Chapter 9: Intermolecular Attractions and the Properties of Liquids and Solids

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Lecture Notes by Ken Marr Chapter 11
(Silberberg 3ed) Covalent Bonding Valence Bond
Theory and Molecular Orbital Theory 11.1
Valence Bond (VB) Theory and Orbital
Hybridization 11.2 The Mode of Orbital Overlap
and the Types of Covalent Bonds 11.3 Molecular
Orbital (MO)Theory and Electron
Delocalization
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Valence Bond Theory

• Covalent Bonds
• Result from the overlap of valence shell atomic
  orbitals to share an electron pair
• s, p, or hybrid orbitals may be used to form
  covalent bonds
• e.g. Predict the Orbitals used for bonding in
• H2, HF, H2S, F2

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Examples of s and p Orbitals involved in Bonding

• Overlap of s orbitals
• H2
• Overlap of s and p orbitals
• HF
• H2S
• Overlap of p orbitals
• F2

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Hybrid Orbitals

• The use of only s and p orbitals does not explain
  bonding in most molecules!!!
• e.g. BeCl2, CH4 , H2O
• hybrid orbitals are used in these cases
• Hybrid Orbitals are used to hold bonding and
  nonbonding electrons!
• s, p, and d orbitals may hybridize to form to
  form hybrid orbitals

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How to Determine an Atoms Hybridization

• Write Lewis structure for the molecule or ion,
  then...
• Determine number of electron pairs around the
  atom in question
• One orbital is needed for each electron pair
• sp hybridization provides 2 orbitals
• sp2 hybridization provides 3 orbitals
• sp3 hybridization provides.....?...........orbita
  ls
• sp3d hybridization provides......?..........orbit
  als
• sp3d2 hybridization provides....?.............orb
  itals

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Examples of Hybrid Orbitals

• Example of sp hybrid orbitals
• BeCl2

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Examples of Hybrid Orbitals

• Example of sp2 hybrid orbitals
• BF3

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Hybrid Orbitals

• Examples of sp3 hybrid orbitals
• CH4, C2H6, H2O, NH3
• Example of sp3d hybrid orbitals
• PCl5
• Example of sp3d2 hybrid orbitals
• SF6

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Bond Angle 92 o
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Mode of Orbital OverlapSigma vs. Pi Bonds

• Sigma Bonds (s-bond)
• Head to head overlap of s, p, or hybrid orbitals
• Responsible for the framework of a molecule
• Single bond one s bond

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Mode of Orbital Overlap Sigma vs. Pi Bonds

• Pi bonds (p-Bonds)
• Side to side overlap of p orbitals
• Restrict rotation
• Double bond one s bond one p bond
• Triple Bond one s bond two p bonds

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Examples Sigma vs. Pi Bonds

• Ethane
• Ethylene (ethene)
• Effect of p-bonding on rotation about the s-bond?
• Acetylene (ethyne)
• Nitrogen
• Formaldehyde

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Predict the hybrid orbitals used in the following

• Nitrogen gas, N2
• Formaldehyde H2CO
• Carbon dioxide, CO2
• Carbon monoxide, CO
• Sulfur dioxide, SO2

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One option for SO2
o o

• S Ne 3s2 3px2 3py1 3pz1
• This structure is...
• Favored by formal charge
• Requires ?? hybridization
• Big Problems with this Structure..
• How many unhybridized p-orbitals are available
  for p bonding?
• How many p-orbitals are needed?

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Another Option for SO2

• S Ne 3s2 3px2 3py1 3pz1
• ??? Hybridization
• Bond order?
• Resonance?

o o
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Resonance Delocalization of electrons

• Shifting of p-bond electrons without breaking the
  s- bond
• Although not favored by formal charge, B.O. 1.5

o o
o o
Resonance
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Molecular Orbital Theory

• p-electron pair found in molecular orbital formed
  from the overlap of p-orbitals
• B.O. 1.5
• same as measured B.O.
• S O bond length is intermediate between S O
  and S O bond lengths

o o
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Strengths and Weaknesses of Valence Bond Theory

• VB Theory ? Molecules are groups of atoms
  connected by localized overlap of valence shell
  orbitals
• VB, VSEPR and hybrid orbital theories work well
  together to explain the shapes of molecules
• ButVB theory inadequately explains
• Magnetic property of molecules
• Spectral properties of molecules
• Electron delocalization
• Conductivity of metals

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Molecular Orbital Theory

• The electrons in a molecule are found in
  Molecular Orbitals of different energies and
  shapes
• Just as an atoms electrons are located in atomic
  orbitals of different energies and shapes
• MOs spread over the entire molecule
• Major drawbacks of MO Theory
• Based on Quantum theory
• Calculations are based on solving very complex
  wave equations? major approximations are needed!
• Difficult to visualize

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Advantages of MO Theory

• VB Theory incorrectly predicts that....
• O2 is diamagnetic with B.O. 2 or....
• O2 is paramagnetic with B.O. 1
• MO Theory correctly predicts that....
• O2 is paramagnetic with B.O. 2
• VB Theory requires resonance structures to
  explain bonding in certain molecules and ions
• MO Theory does not have this limitation

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Formation of Molecular Orbitals

• MOs form when atomic orbitals overlap
• Bonding MOs
• Result from constructive interference of
  overlapping electron waves
• Stabilize a molecule by concentrating electron
  density between nuclei
• MOs ? more stable than AOs ? delocalize
  electron charge over a larger volume

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Overlap of standing electron Waves
Constructive interference
Destructive interference
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(High e- density)
(Low e- density)
Fig. 11.13
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H2 is more stable than the separate atoms
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Antibonding MOs

• Antibonding MOs
• Result from destructive interference of
  overlapping electron waves
• Reduce electron density between nuclei
• Destabilize a molecule
• Higher in energy than bonding MOs of the same type

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Using MO Theory to Calculate Bond Order

• VB definition of Bond Order....
• Number of electron pairs shared between 2 nuclei
• MO Theory
• B.O. ½ (No. Bonding e- - No. Antibonding
  e-)
• Meaning of B.O.
• B.O. gt 0, then molecule more stable than separate
  atoms
• B.O. 0, then zero probability of bond formation
• The greater the B.O., the stronger the bond

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Why Do Some Molecules Exist and Others Do Not?

• Why do H2 and He21 exist , but He2 does not?
  Recall..Bonding results only if there is a net
  decrease in PE
• Molecules with equal numbers of Bonding and
  antibonding electrons are unstable...Why?......
• Antibonding MOs raise PE more than Bonding MOs
  lower PE

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Use MO theory to predict if the following can
form

• Hydride ions H2 1- and H2 2-
• Li2 , Li2 1 , Li2 2 , Li2 1-
• Be2 , Be2 1 , Be2 2 , Be2 1-

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• In He2, the antibonding electrons in s1s cancel
  the PE lowering of s1s

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Sigma vs Pi Molecular Orbitals

• s Molecular Orbitals form when.....
• s - atomic orbitals overlap
• p - atomic orbitals overlap head to head
• p Molecular Orbitals form when.....
• p - atomic orbitals overlap side to side
• Why are s-bonds more stable than p- bonds?

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No mixing of 2s and 2p orbitals
Mixing of 2s and 2p orbitals
AO MO AO MO
Energy Levels for O2, F2 Ne2
AO MO
AO MO Energy Levels for B2, C2 N2
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Explaining MO Energy Levels for Period 2 Elements

• O2, F2 and Ne2
• Paired electrons in 2p sublevel ? Repulsions ? 2s
  and 2p different in Energy
• No mixing occurs between 2s and 2p orbitals
• Raises energy of s2s and s2s MO
• Energy of s2p lt Energy p2p
• B2, C2 and N2
• Only unpaired electrons in 2p sublevel ? 2s and
  2p are very close in energy
• Mixing occurs between 2s and 2p orbitals
• Lowers energy of s2s and s2s MO
• Energy of s2p gt Energy p2p

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Bonding in Diatomic Molecules of Period 2

• Rules for filling of Molecular Orbitals
• Apply the Rules for the filling of Atomic
  Orbitals (Aufbau principle)
• Electrons 1st fill MOs of lowest energy
• Only 2 electrons with opposite spin per MO
• MOs of same energy (sublevel) half fill before
  electrons pair
• Predict the bond order for each of the following
  molecules involving period 2 elements
• Li2, Be2, B2, C2, N2, O2, F2, Ne2, NO

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High Z effective of F results in lower energy or
its AOs
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1. Why are oxygens AOs at lower a energy than
   nitrogens?
2. Bond order?
3. Para- or diamagnetic?

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Delocalized Molecular Orbitals

• MO Theory (unlike VB Theory) does not require
  resonance to explain the bonding in.....
• Carbonate ion, Nitrate ion, Formate ion, Acetate
  ion, Benzene, etc.
• MO Theory Electron pairs can be shared by 3 or
  more atoms .......Why?
• MOs can overlap 3 or more atoms
• Delocalized Bonds form when an electron pair is
  shared by 3 or more atoms
• Offers stability in the same way that resonance
  offers stability

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Bonding in Solids

• Why do metals conduct electricity and nonmetals
  do not?
• Band Theory to the rescue!!

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Band Theory

• Energy Bands form from the overlap of atomic
  orbitals of similar energy from all atoms in a
  solid
• Energy bands containing core (nonvalence)
  electrons are localized
• i.e. Do not extend far from each atom
• Energy Bands containing valence electrons are
  delocalized
• I.e. extent continuously throughout the solid
• Conduction band Valence bands that are either
  partially filled or empty

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Energy Bands (MO Orbitals) for Na
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Fig. 12.37
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Electrical Conductors

• Have a conduction band that is partially filled
  (e.g. Group IA Transition Metals) ....or....
• Have an empty conduction band that overlaps a
  filled valence band (i.e. Have a narrow band gap)
  
• e.g. Group IIA Metals

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Fig. 12.38
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Nonconductors (Insulators)

• All valence electrons are used to form covalent
  bonds
• Have a large band gap between the filled valence
  band and the empty conduction band
• Some examples
• Glass, diamonds, rubber, most plastics

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Semiconductors

• Have a small band gap between the filled valence
  band and the empty conduction band
• Thermal Energy can promote electrons from filled
  valence band to empty conduction band
• e.g. Silicon

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Doping of Semiconductors

• p-type semiconductors
• Doped with a Group IIIA element
• Have one less electron than Si
• Causes positive holes in semi conductor
• Electricity flows through these positive holes
• n-type semiconductors
• Doped with a Group VA element
• Have one more electron than Si
• Causes negative holes in semiconductor
• Electricity flows through these negative holes

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