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Recap Resonance. Elements from the third period and beyond form compounds with 8 or more electrons around the atom. This flexibility may result in quite different resonance structure being possible. - PowerPoint PPT Presentation
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Recap Resonance• Elements from the third period and beyond
form compounds with 8 or more electrons around the atom.
• This flexibility may result in quite different resonance structure being possible.
• The resonance structure(s) with the greatest contribution to the actual structure can be identified using the valency of oxygen as a guide.
Example: ClO4-
O
ClO
O
O
O
ClO
O
O
O
ClO
O
O
O
ClO
O
O
O
ClO
O
O
structure with greatestcontribution to actual structure
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VSEPR TheoryLewis structures give bonding arrangements but do
not imply any molecular shape. For this we use:
Valence Shell Electron Pair Repulsion Theory
This relies on minimising repulsion between areas of electrons (bond pairs and lone pairs) around the central atom.
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VSEPR Theory1. Draw Lewis Structure.
2. Count number of electron pairs.• Count both bonding pairs and non-bonding
pairs.• Count multiple bonds as only one area of
electrons.
3. Determine the arrangement of electron pairs.• Electron pairs want to be as far away from
each other as possible.
4. Use atom positions to name molecular geometry.• This is the atom positions.
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Electron Pair Arrangements
• Two electron pairs:– Atoms at the opposite ends of a line.– 180 degrees between areas of
electrons.– Called linear.– eg CO2
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Electron Pair Arrangements
• Three electron pairs:– Atoms at the corners of a triangle.– 120 degrees between electron pairs.– Called trigonal planar.– Eg BF3
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Electron Pair Arrangements
• Four electron pairs:– Atoms at the corners of a
tetrahedron.– 109.5 degrees between electron
pairs.– Called tetrahedral.– Eg CH4
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Remove one arm from the electron pair arrangement for each lone pair present.
Trigonal Planar (3 e- pairs)
Molecular Geometry
Figure 10.4 Silberberg
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Tetrahedral (4 e- pairs)
Molecular Geometry
Figure 10.5 Silberberg
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Molecular Geometry• Repulsion: lone pair-lone pair > lone
pair-bond pair > bond pair-bond pair.
109.5 107 104.5
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Molecular GeometryFig
ure
10.9
Silb
erbe
rg
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Molecular Geometry - Example
• Molecules with multiple bonds eg COCl2 total 24 e-
3 areas of electrons about C, so trigonal planar arrangement of electrons
No lone pairs so molecular geometry is also trigonal planar
~120
Cl C
O
Cl
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Molecular Geometry - Example
• Cases when there is no single central atom– Just apply the VSEPR rules to each central
atom in turn.
C C
H
H
H
H3 areas of electrons about each C, so trigonal planar arrangement of electrons about each C
~120
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Dipole Moments• Any bond between two different atoms
will be polar.• A molecule has a permanent dipole
moment if it contains polar bonds and it is not a symmetrical shape.
• Note: Cations and anions are not polar – the overall charge overwhelms any local d+ vs d- effects.
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Dipole Moments• Polar molecules
• Non-polar moleculesHF H2O
CHCl3
N2 CO2
CCl4
• By the end of this lecture, you should:− work out the number of bonding and non-
bonding pairs from the Lewis structure of a molecule
− predict the distribution of these pairs around an atom
− predict and describe the molecular shape− determine if a permanent dipole exists
− be able to complete the worksheet (if you haven’t already done so…)
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Learning Outcomes:
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Questions to complete for next lecture:
1. Draw the shapes of the following molecules and ions and give approximate bond angles(a)BH3
(b)NH4+
(c) CS2
(d)CH2O(e)CH3Cl(f) H3O+
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Questions to complete for next lecture:
2. What are the approximate C-C-C bond angles in the two molecules below?
3. Are these molecules flat?
