Bonding II
Lewis Structures
and
Covalent Bond Properties
Lewis Structures
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Introduced by G.N. Lewis (U. C. Berkeley) in 1916
Covalent bonding attributed to shared electron pairs
Electron pairs are localized on or between atoms
Shared pairs of electrons are “owned” by both atoms
C,N,O and F always obey the “Octet Rule”
Elements lighter than C have less than 8 eElements heavier than F may have more than 8 eReview method for writing Lewis structures:
– “simple” examples:CH4, H2O, CO2, etc.
Resonance
• Since Lewis model assumes that electrons are localized,
a single Lewis structure cannot adequately represent a
molecule with delocalized bonding.
• Resonance forms are drawn by assigning bonding and
lone pairs in different ways; the geometric arrangement
of atoms does NOT change! (e.g.: SO2,NO3-, etc.)
• Molecular electronic structure is described as a “hybrid”
(mixture) of the individual resonance forms.
• Equivalent structures contribute equally (same energy).
• Relative contributions of different structures are
determined by assessing their relative energies.
Formal Charge
• Relative merit of different resonance forms may be
assessed by assigning formal charge to each atom.
• Formal charge counting rule: divide bonding pairs of
electrons between atoms (i.e., 100% covalent bonding is
assumed)
• Formal charge does NOT reflect the actual charges
borne by atoms in a structure!
• Favored (lowest energy) structure:
– Minimizes formal charge separation within the molecule.
– Places negative charge (if any) on most electronegative atom.
– Places positive charge (if any) on least electronegative atom.
• Example: cyanate (OCN-) vs. fulminate (CNO-)
Formal Oxidation Number
• The formal oxidation number represents the number of
electrons gained or lost by an element when it forms a
compound.
• Oxidation number counting rule: assign both electrons of
each bonding pair to the more electronegative atom
(100% ionic) (e.g. NH3, NCl3, NO3-, etc.)
• If two atoms of same element are bonded together,
divide the bonding electrons (e.g. C2Cl4, N2O, etc.)
• Formal oxidation number is a “book keeping tool”; it
DOES NOT represent the “true” charge on any atom!
Hypervalence
• Atoms that necessarily exceed 8 e- in their valence shell
to accommodate bonding of peripheral atoms (or lone
pairs) are said to be “hypervalent”, or to have “expanded
octets”
– Examples: PCl5, PF6-, SF6, BrF3, etc.
– Contrary to popular belief, hypervalence does not necessarily
require involvement of d-orbitals in bonding!
– Occurrence of hypervalence may have more to do with the
relative ease of fitting more atoms around a larger central atom.
• It may be possible to draw resonance forms with > 8 ein the valence shell for some molecules (e.g. SO42-, PO43-,
(CH3)2SO, etc.) but these are NOT properly considered to
be “hypervalent” molecules.
Molecular Geometry
• Molecular geometry may be assessed using the Valence
Shell Electron Pair Repulsion (VSEPR) model.
• The description of the molecular geometry is based upon
the arrangement of the bonded atoms.
• The geometry of most common molecules is derived
from 5 basic shapes: linear (AB2), trigonal planar (AB3),
tetrahedral (AB4), trigonal bipyramidal (AB5), and
octahedral (AB6).
• In this notation, “A” denotes the central atom; “B” groups
may be bonded atoms or lone pairs.
Molecular Geometries
Molecular Geometries
Covalent Bond Properties
Bond Length
• The equilibrium bond length corresponds to the
minimum on the molecular potential energy curve
• The mutual attraction between the valence electrons
of each atom and the nucleus of the other pulls the
atoms together until the repulsion between the core
electrons of the atoms begins to dominate.
• The sum of the covalent radii of two bonded atoms
corresponds to the internuclear separation when the
core shells of the two atoms are in contact.
Table of Covalent Radii
Bond Strength
• The bond dissociation enthalpy (BDE) is
the standard reaction enthalpy for the
process:
A-B(g)  A(g) + B(g)
• The mean bond enthalpy is the average
BDE taken over a series of A-B bonds in
different molecules.
Table of Mean Bond Enthalpies
Estimation of DHrxn