Molecular Geometry and Bonding Theories

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Molecular Geometry and Bonding Theories

• Chapter 9
• HW page 388

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Introduction

• Properties of substances are determined by
• Strength of bonds
• Polarity of bonds
• Size of molecules
• Shape of molecule
• Example Article building better aspirin

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Molecular Shapes

• The over shape of a molecule is determined by its
  bond angles
• We will begin with compounds that have a central
  atom (A) that is bonded to n-B atoms ABb
• Shape of ABb depends on value of n

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n 2
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n 3
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Can We Predict the Shape?

• There are three different theories of molecular
  bonding.
• VSEPR
• Valence Bond Theory
• Molecular Orbital Theory
• Why 3 Theories?
• One model does not describe all the properties of
  the molecular bonds

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What is the advantage of each theory?

• VSEPR
• Easy, used for shape of molecule
• Valence Bond Theory
• Good for hybridization
• Molecular Orbital Theory
• Does a good job of predicting electronic spectra
  and parmagnetism (the other 2 do not)
• Limited to talking about diatomic molecules

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The VSEPR Model
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Can We Predict the Shape?

• Using the VSEPR Model
• V valence
• S shell
• E electron
• P pair
• R - repulsion

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VSEPR Model

• Electron Domains
• Bonding pairs vs. lone pairs
• NH3
• Because the electron domains are negatively
  charged regions.

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Balloon analogy
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There are 5 basic Shapes of Electron Domains

• Linear (2)
• Trigonal planar (3)
• Tetrahedrally (4)
• Trigonal bipyramidal (5)
• Octahedral (6)
• Movie on domains

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• Electron domain geometry
• Best arrangement minimizes the e- repulsion
• Electron domain geometry is used to predict
  molecular geometry
• Molecular geometry

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G
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Bent Geometry
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Trigonal Pyramidal
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Tetrahedral
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Steps to predict molecular geometries

• Sketch Lewis structure of molecule or ion
• Count total number of e- domains around central
  atom
• Double/triple bond one domains
• Then use the chart
• Lets try some

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• SnCl3-
• BF3
• SeCl2
• CO32-
• PBr3
• CH3
• SO3

• TeF4
• AsCl5
• BrF5
• H3O
• KrF2
• SF6
• ClO3-

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HW 22
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Effect of nonbonding electrons and multiple bonds
on bond angle
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• Nonbonding e- pairs
• Notice that the angle decreases and the number of
  nonbonding pair of electrons increases
• Why?

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Multiple Bonds

• Contain higher electronic charge density than
  single bonds
• So larger electron domains

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Molecules with expanded Valence Shells

• So far we have looked at molecules with nor more
  than an octet or electrons around the central
  atom
• But when the central atom is from the 3rd period
  and beyond that atom may have more than 4 pairs
  of electrons around it
• When viewing these electron domains there are two
  choices of regions

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• Axial Position of electron domain
• Equatorial position of electron domain
• Note when nonbonding pairs are involved they
  occupy the equatorial position because here they
  lesson the repulsion

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Lets Try some

• SF4
• IF5
• ClF3

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Valance-Bond Theory
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VI. Covalent Bond and orbital overlap

• How can we explain bonding and account for
  geometries of molecules in terms of atomic
  orbitals?
• Valance-Bond Theory
• Orbitals will overlap thus share a region

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Hybrid Orbitals
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Hybrid Orbitals

• A blending of orbitals used to justify certain
  bonds
• 2s2 2p1 2sp2
• 2s2 2p2 2sp3
• 3s2 3p3 3d0 3sp3d
• Movie on hybrid orbitals

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HYBRIDIZATION OF CARBON
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HW 48
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Multiple Bonds
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VII. Multiple Bonds

• Sigma bonds (s) line joining 2 nuclei passes
  through the middle of the overlap region
• Example 2s in hydrogen
• Example s and p in hydrochloric acid

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• pi bonds (p) overlap region lie above and below
  the nuclear axis
• Double bond is both
• 1 sigma s
• 1 pi bond p
• Better orbital overlap will be do to
• Same size atom
• Same energy levels
• Sigma bonds are stronger than pi bonds due to
  overlap

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Molecular orbital (MO) Theory
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Molecular Orbitals

• Valence bond theory and VSPER do not explain
• How molecules absorb light, giving them color
• Also no information about bond energy
• The Molecular orbital (MO) does
• Molecular orbital describes electrons in
  molecules as waves
• Movie

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• Have many characteristics of atomic orbitals
• MO can hold max of 2 electrons with opposite spin
• MO have definite energy (unlike atomic orbitals)
• MO are associated with entire molecule

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MO theory has 5 rules

• 1. The number of MO orbitals number of atomic
  orbitals
• 2. There are 2 MOs
• One is a bonding orbital, of lower energy
• One is an antibond orbital, of high energy
• Node region of zero electron density
• 3. Electrons enter the lowest orbital available
• 4. Max of 2 electrons per orbital (Paulis
  exclusion rule)
• 5. Electrons have opposite spin, fill each
  orbital of same energy before double up (Hunds
  Rule)

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• Energy level diagram for hydrogen
• Electron configuration
• Bond Order (BO)
• BO ½ (bonding e-s antibonding e-s)
• If a molecule is to be stable BO must be greater
  than zero
• If zero, than no bond forms
• BO indicates strength
• BO indicates bond length
• BO indicates stability of the molecules
• Lets solve the BO for hydrogen

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• Lets try some
• Homonuclear diatomic helium
• H2-
• Lets look at molecules with p orbitals
• When bonded only one p orbital will bond end to
  end, or head to head. This is a sigma bond, s
• s2pz or s2pz
• When bonded the other 2 p orbitals will bond by
  overlapping sideways.
• So these are pi, p, bonds
• p 2px or p 2px or p 2py or p 2py

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• Because s overlap is greater, s2p is lower
  energy, or more stable than p2p (which will be
  more energy, less stable)
• Lets try nitrogen
• Energy level diagram
• Bond order
• Electron configuration
• Lets try some others

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oxygen
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Nitrogen
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HW 62
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ParamagnetismDiamagnetism
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X. Electron configuration and molecular
properties

• Most materials have no magnetism until they are
  placed in a magnetic field
• However in the presence of such a field, 2 types
  of magnetism can be induced
• Paramagnetism
• Diamagnetism

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Paramagnetism

• Attraction to inducing magnets
• Molecules with one or more unpaired electrons
• To determine in a lab, sample is weighed in a
  magnetic field, if sample is drawn into the field
  thus appears to gain mass, then paramagnetic
• Example oxygen
• Explains why oxygen is soluble is polar water

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Dimagnetic
Paramagnetic
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Diamagnetism

• No unpaired electrons
• Substances repelled by inducing magnetic fields,
  so less weight in a magnetic field
• Much weaker than paramagnetism
• Example - hydrogen

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HW 68
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