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					14.1 Shapes of molecules
      and ions (HL)
 14.1.1 State and predict the
 shape and bond angles using
 the VSEPR theory for 5 and 6
 negative charge centers.
        Molecules with more than 4
              electron pairs
   Molecules with more
    than 8 valence
    electrons [expanded
    valence shell]
   Form when an atom
    can ‘promote’ one of
    more electron from a
    doubly filled s- or p-
    orbital into an unfilled
    low energy d-orbital
   Only in period 3 or
    higher because that is
    where unused d-
    orbitals begin
 Why does this ‘promotion’ occur?
 When   atoms absorb energy (heat,
  electricity, etc…)their electrons
  become excited and move from a
  lower energy level orbital to a
  slightly higher one.
 How many new bonding sites formed
  depends on how many valence
  electrons are excited.
 Exceptions to the octet rule. Shows
  sulphur achieving 8, 10 and 12
  valence electrons due to energy
  input and excited electrons.
 http://www.saskschools.ca/curr_cont
  ent/chem20/covmolec/exceptns.html
Trigonal Bipyramidal (5 pairs of V.E.)
       Trigonal Bipyramidal
 Normally would have 3 bp, but the lone
  pair has moved from the p-orbital to
  include the d-orbital, allowing for 2
  additional bonding sites.
 Ex: PCl5
Octahedral (6 pairs of V.E.)
BrF5 is square pyramidal



                           SF6 is octahedral




XeF4 is square planar
             Bond angles
 In general, the greater the bond
  angle, the weaker the repulsions.
 Equatorial- equatorial (120 o)
  repulsions are weaker than axial-
  equatorial (90o) repulsions.
  – Equatorial: lie on the trigonal plane
    (straight across)
  – Axial: lies above and below the trigonal
    plane (up and down)
   Remember that lone
    pairs cause more
    repulsion than
    bonding sites, so
    expect the bond
    angle to be changed
    should there be lone
    pairs, or double or
    triple bonds
    involved (multiple
    bonds also cause
    more repulsion than
    expected)
              Practice:
1.   ClF3          1.   T-shaped
2.   PF5           2.   Trigonal
3.   XeO2F2             bipyramidal
4.   SOF4          3.   Seesaw
5.   SCl6          4.   Trigonal
6.   IF4+               bipyramidal
7.   ICl4-         5.   Octahedral
                   6.   Seesaw
                   7.   Square planar
 14.2 Hybridization
                    . and π (pi) bonds
14.2.1 Describe σ (sigma)
14.2.2 State and explain the meaning of the
              term hybridization
  14.2.3 Discuss the relationships between
Lewis structures, molecular shapes and types
        of hybridization (sp, sp2, sp3).
            hybridization
 the concept of mixing atomic orbitals to
  form new hybrid orbitals
 Used to help explain some atomic bonding
  properties and the shape of molecular
  orbitals for molecules.
 The valence orbitals (outermost s and p
  orbitals) are hybridised (mathematically
  mixed) before bonding, converting some
  of the dissimilar s and p orbitals into
  identical hybrid spn orbitals
 We must know sp, sp2, and sp3 hydrid
  orbitals
             Hybrid orbitals
 Carbon has 4
  valence electrons.
 2 electrons paired
  up in the s-orbital,
  and 2 electrons
  unpaired in the p-
  orbital.
 So why does it
  commonly make 4
  bonding sites?
 One  of carbon’s paired s-orbital
  electrons is ‘promoted’ to the empty
  p-orbital
 This produces a carbon in an excited
  state which has 4 unpaired electrons
  (4 equivalent bonding sites)
sp3 hybrid orbital
            formed by mixing the
             outermost s- and all
             three outermost p-
             orbitals to form four
             sp3 hybrids.
            The furthest these
             four [negatively
             charged, and
             therefore repulsive]
             orbitals can get from
             each other is the
             corners of a
             tetrahedron (109°).
Overlap four s-orbitals from four hydrogens (blue)
with four sp3 hybrids on carbon leads to formation
of bonds, each containing one electron from the
        carbon and one from the hydrogen
        Examples of sp3 hybrids




   Methane, ammonia, water and hydrogen fluoride.
   Note that the orbitals not involved in bonding to
    hydrogen are still hybridised, but end up as lone
    pairs of electrons (symbolised by the two dots in
    the diagram above).
sp2 hybrid orbital
          formed when only
           one s- and two p-
           orbitals are
           involved.
          This leaves one
           remaining p
           orbital, which may
           be involved in
           forming a double
           bond.
   The furthest these orbitals can get from one
    another is a trigonal bipyramid, with the sp2
    hybrids arranged at 120° to each other in a
    plane.
   This is characteristic of molecules with double
    bonds.
   Finally, sp hybrids are
    formed using just one s
    and one p orbital.
   Two sp hybrids are formed
    from them, and the two p-
    orbitals remaining may
    contribute to a triple bond.
   These arrange themselves
    at the corners of an
    octahedron, with the two
    sp hybrids diametrically
    opposite one another.
   sp hybridisation is
    characteristic of the triple
    bond. (1 σ-bond and 2 π
    (pi) bonds)
       Sigma bond (σ-bond)
 When s and/or hybrid orbitals overlap
  'end-on', sigma bonds (σ) are formed
 They have a single area of electron
  density between the nuclei of the two
  atoms whose orbitals are overlapping.
 In the diagrams below, σ bond is shown
         Sigma bond (σ-bond)

 results from head-on overlap of orbitals
 electron density is symmetric about the
  internuclear axis: between nuclei.
                 π (pi) bonds
   p orbitals can overlap sideways too: when this
    happens two lobes of electron density are formed
    between the atoms.
   From the diagram, you can see that the double
    bond in ethene is composed of one σ plus one π
    bond,
π (pi) bonds

         results from sideways
          overlap of orbitals
         bonds resulting from
          the combination of
          parallel p orbitals
         electron density is
          above and below the
          internuclear axis.
          Predicting shape
 The shape is dictated by the σ-bonds
  and the non-bonding electron pairs
  (lone pairs)
 π-bonds do not affect the shape of
  the molecule (double bonds or triple
  bonds)
  – That’s why we refer to bonding sites
    when using VSEPR, not paying attention
    to whether it was single, double or triple
    bonded.
14.3 Delocalization of
      electrons
14.3.1 Describe the delocalization
 of (pi) π- electrons and explain
   how this can account for the
    structure of some species
       Delocalised electrons
 The term 'delocalised' refers to an electron
  which is not 'attached' to a particular atom
  or to a specific bond.
 Delocalized electrons are contained within
  an orbital that extends over several
  adjacent atoms.
 Classically, delocalized electrons can be
  found in double bonds and in aromatic
  systems
 Double bonds = 1 sigma and 1 pi bond
 Delocalisation is often represented with
  resonance structures or resonance hybrid
       Resonance structures



 the nitrate ion can be viewed as if it
  resonates between the three different
  structures above.
 Nitrate doesn’t change from one to the
  next, but behaves as a combination of all
  structures
 Resonance is possible whenever a Lewis
  structure has a multiple bond and an
  adjacent atom with at least one lone pair.
 The following is the general form for
  resonance in a structure of this type.
               Practice
 Try to show the
  individual Lewis
  structures for the
  HCO3- ion
 Show its resonance
  structure too
 Practice drawing these resonance
             structures:
1.   NO3-                          TOK
                         Kekule claimed that
2.   NO2-                 the inspiration for the
3.   CO32-                cyclic structure of
                          benzene came from a
4.   O3                   dream.
5.   RCOO-               What role do the less
                          rational ways of
6.   Benzene (C6H6)       knowing play in the
                          acquistion of scientific
                          knowledge?
    Bibliography and sites to visit
   http://www.tutorvista.com/content/chemi
    stry/chemistry-iii/chemical-bonding/types-
    covalent-bonds.php
    – Good site on types of covalent bonds
   http://www.mikeblaber.org/oldwine/chm1
    045/notes/Geometry/VSEPR/Geom02.htm
    – Used for expanded valence shell pictures
   http://www.kentchemistry.com/links/bond
    ing/lewisdotstruct.htm
    – Puts the lewis diagrams together and explain
      them. Including expanded shell
   http://www.mpcfaculty.net/mark_bishop/r
    esonance.htm
    – Resonance structures pictures and notes
   http://en.wikipedia.org/wiki/Delocalization
    – Notes on delocalisation of electrons
   http://www.steve.gb.com/science/atomic_
    structure.html
    – Amazing website for hybrid orbitals
   http://library.thinkquest.org/C006669/dat
    a/Chem/bonding/shapes.html
    – Good review of all shapes

				
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