Covalent bonding #3

Covalent bonding
uThere are many different approaches to
describing covalent bonds
– which one you pick is often a compromise
between ease of use and accuracy
uMolecular orbital theory is the best
approach if good agreement with
experiment is needed
Lewis theory
uLewis’ theory of bonding was one one the
earliest to have any success
uBased on the octet rule
– main group elements like to have eight
electrons when they form compounds (except
hydrogen)
Bond order
uIn many cases Lewis structures can be used
to calculate bond orders that correlate well
with experimentally measured bond
strengths and lengths
– :N:::N: triple bond
– O::O double bond
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Formal charges
uA molecule may have more than one
plausible Lewis structure
uThe best Lewis structure is the one that has
the least charge separation
Resonance structures
uMolecules and ions with more than one
distinct but equivalent (by rotation or
reflection etc.) Lewis structures can occur
the real structure is usually an average of
these different resonance forms
Failures of the Lewis model
uA number of molecules with odd numbers
of electrons exist (no octet) e.g. NO
uAn atom may not have enough electrons to
complete its octet without having ridiculous
formal charges e.g. BF 3
uA central atom may clearly have more than
8 electrons e.g. SF 6
uO 2 is paramagnetic!!
2

Molecular Orbital theory
uIn principle, the electronic structure of
molecules can be worked out in the same
way as for atoms
– solve the schrodinger equation
uThis gives molecular orbitals rather than
atomic orbitals
uHowever, it is difficult to solve the
Schrodinger equation for molecular species
LCAO approximation
uGood approximations to the molecular
orbitals can be obtained by taking linear
combinations of atomic orbitals
Rules for use of MOs
uWhen two AOs mix to give MOs, two MOs
will be produced
uFor mixing AOs must have similar energies
uEach orbital can have two electrons max
uFill lowest energy orbitals first
uIf you have unpaired electrons they should
be spin parallel (Hund’s rule)
uBond order is number of bonding pairs
minus number of antibonding pairs
3

Period 1 diatomics
Overlapping p-orbitals
Oxygen and fluorine
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Nitrogen
Bond orders
Experimental verification
uUV-photoelectron spectroscopy can be used
to verify the MO diagrams.
– Molecules are ionized using monochromatic
light
» N 2(g) + h•---> N 2
+
(g) + e -
– The kinetic energy of the resulting
photoelectrons is measured
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PES spectrum for N 2
Heteronuclear diatomics
uMO diagrams for heteronuclear species are
constructed in a similar fashion to those for
homonuclear species
– However, the AO energies are different
Molecular shapes
uThe prediction of molecular shapes can be
done in a number of ways
– MO theory is very effective. However, it
requires complicated calculations
– VSEPR (Valence Shell Electron Pair
Repulsion) theory is good for main group
compounds and predictions can be made easily
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VSEPR
uVSEPR does not have a solid theoretical
foundation
uIt is based on the idea that pairs of valence
electrons (either bonding or lone pairs) will
try and avoid each other as much as
possible
– molecule adopts a geometry that allows
electron pairs to be as far apart as possible
Electron pair count and geometry
uTwo pairs on central atom - linear
uThree pairs on central atom - trigonal
uFour pairs on central atom - tetrahedral
uFive pairs on central atom - trigonal
bipyramidal
uSix pairs on central atom - octahedral
uSeven pairs - a number of possibilities
Electron counting
uFor molecules with single bonds to the
central atom count all valence electrons of
central atom plus one electron for each
ligand atom
– BCl 3 3 electrons from boron and 1 from each
chlorine so there are 3 pairs of electrons
uWhen there are double bonds present count
the four electrons associated with the
double bond as a single pair
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Linear geometry
Trigonal geometry
Tetrahedral geometry
8

TBP geometry
Octahedral geometry
Seven coordinate species
uIF 7 , XeF 6 , UF 7
2-
, NbF 7
2-
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Species that violate VSEPR
uLike nearly all truly useful sets of rules
there are exceptions
uSpecies that are sterically crowded often do
not obey VSEPR
» XeF 6 obeys VSEPR (7 pairs)
» TeCl 6
2-
does not (7 pairs but is octahedral)
Hybridization
uThe bonding around atoms with different
geometries is often pictures as consisting of
overlapping hybrid orbitals
– a hybrid orbital is a mixture of AOs
– mixing in different ways gives different
geometries
Different types of hybridization
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How useful are hybrid orbitals
uThe use of hybrid orbitals provides a
picture of the bonding around an atom
uHybridization arguments are not predictive,
just descriptive
uMO theory is predictive but complicated to
use
An MO view of H 2 O
uMO theory predicts the lone pairs and
bonds in H 2 O are not equivalent unlike the
hybridization view
Intermolecular forces
uAttractive interactions between molecules
allow the formation of molecular liquids
and solids
uThere are a number of different types of
intermolecular forces
– dispersion/London/van der Waals forces
– dipolar interactions
– hydrogen bonds
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Dispersion forces
uAll molecular and atomic species are
attracted to each other by dispersion forces
uFluctuations in charge distribution polarize
nearby atoms and molecules. This induced
dipole interacts with the original uneven
charge distribution
The strength of dispersion forces
uThe strength of the interaction depends on
– number of electrons in atom or molecule
– spatial extent of atom or molecule
uIncreasing electron count tends to increase
the strength of the interaction
– more charge to move around
uIncreasing spatial extent increase the
interaction
– can get better charge separation
The BPs of MH 4 species
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Interactions between polar molecules
uSome molecules have a permanent dipole
moment
– due to a low symmetry charge distribution
uTo get low symmetry charge distribution
you must have polar bonds
– uneven distribution of charge between two
bonded atoms
uPolar bonds do not guarantee a polar
molecule!
Electronegativity
uSome atoms are better at attracting
electrons than others
– the ability to attract electrons is called
electronegativity
uElectronegativity was quantified by Pauling
Dipole-dipole versus dispersion
uMolecular dipoles can give a significant
contribution to intermolecular forces
– CO (BP 82K) and N 2 (BP 77K)
uHowever, dispersion forces are usually
more important
– HCl (BP 188K) and HBr (BP 206 K)
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The effect of hydrogen bonding
Hydrogen bonding
uVery polar molecules (HF, H 2 O, NH 3 ) often
have anomalous properties
uThe interactions between molecules are
strong
– intermolecular distances drop down into the
region where a covalent contribution is to be
expected
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