14.1 Shapes of molecules and ions (HL)

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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.

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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

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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.

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• 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_content/chem20/covm
  olec/exceptns.html

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Trigonal Bipyramidal (5 pairs of V.E.)
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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

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Octahedral (6 pairs of V.E.)
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BrF5 is square pyramidal
SF6 is octahedral
XeF4 is square planar
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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)

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• 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)

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Practice

• ClF3
• PF5
• XeO2F2
• SOF4
• SCl6
• IF4
• ICl4-

• T-shaped
• Trigonal bipyramidal
• Seesaw
• Trigonal bipyramidal
• Octahedral
• Seesaw
• Square planar

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14.2 Hybridization.

• 14.2.1 Describe s (sigma) and p (pi) bonds
• 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).

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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

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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?

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• One of carbons 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)

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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).

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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
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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).

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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.

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• 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.

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• 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 s-bond and 2 p (pi) bonds)

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Sigma bond (s-bond)

• When s and/or hybrid orbitals overlap 'end-on',
  sigma bonds (s) 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, s bond is shown

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Sigma bond (s-bond)

• results from head-on overlap of orbitals
• electron density is symmetric about the
  internuclear axis between nuclei.

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p (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 s plus one p
  bond,

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p (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.

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Predicting shape

• The shape is dictated by the s-bonds and the
  non-bonding electron pairs (lone pairs)
• p-bonds do not affect the shape of the molecule
  (double bonds or triple bonds)
• Thats why we refer to bonding sites when using
  VSEPR, not paying attention to whether it was
  single, double or triple bonded.

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14.3 Delocalization of electrons

• 14.3.1 Describe the delocalization of (pi) p-
  electrons and explain how this can account for
  the structure of some species

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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

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Resonance structures

• the nitrate ion can be viewed as if it resonates
  between the three different structures above.
• Nitrate doesnt change from one to the next, but
  behaves as a combination of all structures

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• 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.

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Practice

• Try to show the individual Lewis structures for
  the HCO3- ion
• Show its resonance structure too

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Practice drawing these resonance structures

• NO3-
• NO2-
• CO32-
• O3
• RCOO-
• Benzene (C6H6)

• TOK
• Kekule claimed that the inspiration for the
  cyclic structure of benzene came from a dream.
• What role do the less rational ways of knowing
  play in the acquistion of scientific knowledge?

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Bibliography and sites to visit

• http//www.tutorvista.com/content/chemistry/chemis
  try-iii/chemical-bonding/types-covalent-bonds.php
• Good site on types of covalent bonds
• http//www.mikeblaber.org/oldwine/chm1045/notes/Ge
  ometry/VSEPR/Geom02.htm
• Used for expanded valence shell pictures
• http//www.kentchemistry.com/links/bonding/lewisdo
  tstruct.htm
• Puts the lewis diagrams together and explain
  them. Including expanded shell

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• http//www.mpcfaculty.net/mark_bishop/resonance.ht
  m
• Resonance structures pictures and notes
• http//en.wikipedia.org/wiki/Delocalization
• Notes on delocalisation of electrons
• http//www.steve.gb.com/science/atomic_structure.h
  tml
• Amazing website for hybrid orbitals
• http//library.thinkquest.org/C006669/data/Chem/bo
  nding/shapes.html
• Good review of all shapes