Additional Aspects of
Molecular Bonding & Structure
Chapters 8 and 9
BLB 12th
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 ↓`
Hess’s Law for lattice energy (Born-Haber cycle)
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: s electron(s) are lost
first.
8.3 Covalent Bonding
• Atoms share electrons.
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
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
The greater the difference in electronegativity, the more
polar the bond.
Electronegativity
Electronegativity difference • Examples:
• < 0.5 nonpolar
C–H
• 0.5-2.0 polar
• > 2.0 ionic
N–O
Rough guidelines only.
See p. 304.
Na–Cl
Cl–Cl
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.
CO2 a nonpolar molecule
H2O a polar molecule
9.3 Molecular Polarity
• Examples
KrF2
SO2
ICl3
XeO4
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)
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.
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!
Resonance Structures
• The measured bond lengths are an average of the
representative structures; somewhere between a
single and double bond length.
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.
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
413 kJ + 242 kJ
−104 kJ
328 kJ + 431 kJ
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
Orbital Overlap
σ bonds – single area of overlap
Orbital Overlap in the formation of H2
sp Hybrid Orbitals
sp2 Hybrid Orbitals
sp3 Hybrid Orbitals
Hybrid Orbital Overlap
9.6 Multiple Bonds
C2H4 – the whole picture
C2H2 – the whole picture