VSEPR Theory

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VSEPR Theory

• Valence Shell Electron Pair Repulsion

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VSEPR THEORYAT THE CONCLUSION OF OUR TIME
TOGETHER, YOU SHOULD BE ABLE TO

1. Use VSEPR to predict molecular shape
2. Name the 6 basic shapes that have no unshared
   pairs of electrons
3. Name a few variations off of these basic shapes

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VSEPR Theory
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Molecular Shapes

• Lewis structures show which atoms are connected
  where, and by how many bonds, but they don't
  properly show 3-D shapes of molecules.
• To find the actual shape of a molecule, first
  draw the Lewis structure, and then use VSEPR
  Theory.

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Valence Shell Electron-Pair Repulsion Theory or
VSEPR

• Molecular Shape is determined by the repulsions
  of electron pairs
• Electron pairs around the central atom stay as
  far apart as possible.
• Electron Pair Geometry - based on number of
  regions of electron density
• Consider non-bonding (lone pairs) as well as
  bonding electrons. Unshared repel the most.
• Electron pairs in single, double and triple bonds
  are treated as single electron clouds.
• Molecular Geometry - based on the electron pair
  geometry, this is the shape of the molecule

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Electron-group Repulsions And The Five Basic
Molecular Shapes.
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LETS CONSIDER THESE BASIC SHAPES AND SOME
VARIATIONS OF THEM
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Linear

• 2 atoms attached to central atom
• 0 unshared pairs (lone pairs)
• Bond angle 180o
• Type AX2
• Ex. BeF2

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Linear

• Carbon dioxide
• CO2

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The Single Molecular Shape Of The Linear
Electron-group Arrangement.
Examples CO2, BeF2
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Trigonal Planar

• Boron Trifluoride
• BF3

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Trigonal Planar

• 3 atoms attached to central atom
• 0 lone pairs
• Bond angle 120o
• Type AX3
• Ex. AlF3

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The Two Molecular Shapes Of The Trigonal Planar
Electron-group Arrangement.
Examples H2CO, BCl3, NO3-, CO32-
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Factors Affecting Actual Bond Angles
Bond angles are consistent with theoretical
angles when the atoms attached to the central
atom are the same and when all electrons are
bonding electrons of the same order. Some
exceptions follow
1200
larger EN
1200
ideal
greater electron density
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Factors Affecting Actual Bond Angles
Lone pairs repel bonding pairs more strongly than
bonding pairs repel each other
950
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Trigonal Planar Variation
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The Second Molecular Shape Of The Trigonal Planar
Electron-group Arrangement.
Examples SO2, O3
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Bent

• Trigonal Planar variation 1
• 2 atoms attached to central atom
• 1 lone pair
• Bond angle lt120
• Type AX2E
• Ex. SO2

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Tetrahedral

• 4 atoms attached to central atom
• 0 lone pairs
• Bond angle 109.5o
• Type AX4
• Ex. CH4

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Tetrahedral

• Carbon tetrachloride
• CCl4

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2 Tetrahedral Variations
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The Three Molecular Shapes Of The Tetrahedral
Electron-group Arrangement.
Examples CH4, SO42-
NH3 PF3
H2O OF2
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Trigonal Pyramidal

• Tetrahedral variation 1
• 3 atoms attached to central atom
• 1 lone pair
• Bond angle 107o
• Type AX3E
• Ex. NH3

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Trigonal Pyramidal

• Nitrogen trifluoride
• NF3

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Bent

• Tetrahedral variation 2
• 2 atoms attached to central atom
• 2 lone pairs
• Bond angle 104.5o
• Type AX2E2
• Ex. H2O

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Bent

• Chlorine difluoride ion
• ClF2

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Remember the 3 exceptions to the octet rule?

• Molecules with atoms near the boundary between
  metals and nonmetals will tend to have less than
  an octet on the central atom. (i.e. B, Be, Al,
  Ga)
• Molecules with a central atom with electrons in
  the 3rd period and beyond will sometimes have
  more than an octet on the central atom, up to 12,
  called an extended or expanded octet.
• Molecules with an odd number of electrons

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5 Bond Sites on the Central Atom
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Trigonal Bipyramidal

• 5 atoms attached to central atom
• 0 lone pairs
• Bond angle
• equatorial -gt 120o
• axial -gt 90o
• Type AX5
• Ex. PF5

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Trigonal Bipyramidal

• Antimony Pentafluoride
• SbF5

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3 Bipyramidal Variations
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The Four Molecular Shapes Of The Trigonal
Bipyramidal Electron-group Arrangement.
PF5 AsF5
SF4 XeO2F2
XeF2 I3-
ClF3 BrF3
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See Saw

• Trigonal Bipyrimid Variation 1
• Sulfur tetrafluoride
• SF4

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T-Shaped

• Trigonal Bipyramid Variation 2
• Chlorine tribromide

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Linear

• Trigonal Bipyramid Variation 3
• Xenon difluoride
• XeF2

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6 Bond Sites on the Central Atom
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Octahedral

• 6 atoms attached to central atom
• 0 lone pairs
• Bond angle 90o
• Type AX6
• Ex. SF6

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Octahedral

• Sulfur hexafluoride
• SF6

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2 Octahedral Variations
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The Three Molecular Shapes Of The Octahedral
Electron-group Arrangement.
SF6 IOF5
IF5 XeOF4
XeF4 (BrF4)-
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Square Pyramidal

• Octahedral Variation 1
• Chlorine pentafluoride
• ClF5

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Square Planar

• Octahedral Variation 2
• Xenon tetrafluoride
• XeF4

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Octahedral

• Do not need to know
• T-shape
• Linear

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Lets Review VSEPR Theory

• Predicts the molecular shape of a bonded molecule
• Electrons around the central atom arrange
  themselves as far apart from each other as
  possible
• Unshared pairs of electrons (lone pairs) on the
  central atom repel the most
• So only look at what is connected to the central
  atom

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Remember the 3 exceptions to the octet rule?

• Molecules with atoms near the boundary between
  metals and nonmetals will tend to have less than
  an octet on the central atom. (i.e. B, Be, Al,
  Ga)
• Molecules with a central atom with electrons in
  the 3rd period and beyond will sometimes have
  more than an octet on the central atom, up to 12,
  called an extended or expanded octet.
• Molecules with an odd number of electrons

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Bent

• Nitrogen dioxide
• NO2

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VSEPR THEORYLETS SEE IF YOU CAN

1. Use VSEPR to predict molecular shape
2. Name the 6 basic shapes that have no unshared
   pairs of electrons
3. Name a few variations off of these basic shapes

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Sample Problems
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The Steps In Determining A Molecular Shape.
Molecular formula
Step 1
Lewis structure
Count all e- groups around central atom (A)
Step 2
Electron-group arrangement
Step 3
Note lone pairs and double bonds
Bond angles
Count bonding and nonbonding e- groups separately.
Step 4
Molecular shape (AXmEn)
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Review of Lewis Structures

• Step 1 Count the number of valence electrons.
  For a neutral molecule this is equal to the
  number of valence electrons of the constituent
  atoms.
• Example (CH3NO2)Each hydrogen contributes 1
  valence electron. Each carbon contributes 4,
  nitrogen 5, and each oxygen 6 for a total of 24.

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Review of Lewis Structures

• Step 2 Connect the atoms by a covalent bond
  represented by a dash.
• ExampleMethyl nitrite has the partial
  structure

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Review of Lewis Structures

• Step 3 Subtract the number of electrons in
  bonds from the total number of valence electrons.
• Example24 valence electrons 12 electrons in
  bonds. Therefore, 12 more electrons to assign.

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Review of Lewis Structures

• Step 4 Add electrons in pairs so that as many
  atoms as possible have 8 electrons. Start with
  the most electronegative atom.
• Example The remaining 12 electrons in methyl
  nitrite are added as 6 pairs.

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Review of Lewis Structures

• Step 5 If an atom lacks an octet, use electron
  pairs on an adjacent atom to form a double or
  triple bond.
• Example There are 2 ways this can be done.

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Review of Lewis Structures

• Step 6 Calculate formal charges.
• Example The left structure has formal charges
  that are greater than 0. Therefore it is a less
  stable Lewis structure.

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SAMPLE PROBLEM
Predicting Molecular Shapes with Two, Three, or
Four Electron Groups
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The shape is based upon the tetrahedral
arrangement.
The F-P-F bond angles should be lt109.50 due to
the repulsion of the nonbonding electron pair.
The final shape is trigonal pyramidal.
lt109.50
The type of shape is AX3E
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(b) For COCl2, C has the lowest EN and will be
the center atom.
There are 24 valence e-, 3 atoms attached to the
center atom.
C does not have an octet a pair of nonbonding
electrons will move in from the O to make a
double bond.
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The shape for an atom with three atom attachments
and no nonbonding pairs on the central atom is
trigonal planar.
The Cl-C-Cl bond angle will be less than 1200 due
to the electron density of the CO.
Type AX3
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SAMPLE PROBLEM
Predicting Molecular Shapes with Five or Six
Electron Groups
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(b) BrF5 - 42 valence e- 5 bonding pairs and 1
nonbonding pair on central atom. Shape is
AX5E, square pyramidal.
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