with incomplete d-subshells (qualifying Sc as [Ar] 3d1 4s2, while Zn is excluded under both conventions).
  14. (HL) Why do transition metals have higher melting points than s-block metals?
    They delocalise both s AND d electrons into the electron sea. More delocalised electrons + higher cation charge + smaller radius = much stronger metallic bonding. E.g. Fe (1538 C) vs Na (98 C).
  15. (HL) Why are Zn and Sc often excluded from transition metals despite being d-block?
    Under the stable-ion convention, Zn forms only Zn2+ (3d10, completely full) and Sc forms only Sc3+ (3d0, completely empty). However, under the IUPAC neutral-atom definition, Sc ([Ar] 3d1 4s2) has an incomplete d-subshell and qualifies, whereas Zn is excluded under both conventions.
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