CHM 110 CHAPTER 10: The Shape of Molecules

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CHM 110CHAPTER 10 The Shape of Molecules

• Dr. Floyd Beckford
• Lyon College

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ELECTRON-DOT STRUCTURES

• Bonding involves only the valence electrons
• Electron sharing can be represented by electron-
• dot structures

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• Writing electron-dot structures
• 1. Select a reasonable (symmetrical) skeleton
• (a) Central element is least electronegative
• (b) Oxygen atoms do not bond to each other
• 2. Find total number of electrons in molecule
• (a) Add one e- for each negative charge
• (b) Subtract one e- for each positive charge

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3. Form single bonds between the atoms - assign
remaining electrons to form octets on all
atoms 4. If any electron remain, place them the
central atom as lone pairs 5. If central atom
does not have a filled octet, form multiple bonds
with a neighboring atom
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RESONANCE

• Consider the Lewis structure for CO32-

• The three individual forms are called resonance
• structures
• The correct structure is a blend of all three
  the
• resonance hybrid

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FORMAL CHARGES

• It may be possible to draw more than one Lewis
• The likely structure can be determined from the
• concept of formal charges
• Formal charge the charge on an atom in a
• molecule or polyatomic ion
• Most favorable structure has atoms with zero or
• near-zero formal charges on each atom

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• Guidelines for assigning formal charges
• (a) In a molecule the sum of the formal charges
• is zero
• (b) In a polyatomic ion, the sum of the formal
• charges is equal to the charge on the ion
• FC Group number ( of bonds) ( of
• unshared electrons)

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• Not all compounds obey the octet rule
• e.g. BF3, BeCl2
• - these are examples of electron deficient
  species
• e.g. NO2
• - this is an example of an odd-electron species
• e.g. SF6
• - an example of a compound with the central
  atom having an expanded valence shell

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MOLECULAR SHAPESVSEPR MODEL

• Lewis structures are not true representations of
  
• the molecular shape
• Molecular shape can be predicted from the
• valence-shell electron-pair repulsion model
• In this model
• 1. Count number of electron clouds around the
• central atom

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• 2. Arrange the electron clouds around the central
• atom so that they are as far apart as possible
• This arrangement is called the atoms electronic
• geometry
• Molecular geometry arrangement of the bonded
• atoms
• Double and triple bonds count as one

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Two charge clouds e.g. CO2 - electron clouds
180 apart give rise to LINEAR molecules

• Three charge clouds
• e.g. BF3
• Most stable if electron clouds are 120 apart

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• Give TRIGONAL PLANAR electronic geometry

BENT molecular geometry

• Four charge clouds
• e.g. CH4, NH3, H2O
• All have TETRAHEDRAL electronic geometry
• all angles are 109

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• Molecular geometry is different

• Bond angles are 107 and 104.5 in ammonia and
• water respectively

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• Five charge clouds
• e.g. PCl5
• Have TRIGONAL BIPYRAMIDAL electronic
• geometry

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• In general (for molecular geometry and L lone
• pair)

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• Six charge clouds
• OCTAHEDRAL electronic
• geometry bond angles of
• 90 and 180

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BOND POLARITY

• Recall the nature of bond in HCl

• The separation of charge in a polar covalent
• bond creates an electric dipole

• A molecule typically has a number of polar
• bonds

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• The arrangement of the bonds determines if
• the molecule is polar
• For a molecule to be polar
• 1. There must be at least one polar bond or one
• lone pair on the central atom
• and
• 2. The polar bonds, if more than one, must not be
• symmetrically arranged
• or

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• If there are two or more lone pairs on the
• central atom, they must not be symmetrically
• arranged

nonpolar
polar
nonpolar