Chapter 10 The Shapes of Molecules Lecture Notes by K. Marr (Silberberg 3rd Edition)

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Chapter 10 The Shapes of Molecules Lecture
Notes by K. Marr (Silberberg 3rd Edition)

• 10.1 Depicting Molecules and Ions with Lewis
  Structures
• 10.2 Using Lewis Structures and Bond Energies to
  Calculate Heats of Reaction
• 10.3 Valence-Shell Electron-Pair Repulsion
  (VSEPR) Theory and Molecular Shape
• 10.4 Molecular Shape and Molecular Polarity

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Lewis Structures..

• Indicate the kind of bonding and which atoms are
  bonded in molecules and polyatomic ions
• Do NOT indicate the molecular shape or structure.
  However.
• VSEPR theory uses Lewis structures to predict 3-D
  structure

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Guidelines for Writing Lewis Structures

1. Decide which atoms are bonded
2. Count all valence electrons (account for the
   charge of ions!!)
3. Place 2 electrons in each bond
4. Complete the octets of the atoms attached to the
   central atom by adding electrons in pairs
5. Place any remaining electrons on the central atom
   in pairs
6. If the central atom does not have an octet, form
   double bonds, or if necessary, a triple bond.
7. Write the Lewis Structures for ClF5, TeF4, CO32-,
   CH3COO1-

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The Octet Rule is Often Violated

• H, Be, B, Al violate the octet rule (lt 8 valence
  electrons)
• e.g. BeCl2, BH3, AlCl3
• Nonmetals with a valence shell greater than n
  2 (e.g. P, Cl, Br, I, etc.)
• May violate the octet rule when they are the
  CENTRAL atom (e.g. ClF5 )
• How can they do this?
• Why doesnt Fluorine violate the octet rule?

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Lewis Structures for Organic Compounds

• Alkanes CnH2n2
• Methane, Ethane, Propane, Butane, Pentane,
  Hexane
• What are isomers?
• Alkenes CnH2n have double bond(s)
• One double bond Ethene (ethylene), Propene
  (propylene)
• Alcohols CnH2n1OH have hydroxyl group(s)
• methanol, ethanol
• Carboxylic Acids CnH2n1COOH have carboxyl
  group(s)
• Methanoic acid (formic acid), HCOOH
• Ethanoic acid (acetic acid, CH3COOH

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Using Formal Charge to Select the Favored Lewis
Structure

• Sometimes more than one Lewis Structure is
  possible for a compound e.g. sulfuric acid,
  H2SO4 phosphate ion, PO4-3
• Formal Charge
• Apparent charge on a bonded atom
• An atom owns all of its nonbonding electrons
  and half of its bonding electrons.
• The Lewis Structure with the lowest total formal
  charge is favored
• Formal charge of atom
• valence e- unshared e- 1/2 shared
  e-
• OR
• F.C. of valence e- - of unshared
  bonds formed

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Use of Formal Charge to Select the Favored Lewis
Structure

• Formal Charge
• Apparent charge on a bonded atom
• An atom owns all of its nonbonding electrons
  and half of its bonding electrons.
• The Lewis Structure with the lowest total formal
  charge is favored
• Formal charge of atom
• valence e- unshared e- 1/2 shared
  e-
• OR
• F.C. of valence e- - of unshared
  bonds formed

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Use of Formal Charge to Select the Favored Lewis
Structure

• Use formal charge to determine the correct Lewis
  structure for
• sulfuric acid, H2SO4
• phosphate ion, PO4-3
• Recall
• F.C. valence e- unshared e- 1/2
  shared e-
• OR
• F.C. of valence e- - of unshared
  bonds formed

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Formal Charge Three criteria for choosing the
more important structure

• Smaller formal charges (either positive or
  negative) are preferable to larger charges
• Avoid like charges ( or - - ) on adjacent
  atoms
• A more negative formal charge should exist on an
  atom with a larger EN value.

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ResonanceWhen Lewis Structures Fail.....

• Write the Lewis Structure for the nitrate ion,
  NO3-
• Based on your Lewis structure, what kind of
  bonding would be expected ?
• Experimental measurements indicate....
• All bond lengths and energies are the same!!
  (B.O. 1.33)
• The NO3- is a Resonance Hybrid of 3 different
  Lewis structures....
• Just as mule is neither a horse or a donkey, none
  of the 3 structures represent NO3-

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Resonance Hybrids

• Each resonance structure does not actually
  exist!!
• The actual molecule or ion is a hybrid or average
  of each resonance structure
• Electron-Pair Delocalization
• Each bonding electron pair is delocalized or
  spread over the entire molecule or ion.
• Results in identical bonds with extra stability
  since electron repulsions reduced

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Resonance Structures Practice Makes Perfect?

• Draw the resonance structures for the nitrite
  ion, NO2 - and the phosphite ion, PO3-3
• How do you know when to use resonance?
• How do you know how many resonance structures are
  possible?
• Draw the Lewis structures for ......
• The oxalate ion, C2O4-2
• Benzene, C6H6
• Benzene has a hexagonal ring structure

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Using Bond Energies to Calculate Heats of
Reaction, DHrxn

• Lewis structures can be used to calculate DHrxn
• For a reaction to occur.
• Bonds within the reactants must be broken
  (endothermic)
• Bonds within the reactants must be made
  (exothermic)
• DHrxn S DHreactant bonds broken S DHproduct
  bonds formed
• Reactants and products must be in gaseous state!!
  Why??

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Using Bond Energies to Calculate Heats of
Rxn DHrxn S DHreactant bonds broken S
DHproduct bonds formed
e.g. CH4 (g) 2 O2 (g) ? CO2 (g) 2 H2O (g)
DH0rxn -818 kJ/mol
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Figure 10.3
Using bond energies to calculate DH0comb. of
Methane, CH4
BOND BREAKAGE
4BE(C-H) 1652kJ
2BE(O2) 996kJ
DH0(bond breaking) 2648kJ
BOND FORMATION
4-BE(O-H) -1868kJ
Enthalpy,H
DH0(bond forming) -3466kJ
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Examples Using Bond Energies to Calculate Heats
of Reaction, DHrxn

• Use bond energies (see table 9.2, page 340 3rd
  ed) to calculate in kJ/mole the
• Standard heat of formation of water (compare your
  answer with Appendix Bthey should be the same)
• Standard heat of combustion of propane, C3H8
  (ans. -2042 kJ/mol)
• Now use standard heats of formation, DHof, to
  calculate the heat of combustion of propane, C3H8
  (ans. - 2043 kJ/mol)

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Predicting the Shapes of Molecules VSEPR Theory

• Valence Shell Electron Pair Repulsion Theory
• In order to limit electrostatic repulsion,
  electron pairs in the orbitals around the central
  atom stay as far apart as possible

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VSEPR A balloon analogy for the mutual repulsion
of electron groups.
Linear
Trigonal Planar
Tetrahedral
Trigonal Bipyramidal
Octahedral
Figure 10.4
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VSEPR Theory

• The Number of Electron Pairs around the Central
  Atom Determine Molecular Geometry....
• 2 bonding pairs ? linear (Bond angle 180o)
• 3 bonding pairs ? planar triangle (Bond angle
  120o)
• 4 bonding pairs ? tetrahedral (Bond angle
  109.5o)
• 5 bonding pairs ? trigonal bipyramidal (Bond
  angles 90o and 120o )
• 6 bonding pairs ? octahedral (Bond angle 90o)

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Figure 10.5
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Predicting Molecular Geometry

• Use Lewis structures and VSEPR Theory to explain
  the following molecular geometries....
• H2O and SnCl2
• Are they Bent or V-shaped molecules?
• BeCl2 and CO2
• Bent or linear molecules?
• Treat double bonds as if only one pair...Why?

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Predicting Molecular Geometry

• Use Lewis structures and VSEPR Theory to predict
  the following molecular geometries....
• BH3
• NH3
• ClF3
• ClF3 T-Shaped and NOT trigonal planar. Why??
• Nonbonding pairs take up more space than bonding
  electrons......why?
• Therefore, nonbonding pairs need to be separated
  as much as possible.

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Predicting Molecular Geometry

• Use Lewis structures and VSEPR Theory to predict
  the following molecular geometries....
• CH4 and PO43- (Ans. Tetrahedral)
• XeF4 (Ans. Square
  planar. Why not tetrahedral?)
• PCl5 (Ans. Trigonal
  bipyramidal)
• BrF5 (Ans. Square
  pyramidal)

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SAMPLE PROBLEM 10.9
Predicting Molecular Shapes with More Than One
Central Atom
SOLUTION
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Predicting Molecular Shapes with More Than One
Central Atom
The tetrahedral centers of ethanol.
Figure 10.13
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Figure 10.9
Lewis structures and molecular shapes
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Molecular Polarity

• Influences Chemical and Physical Properties
• Polar molecules have higher MPs and BPs than
  nonpolar molecules.....Why?
• Magnitude of Dipole moment influences MP and BP
• e.g. H2O vs H2S
• Solubility Like dissolves Like
• Polar solutes dissolve in polar solvents
• Nonpolar solutes dissolve in nonpolar solvents

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Nonpolar Molecules

• Any Molecule with only nonpolar bonds
• e.g. F2 and C8H18
• Symmetrical Molecules with Polar Bonds of equal
  dipole moment.....
• CO2 , BCl3, and CCl4
• PCl5 and SF6

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Polar Molecules

• Asymmetrical Molecules with Polar Bonds
• H2O and NH3
• HCl
• Symmetrical Molecules with Polar Bonds of unequal
  dipole moment
• e.g. CHCl3 and CF2Cl2
• Note CCl2F2 ? CFC-12 ? once used in
  refrigerators ?Ozone depletion

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Electronegativities of the Elements
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Figure 10.14
The orientation of polar molecules in an electric
field.
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