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Metallic Bonding

10. Metallic Bonding. 10.1 Metallic Bonding 10.2 Metallic Radius 10.3 Factors Affecting the Strength of Metallic Bond 10.4 Metallic Crystals 10.5 Alloys. Nature of Metallic Bonding. Li(g) + Li(g) Li + Li  (g) (1). Li(g) + Cl(g) Li + Cl  (g) (2).

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Metallic Bonding

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  1. 10 Metallic Bonding 10.1 Metallic Bonding 10.2 Metallic Radius 10.3 Factors Affecting the Strength of Metallic Bond 10.4 Metallic Crystals 10.5 Alloys

  2. Nature of Metallic Bonding

  3. Li(g) + Li(g) Li+ Li(g) (1) Li(g) + Cl(g) Li+ Cl(g) (2) Not ionic : - ∵ atoms of the same electronegativity No favourable favourable

  4. Li(g) + Li(g) Li+ Li(g) (1) Li(g) + Cl(g) Li+ Cl(g) (2) Q.57 Explain, with the aid of suitable enthalpy change cycles, why reaction (1) is energetically less favourable than reaction (2). (1st EA : Li = 60 kJ mol1, Cl = 349 kJ mol1)

  5. H2 H4 Li+(g) + Li(g) Li+(g) + Cl(g) H1 Li(g) + Li(g) Li+ Li(g) 1st E.A. of Li 1st I.E. of Li H1 = 1st I.E. of Li + (60 kJ mol1) + H2 H3 Li(g) + Cl(g) Li+ Cl(g) 1st E.A. of Cl 1st I.E. of Li H3 = 1st I.E. of Li + (349 kJ mol1) + H4

  6.  H4 is more negative than H2  Cl is smaller than Li

  7. H1 = 1st I.E. of Li + (60 kJ mol1) + H2 H3 = 1st I.E. of Li + (349 kJ mol1) + H4 Li(g) + Li(g) Li+ Li(g) (1) Li(g) + Cl(g) Li+ Cl(g) (2) H1 No favourable H3 favourable H4 is more negative than H2 H3 is more negative than H1

  8. Not covalent : - ∵ efficiency of orbital overlap  as the bonding atoms get larger

  9. Explanation 1. Classical approach Distance between shared pair and the bonding nuclei : H2 < Li2 < Na2 < K2 < Rb2 < Cs2 Bond strength : H–H > Li–Li > Na–Na > K–K > Rb–Rb > Cs–Cs

  10. LiA Li2 LiB 1s*   2sA 2sB  1s 2. MO approach (Not required in AL)

  11. Be(A) Be2 Be(B) 1s* 2s* 2sA 2sB 1s 2s Q.58 There is no gain of stability when the AOs of two Beryllium atoms overlap. 2 e involved in bonding 2 e involved in antibonding Overall : - No e involved in bonding Bond order = 0

  12. Conclusion : - 1. Metals tend to form giant structures rather than discrete molecules. 2Li  Li – Li (low bond enthalpy) nLi  Lin (high bond enthalpy) Stronger bonds are formed due to extensive delocalization of valence electrons.

  13. Conclusion : - 2. The electron-sea model The valence electrons do not belong to any specific atoms (not localized) but delocalize throughout the whole crystal structure.

  14. Conclusion : - 2. The electron-sea model Mobile es  electron sea Stationary +ve ions

  15. Conclusion : - 2. The electron-sea model The electrostatic attractive forces between the delocalized electron cloud and the positive ions are called the metallic bonds

  16. Since metallic bonds are non-directional, they exist in significant extent even in molten state. The boiling points of metals are much higher than the corresponding melting points. E.g. Na m.p.  97.8oC ; b.p.  903.8oC NaCl m.p. = 801C ; b.p. = 1413C

  17. Conclusion : - 3. MO approach : Band theory Spacing  as the no. of molecular orbitals  Li Li2 Li3 Li4 Lin

  18. Lin : n orbitals overlap  continuous band Half-filled band

  19. Metallic Radius

  20. 10.2 Metallic radius (SB p.262) Metallic radius (r) is defined as half of the internuclear distance between adjacent atoms in a metal crystal.

  21. 10.2 Metallic radius (SB p.262) Trend of metallic radius in the Periodic Table • Moving down a group, metallic radii increase • Going across a period, metallic radii decrease

  22. Factors Affecting the Strength of Metallic Bond

  23. The strength of metallic bond can be estimated by melting point, boiling point, enthalpy change of fusion or enthalpy change of vapourization. Higher m.p./b.p./Hfusion/Hvap  stronger metallic bond

  24. 10.3 Factors affecting the strength of metallic bond (SB p.262) The metallic bond strength increases with: 1. decreasing size of the metal atom (i.e. the metallic radius); 2. increasing number of valence electrons of the metal atom.

  25. Let's Think 1 10.3 Factors affecting the strength of metallic bond (SB p.263) Effect of number of valence electrons on metallic bond strength

  26. 10.3 Factors affecting the strength of metallic bond (SB p.263) Effect of metallic radius on metallic bond strength of Group IA metals

  27. Metal Ni Cu Ag Pb Hg Au Density (g cm3) 8.91 8.94 10.49 10.66 13.53 19.30 Typical properties of metals 1. High density due to close packing of atoms in metallic crystal (h.c.p./f.c.c. co-ordination number  12)

  28. Typical properties of metals 1. High density Exception : Alkali metals have low densities (< 1 for Li, Na and K ) (a) they have more open structures (b.c.c. /co-ordination number  8) (b) their atomic radii are the highest in their own Periods. E.g. Size : Na > Mg > Al

  29. Typical properties of metals 2. High melting point and boiling point Extensive delocalization of valence electrons  stronger bonds Bond strength : - ionic bond  covalent bond  metallic bond

  30. Typical properties of metals 3. High flexibility Malleability : The ability to be deformed under compression Ductility : The ability to be deformed under tension

  31. Typical properties of metals 3. High flexibility Reasons : - (a) The presence of layers in the crystal lattice i.e. the layers can slide over one another under strain (b) Metallic bonds are non-directional. i.e. electrons can take up new positions and reform metallic bond after the deformation

  32. Typical properties of metals

  33. Lin Half-filled 2s band Since the gap between energy levels are extremely small, radiation of any frequency in visible region can be absorbed and emitted. Typical properties of metals  Silvery and shiny 4. Surface lustre

  34. Half-filled s band E = 220 kJ mol1 full-filled d band Typical properties of metals Cu : 3d10, 4s1 Reddish brown

  35. Half-filled s band E = 300 kJ mol1 full-filled d band Typical properties of metals Au : 5d10, 6s1 Golden yellow

  36. UV light absorbed Half-filled s band E = 380 kJ mol1 full-filled d band Typical properties of metals Ag : 4d10, 5s1 Silvery

  37. Typical properties of metals 5. High Thermal and Electrical Conductivity Due to the free movement of delocalized electrons

  38. Metallic Crystals

  39. 10.4 Metallic crystals (SB p.263) Closed-packed structure • Closed-packed structure is made possible with identical particles • Two types : - • 1. Hexagonal closed-packed, h.c.p. • Cubic closed-packed, c.c.p. Or • Face-centred cubic, f.c.c.

  40. 10.4 Metallic crystals (SB p.263) Hexagonal close-packed structure abab…

  41. 10.4 Metallic crystals (SB p.265) Hexagonal close-packed structure (a) normal side view (b) exploded view (c) a unit cell Packing efficiency = 74 % Co-ordination no. = ?

  42. Co-ordination number = 12

  43. rotate by 45 10.4 Metallic crystals (SB p.265) Cubic close-packed / Face-centred cubic structure c.c.p. or f.c.c. abcabc… Packing efficiency = 74 % Co-ordination no. = 12

  44. 10.4 Metallic crystals (SB p.266) Open structure • Structures with more empty space between the atoms • Most common: body-centred cubic structure

  45. 10.4 Metallic crystals (SB p.267) Body-centred cubic structure (a) normal side view (b) exploded view (c) a unit cell Packing efficiency = 68 % Co-ordination no. = 8

  46. Given : Density of Cu = 8.94 g cm3 Relative atomic mass of Cu = 63.546 Atomic radius of Cu = 0.128 nm Cu adopts f.c.c. structure Calculate the Avogadro’s constant a2 + a2 = (4r)2

  47. Since one unit cell contains 4 Cu atoms = 5.99  1023 mol1

  48. Alloys

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