Additional Aspects of Molecular Bonding & Structure Chapters 8 and 9 BLB 12 th.
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Transcript of Additional Aspects of Molecular Bonding & Structure Chapters 8 and 9 BLB 12 th.
8.2 Ionic BondingEnergetics 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 ↓`
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.
8.4 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
Electronegativity
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
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.
8.5 Drawing Lewis StructuresFormal 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 StructuresFormal 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
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