1 / 59

Ionic bonding evidence for ionic bonding, electron density maps trends in radii Born Haber cycles

Ionic bonding evidence for ionic bonding, electron density maps trends in radii Born Haber cycles explaining formulae, why AlO is incorrect polarisation Metallic Bonding. Covalency electron density maps giant atomic structures dot-cross diagrams shapes of molecules VSEPR

edda
Download Presentation

Ionic bonding evidence for ionic bonding, electron density maps trends in radii Born Haber cycles

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Ionic bonding evidence for ionic bonding, electron density maps trends in radii Born Haber cycles explaining formulae, why AlO is incorrect polarisation Metallic Bonding Covalency electron density maps giant atomic structures dot-cross diagrams shapes of molecules VSEPR electronegativity polarity of covalent bonds polarity of molecules • Intermolecular forces • trends in physical properties • Solutions and dissolving • why certain substances dissolve in particular solvents

  2. Melting Points of Period 3 elements

  3. Melting Points of Period 2 elements

  4. Ionic bonding • The electrostatic attraction between oppositely charged ions • Metals, hydrogen and ammonium form positive ions (cations). • Non-metals form negative ions (anions).

  5. Evidence for ionic compounds • High melting points • strong electrostatic attractions between oppositely charged ions • Electrical conductivity only in liquid state or aqueous solution because ions need to move. • Coloured ions can be observed migrating to electrodes during electrolysis (e.g. CuCr2O7) • green / blue Cu2+ (aq) moves to cathode • yellow Cr2O72- (aq) moves to anode • Electron density maps show low electron density between the oppositely charged ions.

  6. DHat = + 107.3 kJ mol-1 DHformation = - 411.2 kJ mol-1 DHlatt (=348.8-121.7-496-107.3-411.2) = - 787.4 kJ mol-1 Na+ (g) + e- + Cl (g) DHat [Cl]= + 121.7 kJ mol-1 Eaff[Cl] = - 348.8 kJ mol-1 Na+ (g) + e- + ½ Cl2 (g) Na+ (g) + Cl- (g) Em1[Na] = + 496 kJ mol-1 Enthalpy (H) Born Haber cycle for sodium chloride Na (g) + ½ Cl2 (g) Na (s) + ½ Cl2 (g) NaCl (s)

  7. DHat = + 89.2 kJ mol-1 DHformation = - 436.7 kJ mol-1 DHlatt (=348.8-121.7-419-89.2-436.7) = - 717.8 kJ mol-1 K+ (g) + e- + Cl (g) DHat [Cl]= + 121.7 kJ mol-1 Eaff[Cl] = - 348.8 kJ mol-1 K+ (g) + e- + ½ Cl2 (g) K+ (g) + Cl- (g) Em1[K] = + 419 kJ mol-1 Enthalpy (H) Born Haber cycle for potassium chloride K (g) + ½ Cl2 (g) K (s) + ½ Cl2 (g) KCl (s)

  8. DHat [K]= + 89.2 kJ mol-1 DHformation = - 393.8 kJ mol-1 DHlatt =(324.6-111.9-419-89.2-393.8) = - 689.3 kJ mol-1 K+ (g) + e- + Br (g) DHat [Br]= + 111.9 kJ mol-1 Eaff[Br] = - 324.6 kJ mol-1 K+ (g) + e- + ½ Br2 (l) K+ (g) + Br - (g) Em1[K] = + 419 kJ mol-1 Enthalpy (H) Born Haber cycle for potassium bromide K (g) + ½ Br2 (l) K (s) + ½ Br2 (l) KBr (s)

  9. 2 Li+ (g) + O2- (g) 2 DHat = + 318.8 kJ mol-1 DHformation = - 597.9 kJ mol-1 DHlatt (=-798+141.1-249.2-1040-318.8-597.6) = - 2862.8 kJ mol-1 2 Li+ (g) + 2 e- + O (g) Eaff[O-] +798 kJ mol-1 Eaff[O]= - 141.1 kJ mol-1 2 Li+ (g) + e- + O- (g) DHat [O]= + 249.2 kJ mol-1 2 Li+ (g) + 2 e- + ½O2 (g) Enthalpy (H) 2 Em1[Li] = + 1040 kJ mol-1 2 Li (g) + ½O2 (g) Born Haber cycle for lithium oxide 2 Li (s) + ½ O2 (g) Li2O (s)

  10. DHformation=+148+738+1451+243.4-679.6-2526 - 625.2 kJ mol-1 DHlatt = = - 2526 kJ mol-1 Mg2+ (g) + 2e- + 2 Cl (g) 2 DHat [Cl]= + 243.4 kJ mol-1 2 Eaff [Cl] = - 697.6 kJ mol-1 Mg2+ (g) + 2e- + Cl2 (g) Mg2+ (g) + 2 Cl- (g) Enthalpy (H) Em2[Mg] = + 1451 kJ mol-1 Mg+ (g) + e- + Cl2 (g) Em1[Mg] = + 738 kJ mol-1 Mg (g) + Cl2 (g) DHat(Mg) = + 148 kJ mol-1 Mg (s) + Cl2 (g) Born Haber cycle for magnesium chloride MgCl2 (s)

  11. DHat = + 148 kJ mol-1 DHformation[MgCl (s)] (=+148+738+121.7-248.8-780) = - 21.1 kJ mol-1 DHlatt[MgCl (s)] = - 780 kJ mol-1 Mg+ (g) + e- + Cl (g) DHat [Cl]= + 121.7 kJ mol-1 Eaff[Cl] = - 348.8 kJ mol-1 Mg+ (g) + e- + ½ Cl2 (g) Mg+ (g) + Cl- (g) Em1[Mg] = + 738 kJ mol-1 Enthalpy (H) Born Haber cycle for MgCl Mg (g) + ½ Cl2 (g) Mg (s) + ½ Cl2 (g) MgCl (s)

  12. DHat = + 148 kJ mol-1 DHformation= +148+9927+366-1047-4500 = + 4894 kJ mol-1 DHlatt = - 4500 kJ mol-1 Mg3+ (g) + 3e- + 3 Cl (g) 3 DHat [1/2Cl2]= + 366 kJ mol-1 3 Eaff[Cl] = - 1047 kJ mol-1 Mg3+ (g) + 3e- + 3/2 Cl2 (g) Mg3+ (g) + 3 Cl- (g) Em1+ Em2+Em3 [Mg] = (738+1451+7738) kJ mol-1 = + 9927 kJ mol-1 Enthalpy (H) Born Haber cycle for MgCl3 Mg (g) + 3/2 Cl2 (g) MgCl3 (s) Mg (s) + 3/2 Cl2 (g)

  13. DHat [Ca] = + 178.2 kJ mol-1 DHformation = - 533.5 kJ mol-1 DHlatt =(590.8-214-1145-590-178.2-533.5) = - 2069.9 kJ mol-1 Ca2+ (g) + 2e- + 2 I (g) 2 DHat [I]= + 214 kJ mol-1 2 Eaff[I]= - 590.8 kJ mol-1 Ca2+ (g) + 2e- + I2 (g) Ca2+ (g) + 2 I- (g) Enthalpy (H) Em2[Ca] = + 1145 kJ mol-1 Ca+ (g) + e- + I2 (g) Em1[Ca] = + 590 kJ mol-1 Born Haber cycle for Calcium iodide Ca (g) + I2 (g) Ca (s) + I2 (g) CaI2 (s)

  14. Mg2+ (g) + O2- (g) DHlatt = (-798+141.1-249.2-1451-738-147.7-601.7) = - 3844.5 kJ mol-1 DHformation = - 601.7 kJ mol-1 Mg2+ (g) + 2 e- + O (g) Eaff[O-] +798 kJ mol-1 Eaff[O]= - 141.1 kJ mol-1 DHat [O]= + 249.2 kJ mol-1 Mg2+ (g) + e- + O- (g) Mg2+ (g) + 2 e- + ½O2 (g) Enthalpy (H) Em2[Mg] = + 1451 kJ mol-1 Mg+ (g) + e- + ½O2 (g) Em1[Mg] = + 738 kJ mol-1 Mg (g) + ½O2 (g) DHat(Mg) = + 147.7 kJ mol-1 Born Haber cycle for magnesium oxide Mg (s) + ½ O2 (g) MgO (s)

  15. DHat = + 326.4 kJ mol-1 DHformation = - 1504.1 kJ mol-1 DHlatt (984-237-5140-326.4-1504.1) = - 6223.5 kJ mol-1 B3+ (g) + 3e- + 3 F (g) 3 DHat [F]= + 237 kJ mol-1 3 Eaff[F] = - 984 kJ mol-1 B3+ (g) + 3e- + 3/2 F2 (g) B3+ (g) + 3 F- (g) Em1+ Em2+Em3 [B] = (578+1817+2745) kJ mol-1 = + 5140 kJ mol-1 Enthalpy (H) Born Haber cycle for Boron fluoride B (g) + 3/2 F2 (g) B (s) + 3/2 F2 (g) BF3 (s)

  16. 2 Al3+ (g) + 3 O2- (g) DHformation = - 1675.7 kJ mol-1 DHlatt = (-2394+423.3-747.6-10280-652.8-1675.7) = - 15 327 kJ mol-1 2 Al3+ (g) + 6 e- + 3 O (g) 3 Eaff[O-] +2394 kJ mol-1 3 Eaff[O]= - 423.3 kJ mol-1 3 DHat [O] = + 747.6 kJ mol-1 2 Al3+(g)+3e-+3O- (g) 2 Al3+ (g) + 3 e- + 3/2O2 (g) Enthalpy (H) 2(Em1+ Em2+ Em3 )[Al] = 2(578+1817+2745) kJ mol-1 = + 10 280 kJ mol-1 2 Al (g) + 3/2O2 (g) 2 DHat(Al) = + 652.8 kJ mol-1 Born Haber cycle for aluminium oxide 2 Al (s) + 3/2O2 (g) Al2O3 (s)

  17. 2 B3+ (g) + 3 O2- (g) DHformation = - 1273 kJ mol-1 DHlatt = (-2394+423.3-747.6-13776-1025.4-1273) = - 18 800 kJ mol-1 2 B3+ (g) + 6 e- + 3 O (g) 3 Eaff[O-] +2394 kJ mol-1 3 Eaff[O]= - 423.3 kJ mol-1 3 DHat [½ O2 (g)] = + 747.6 kJ mol-1 2 B3+(g)+3e-+3O- (g) 2 B3+ (g) + 3 e- + 3/2O2 (g) Enthalpy (H) 2(Em1+ Em2+ Em3 )[B] = 2(801+2427+3660) kJ mol-1 = + 13 776 kJ mol-1 2 B (g) + 3/2O2 (g) 2 DHat(B) = + 1025.4 kJ mol-1 Born Haber cycle for boron oxide 2 B (s) + 3/2O2 (g) B2O3 (s)

  18. Lattice energies

  19. Not Al2+ O2- Whilst successive ionisation energies of Al increase, Em3is not especially large. Al3+ is very much smaller than Al2+, since its 3rd principal quantum shell is now empty. Consequently, the ions pack more tightly in (Al3+)2(O2-)3. Al3+ also carries a greater charge than Al2+,,increasing the attraction to O2- anions. The lattice energy of (Al3+)2(O2-)3 is therefore much greater in magnitude than that of Al2+O2-. AlO is not the formula of aluminium oxide

  20. Not Al3+ O3- O3- would have a greater radius than O2- since its extra electron occupies a new principal quantum shell, further from the nucleus and more shielded by inner quantum shells. In spite of the increased charge of the O3- anion, its large size reduces packing density of the solid. The electron affinity required to form O3- from O2- would be large and endothermic. Therefore the lattice energy of Al3+O3- does not make up for the endothermic steps in the Born Haber cycle. AlO is not the formula of aluminium oxide

  21. Y- X+ Y- X+ Polarisation of the anion

  22. Factors leading to anion polarisation • Cation polarizing power increases with • small radius (increasing charge density) • large positive charge (increasing charge density) • Anion polarizability increases with • large radius (outer electrons far from nucleus and shielded by inner shells) • increasing negative charge increases its size • Increasing anion polarisation means increasing covalent character to the bonding • indicated by large difference between theoretical and experimental lattice energies

  23. Metallic bonding • The attraction between 'positive ions' and a sea of delocalised electrons. • Why does the melting point increase across a period, Na<Mg<Al? • Electrical and thermal conductivity due to transfer of charge and energy by the movement of delocalised electrons.

  24. The covalent bond • The attraction between two nuclei and a shared pair of electrons. • One electron of the shared pair originating from each atom in a 'standard' covalent bond. • Both electrons of the shared pair originate from the same atom in a dative bond. • 'Standard' and dative covalent bonds are indistinguishable.

  25. Hydrogen molecule (H2) H H H-H

  26. Valence Shell Electron Pair Repulsion • Sigma bond electron pairs and lone pairs all repel each other around the central atom. • The electron pairs move into positions of maximum separation. • 2 pairs gives 180: 3 pairs gives 120: 4 pairs gives 109.5, 6 pairs gives 90. • Lone pairs have a greater repulsion than sigma bond pairs. • Each lone pair reduces the expected bond pair - bond pair angle by about 2.5.

  27. Chlorine molecule, Cl2 Cl-Cl Cl Cl

  28. Cl H Hydrogen chloride molecule, HCl (g) {not HCl (aq) which is ionic} H-Cl Why is this covalent?

  29. O H H 109.5 H O H Water molecule, H2O

  30. N N N N Nitrogen molecule, N2

  31. N H H H 107 H N H H Ammonia, NH3

  32. H H H C H H H C H H Methane, CH4 H H H C 109.5 H

  33. Ethane, C2H6 H C H H H3C H H C 109.5 C H H H H H H H C C H H H

  34. Ethene, C2H4 H H C 121 118 C H H H H C = C H H

  35. O C O Carbon dioxide, CO2 O = C = O

  36. H S H H2S 104

  37. H H Si H H SiH4 109.5

  38. H C H O Methanal, HCHO 120 120

  39. HCN H C N 180

  40. H H 104 O O 104

  41. + H + H N H 109.5 H

  42. - H O

  43. 2- O 2- O S O 109.5 O

  44. 2- O 2- S S O 109 O

  45. H H C C H H H H 121 118 C = C H H Ethene • 3 s-bond pairs around each C atom repel to positions of maximum separation. • trigonal planar

  46. H H H H H H H C C C C H H H H H H H H C C C C H Ethene or s bonds shown as lines and wedges

  47. Benzene, C6H6 neither nor but

  48. Benzene • p bonding is delocalised over the whole ring because all 6 p orbitals are coplanar and overlap. • not 3 separate p bonds • Benzene is more stable than alkenes and tends to react by substitution rather than addition.

  49. Pyrene, C16H10

  50. Graphite Flat sheet of C atoms Weak forces between sheets

More Related