RESONANCE STRUCTURES

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• RESONANCE STRUCTURES
• For certain molecules or molecular ions, two or
  more EQUIVALENT Lewis structures can be drawn.
• The equivalent structures have the same formal
  charge and are equal in energy.

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Three equivalent structures for the carbonate ion
all have the same formal charge.
These equivalent structures are called RESONANCE
STRUCTURES.
The double bond electron pair is DELOCALIZED over
all three C-O bonds and hence all three C-O bonds
are between pure-single and pure-double bonds.
Note This differs from the two structures for
N3- which had different bonding arrangements for
the atoms, and hence not equivalent.
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• The carbonate ion is NOT a mixture of the three
  structures, nor does any one of the three exist
  independently.
• The real structure is an average of the three
  structures, and properties of these bonds are
  such that they are intermediate between single
  and double bonds.
• All three C-O bonds identical, the bond lengths
  are between a single bond between C and O and a
  double bond between C and O
• Other examples of molecules that have more than
  one equivalent resonance structures are NO3-, SO3

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• EXCEPTIONS TO THE OCTET RULE
• Although the octet rule is very useful, it does
  not apply in all cases. There are three classes
  of exceptions
• 1) Compounds with odd number of electrons
• 2) Compounds with a shortage of electrons, less
• than the octet.
• 3) Compounds with an abundance of electrons,
• exceeding the octet

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• 1) Molecules with ODD number of valence electrons
• An example is nitrogen monoxide, NO.
• Molecules with an odd number of electrons cannot
  satisfy the octet rule for all its atoms.
• In this case the octet rule must be given up for
  one of the atoms by leaving an unpaired lone
  electron.
• ALL BONDING ELECTRONS MUST BE PAIRED.

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• NO is stable molecule with 11 valence electrons

In (i) N has an octet, but not O in (ii) O has
an octet but not N
Calculate the formal charge of each atom for (i)
and (ii) to determine which is the more favored
structure.
Structure (ii) is favored since the formal
charges on both atoms is zero in (i) N is -1 and
O is 1, with the more electronegative atom, O,
possessing a 1 formal charge
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• 2) Electron Deficient Molecules
• Example BF3

Lewis model predicts the following structure for
BF3
However, experimental evidence suggests that the
structure of BF3 is
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• (i) assigns a 1 formal charge to a very
  electronegative atom which is undesirable
• (ii) is the observed structure, even though B
  does not have an octet, the formal charges on all
  atoms are zero.

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• 3) Valence-Shell Expansion
• The elements in Groups IIIA to VII A in the third
  period and beyond show a tendency to surround
  themselves with more than 8 electrons.
• It is possible for these elements through
  valence shell expansion for the central atom to
  have more than 8 electrons
• In cases like this, after accounting for bonding
  electrons, assign lone pairs to the outer atoms
  to give them octets.
• If any electrons still remain, assign them to the
  central atoms as lone pairs.

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• Example SF6
• Number of valence electrons 6 (6x7) 48
• Total number of electrons needed for each atom to
  satisfy an octet 7x8 56
• Number of shared electrons 56-48 8
• 8 electrons, 4 bonds, is not enough bonds for SF6
• Hence, assign one bond pair to each bond, which
  for SF6 is 6.
• Then assign lone pairs to the outer atoms to give
  them an octet. If any electrons remain, assign
  them to the central atom.

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• Count all electrons
• (6x6) (6x2) 48 electrons

Formal charge on all atoms 0
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POCl3 - phosphoryl chloride
(ii)
(i)
i) non zero formal charges ii) valence shell
expansion for P
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Shapes of Molecules

• Molecules have definite shape i.e. definite
  geometries.
• For example, H2O is bent and CO2 is linear.

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• Lewis stuctures do not provide information on a
  molecules geometry, i.e. shape.
• VALENCE SHELL ELECTRON-PAIR REPULSION (VSEPR)
  THEORY predicts molecular shapes based on Lewis
  stuctures
• VSEPR derives from the realization that bonding
  and non-bonding electron pairs will locate
  themselves in space so as to minimize repulsive
  interactions.
• The geometry of the electron pairs, and the
  covalent bonds that form, can be predicted using
  VSEPR

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• Molecular geometry depends on the total number of
  electron pairs
• To use the VSEPR theory to predict molecular
  geometry, need to know how electron pairs can be
  arranged around a central atom so as to minimize
  electron-electron repulsion.
• An electron pairs arranges itself in
  three-dimensional space such that it is equally
  far apart from all the other electron pairs.
• The resulting arrangement of the electron pairs
  in a molecule depends on the NUMBER of electron
  pairs.
• Note For VSEPR, the electron pairs we refer to
  are the valence electrons

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For a central atom surrounded by TWO electron
pairs, the optimum geometry which minimizes
repulsion between the electrons pairs is LINEAR
For a central atom surrounded by THREE electron
pairs, the optimum geometry which minimizes
repulsion between the electrons pairs is TRIGONAL
PLANAR
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For a central atom surrounded by FOUR electron
pairs, the optimum geometry which minimizes
repulsion between the electrons pairs is
TETRAHEDRAL
For a central atom surrounded by FIVE electron
pairs, the optimum geometry which minimizes
repulsion between the electrons pairs TRIGONAL
BIPYRAMIDAL
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For a central atom surrounded by SIX electron
pairs, the optimum geometry which minimizes
repulsion between the electrons pairs OCTAHEDRAL
We have been talking about the location of the
electron pairs about the central atom. This
maybe different from molecular geometries i.e.
the structure or shape of a molecule
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• Water has a BENT structure with the H-O-H angle
  measured to be 104.5o.

If O is bonded to 2 H atoms, why is the structure
of water not linear?
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MOLECULAR GEOMETRY Molecular geometry is the
arrangement of ATOMS about the central atom. If
there are no non-bonding electrons, then the
electron geometry and the molecular geometry is
the same.

• If there are non-bonding electrons on the central
  atom then the molecular geometry differs from the
  electron geometry but is obtained from knowledge
  of the electron geometry about the central atom.

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• To determine the molecular geometry about the
  central atom of a molecule
• 1) Write the best Lewis structure for the
  molecule or ion
• 2) Count the total number of ATOMS attached to
  the central atom plus the number of NONBONDING
  ELECTRON PAIRS on the central atom .
• 3) Using the VSEPR geometries establish the
  ELECTRON geometry about the central atom.
• 4) Finally examine the placement of atoms and
  identify the molecular geometry.

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• Step 1 Write the best Lewis structure for the
  molecule
• Total number of valence electrons 116 8
• Number of electrons needed for each to complete
  an octet (remember for H need two electrons) 2
  2 8 12
• Shared electrons 12- 84 gt two bonds
• Number of unpaired electrons 4 gt 2 lone pairs
  on O
• Lewis structure of H2O is

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• 2) Count the total number of atoms attached to
  the central atom plus the number of non-bonding
  electron pairs on the central atom

For water, Total number of electron pairs
2
pairs of bonding electrons (bp) 2 pairs of
non-bonding electrons (np) 4 pairs
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3) Then using the VSEPR geometries establish the
electron geometry about the central atom.
According to VSEPR theory For a central atom
surrounded by four electron pairs, the optimum
geometry which minimizes repulsion between the
electrons pairs is tetrahedral
Arrange the two nonbonding electron pairs and the
two H atoms at the apices of the tetrahedron,
with O at the center
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4) Finally examine the placement of atoms and
identify the molecular geometry.
To determine the molecular geometry of water,
look at the arrangement of just the atoms (in
this case, H,O,H)
Looking at the arrangement of the atoms in H2O,
we see that the molecule is bent.
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• For a tetrahedral geometry, expect that the H-O-H
  angle is 109.5o
• The measured H-O-H angle is 104.5o, slightly
  smaller than predicted.
• Non-bonding electron pairs are more diffuse than
  bonding electron pairs, taking up more space.

Model
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• Draw the Lewis structure to determine the number
  of bonding electron pairs and number of
  nonbonding pairs on the central atom (a double or
  triple bond counts as a single electron pair when
  using VSEPR theory).
• Arrange the bonding and nonbonding electron pairs
  around the central atom in a geometry predicted
  by VSEPR.
• Position the atoms bonded to the central atom
  where the bonding electrons are positioned.
• Determine the molecular geometry based on the
  arrangement of the atoms

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• Two electron pairs
• Example CO2

Number of electron pairs around C 2 (count the
double bond as a single electron pair). Both
are bonding pairs VSEPR predicts a linear
molecular geometry for CO2
Model
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• Three electron pairs

The three electron pairs around the central B
atom are all bonding pairs. According to VSEPR
the three electron pairs around the central atom
are positioned at the vertices of an equilateral
triangle In the case of BF3, since the three
electron pairs are bonding pairs, the three F
atoms are positioned at the vertices of the
equilateral triangle.
Model
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• For VSEPR count two bonding electron pairs and
  one nonbonding electron pair.
• Hence the three electron pairs are situated at
  the vertices of an equilateral triangle.

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• The two O atoms are situated at two of the
  vertices and the nonbonding electron pair at the
  third.

The molecular geometry if NO2- is not trigonal
planar, but bent. The O-N-O angle is less than
120o due to the presence of the nonbonding pair
on N
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• Four electron pairs
• For molecules with four electron pairs, the
  electron pairs surround the central atom in a
  TETRAHEDRAL geometry.
• The molecular geometry depends on the number of
  bonding and nonbonding electron pairs on the
  central atom.
• Examples

Model
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• Five Electron Pairs
• Electronic geometry is trigonal bipyramidal.
• However, unlike the 2, 3 and 4 electron
  geometries all vertices of a trigonal bipyramid
  are NOT equivalent.
• There are two AXIAL positions and three
  EQUATORIAL positions.

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• Angle between
• axial-central atom-axial 180o
• axial-central atom-equatorial 90o
• equatorial -central atom- equatorial 120o

The positions of the nonbonding electron pairs
relative to each other and the bonding pairs have
to account for the non-equivalence of the axial
and equatorial positions. To position the
nonbonding and bonding electrons nb-nb
repulsion gt nb-b gt b-b
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• Hence, position nonbonding pairs first, then
  bonding pairs accounting for the repulsion
  between the electron pairs.

PCl5
Model
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• SF4

Position the nonbonding electron pair on S at the
equatorial position since this minimizes nb-b
repulsions
F
F
S
F
see-saw
F
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Other examples ClF3 - distorted T XeF2 - linear
Model
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• Six electron pairs
• Octahedral geometry
• All vertices are the equivalent, hence bonding
  and nonbonding electron pairs can be positioned
  at any vertex

Molecular geometry determined by number and
position at atoms and nonbonding electron pairs
Model
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• Examples

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• Molecular Properties
• Bond Lengths
• The bond distance is the distance between the two
  nuclei of the two atoms bonded together.
• Bond lengths depend on the elements and whether
  the bond between the two elements is single,
  double or triple
• Typically
• single bond length gt double bond length gt triple
  bond length

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• Bond Energies
• If two atoms in a diatomic molecule are pulled
  far enough apart the bond between the atoms
  breaks.
• The energy needed to break the bond is called the
  bond energy or bond dissociation energy.
• Bond energies depend on the nature of the atoms
  bonded together and whether the bond is single,
  double or triple.
• Typically single bond energies lt double bond
  energy lt triple bond energy

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• Dipole Moments
• The bonding electron pair in HCl is not shared
  equally between the H and Cl atoms.
• This is because Cl is more electronegative than
  H, and tends to attract the electron pair closer
  towards it.
• Hence the electron distribution for the bonding
  pair of electrons is not evenly distributed over
  both the H and Cl atoms, but is more concentrated
  over the Cl atom.
• The more electronegative atom develops a slight
  negative charge and the more electropositive atom
  a slight positive charge

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• This charge separation results in a DIPOLE
  MOMENT, and a POLAR HCl bond.
• The dipole moment is defined as the product of
  the magnitude of the separated charges and the
  distance between the centers of the charges.
• The dipole moment is a vector

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• The larger the electronegative difference between
  the bonded atoms, larger is the magnitude of the
  dipole moment and hence more polar is the bond.
• All heteronuclear diatomic molecules (A-B. where
  A?B) are polar molecules with the degree of
  polarity depending on the electronegativity
  differences between the two atoms.
• Heteronuclear diatomics have non-zero dipole
  moments
• All homonuclear diatomic molecules (A2) are
  non-polar, with zero dipole moments

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• Polyatomic molecules
• Bonds between atoms of different elements will
  each have a dipole moment associated with them.
• The magnitude of the dipole moment of a bond
  depends on the electronegativity differences
  between the bonded atoms.
• However, whether the molecule as a whole has a
  dipole moment depends on the geometry of the
  molecule.
• The dipole moment is a vector and hence the
  dipole moment of the molecule is the vector sum
  of the dipole moments of each bond in the
  molecule.

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• Methane, CH4

The net dipole moment for CH4 is zero
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• Water, H2O

Model