Advanced Molecular Bonding: Ionic & Covalent Energetics, Electronegativity, Polarity, Orbitals

1. Additional Aspects of Molecular Bonding & Structure Chapters 8 and 9 BLB 12th

2. 8.2 Ionic Bonding Energetics of Ionic Bond Formation Na(s) + ½ Cl2(g) → NaCl(s) ΔHf° = −410.9 kJ ΔH = 147 kJ/mol • Lattice energy – energy required to completely separate the ions in one mole of an ionic compound Na+(g) + Cl‾(g) → NaCl(s) ΔHlattice = −788 kJ • Lattice energy ↑ as ion charges ↑ and size ↓`

4. Hess’s Law for lattice energy (Born-Haber cycle)

5. 8.2 Ionic Bonding • Electron Configurations of Ions (also 7.4) • Atoms will gain or lose electrons to achieve a noble gas configuration. • Transition metal ions: selectron(s) are lost first.

6. 8.3 Covalent Bonding • Atoms share electrons.

7. 8.4 Bond Polarity • Unequal sharing of electrons in a covalent bond • Electronegativity – the ability of an atom in a molecule (bonded) to attract electrons to itself

9. 8.4 Bond Polarity • Nonpolar covalent bond – electrons shared equally • Polar covalent bond – electrons shared unequally due to different electronegativity values • Greater electronegativity difference, more polar the bond (higher dipole moment) • Dipole moment – measured magnitude of a dipole. • Predict direction • No calculations

10. The greater the difference in electronegativity, the more polar the bond.

11. Electronegativity difference < 0.5 nonpolar 0.5-2.0 polar > 2.0 ionic Rough guidelines only. See p. 304. Examples: C–H N–O Na–Cl Cl–Cl Electronegativity

12. 9.3 Molecular Polarity • Depends upon the polarities of the bonds and the molecular geometry of the molecule • Bond dipole moments are vector quantities. • A molecular dipole moment is the vector sum of its bond dipoles. • A molecule can be nonpolar, that is, have a net dipole moment of zero, even if bond dipole(s) exist.

13. CO2 a nonpolar molecule H2O a polar molecule

15. 9.3 Molecular Polarity • Examples KrF2 SO2 ICl3 XeO4

16. 8.5 Drawing Lewis Structures • Formal Charges (p. 307) • There may be more than one valid Lewis Structure for a given molecule. • Formal charges are used to determine the most reasonable structure. • Calculate a formal charge (FC) for each atom: • FC = (# valence e¯) − (# e¯ belonging to atom)

17. 8.5 Drawing Lewis Structures • Formal Charges • Best structure? The one with the formal charges closest to zero and where the most negative charges reside on the most electronegative atoms. • For H, Be, and B, formal charges indicate that the Lewis structure with an incomplete octet is more important than the ones with double bonds.

19. 8.6 Resonance Structures • Lewis structures and the VSEPR theory are means by which we try to mimic or predict the experimentally determined properties of molecules. • When a combination of single and multiple bonds are used, it implies that the bond lengths are unequal. • They’re not!

20. Resonance Structures • The measured bond lengths are an average of the representative structures; somewhere between a single and double bond length.

22. Resonance • The organic compound benzene, C6H6, has two resonance structures. • It is commonly depicted as a hexagon with a circle inside to signify the delocalized electrons in the ring.

23. 8.8 Strengths of Covalent Bonds • Bond enthalpy – energy required to break a bond • Enthalpy of reaction – estimate from difference of the bonds broken minus the bonds formed

24. 413 kJ + 242 kJ 328 kJ + 431 kJ −104 kJ

26. 9.4 Orbital Overlap & 9.5 Hybrid Orbitals • Electrons exist in orbitals. • Valence Bond Theory – bonding model where orbitals overlap to form bonds • Hybridization combines orbitals into hybrid orbital sets that match experimentally determined geometried. • σ bond – one area of overlap • π bond – two areas of overlap

27. Orbital Overlap σ bonds – single area of overlap

28. Orbital Overlap in the formation of H2

29. sp Hybrid Orbitals

30. sp2 Hybrid Orbitals

31. sp3 Hybrid Orbitals

32. Hybrid Orbital Overlap

33. 9.6 Multiple Bonds

34. C2H4 – the whole picture

35. C2H2 – the whole picture