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Molecular Structure and Bonding
Dr.Christoph
Phayao University June 2014
Part 1
What is a chemical bond ?
Ionic Bond
Normally between a metal and a non-metal:
They exchange electrons and become ions
(charged atoms) which attract...
Covalent Bond
Two non-metals share
(valence) electrons:
(Remark: Transition metals can form covalent bonds also !)
Polar Covalent Bond
Two non-metals share electrons unevenly because
of electronegativity difference.
Electrons are closer ...
Metallic Bond
Metal atoms share all
their valence electrons,
which freely move
between all atoms which
form a network.
The...
Bond Polarity
Polar Bonds
Uneven sharing of electrons due to
differences in Electronegativity
The “pull”
an atom has
for electrons
Electronegativity Trends
Common Electronegativites
Highest value,
set to 4
Polar Molecules
Electrons are not equally shared in a bond,
which can lead to a dipolmoment of the whole molecule
Polar Bonds and Geometry
Which bond type ?
(exception:
Transition
metals !)
Electron counting
Formal Charge
Split all bonds in the middle
=> “real” charge on atoms
(2) Octet Rule
Count all bonding electrons for
one a...
Special Cases
“Extended
octet”
Especially P and
S can use d-
orbitals to make
more than 3
resp. 2 bonds !
6 VE:
Especially...
Part 2: Valence Bond Theory
(VB)
“Valence Electrons are located in
bonds and lone pairs”
Sigma bonds
Pi Bonds
“Resonance”
Write the resonance formula for OZONE !
Does the molecule have a charge ?
Important exception: Carbon Monoxide !
Homework(2)
Draw Lewis structure(s) and find
formal charges (all atoms)
and hybridization (central atom) in:
1. NO3 (-)
2....
***** Break *****
Prentice Hall © 2003 Chapter 9
• Atomic orbitals can mix or hybridize in order to adopt an
appropriate geometry for bondin...
Figure 11.2 The sp hybrid orbitals in gaseous BeCl2.
atomic
orbitals
hybrid
orbitals
orbital box diagrams
Figure 11.2 The sp hybrid orbitals in gaseous BeCl2(continued).
orbital box diagrams with orbital contours
Figure 11.3 The sp2 hybrid orbitals in BF3.
sp2 and sp3
Hybrid Orbitals
Figure 11.4 The sp3 hybrid orbitals in CH4.
Figure 11.5 The sp3 hybrid orbitals in NH3.
Figure 11.5 continued The sp3 hybrid orbitals in H2O.
Including d-orbitals
3d orbitals can be filled as well
=> Al acts as Lewis acid
=> P and S have “hypervalence”
Figure 11.6 The sp3d hybrid orbitals in PCl5.
Figure 11.7 The sp3d2 hybrid orbitals in SF6.
Acid or Base ?
Compare AlCl3 and PCl3 ?
Which acts as acid and which as base – and why ?
Why is FeCl3 a strong Lewis acid ?
SOLUTION:
PROBLEM: Describe the types of bonds and orbitals in acetone, (CH3)2CO.
PLAN: Use the Lewis structures to ascert...
Tasks
• Draw the Lewis Structures and the Hybrid
Orbitals for Ethane, Ethene and Ethyne
(mark the hybrid orbitals)
• Which...
Chemical Reactivity
From the hybrid orbitals we can estimate if a
molecule acts as Lewis acid or base
(if there is an elec...
***** Break *****
VSEPR
VSEPR Theory
Clip:
http://www.youtube.com/watch?v=nxebQZUVvTg
MO Theory
The Central Themes of MO
Theory
A molecule is viewed on a quantum mechanical level as a collection of nuclei
surrounded by...
Amplitudes of wave
functions added
Figure 11.14
An analogy between light waves and atomic wave functions.
Amplitudes of
wa...
Prentice Hall © 2003 Chapter 9
Molecular Orbitals
• Molecular orbitals:
• each contain a maximum of two electrons
• have d...
Prentice Hall © 2003 Chapter 9
Molecular Orbitals
The Hydrogen Molecule
Prentice Hall © 2003 Chapter 9
Figure 11.15 The MO diagram for H2.
Energy
MO
of H2
*1s
1s
AO
of H
1s
AO
of H
1s
H2 bond ...
Prentice Hall © 2003 Chapter 9
Electron Configurations and Molecular
Properties
• Two types of magnetic behavior:
• parama...
Diatomic molecules
The energy level
is the lower, the
higher the EN of
the atom is !
Naming of MO’s: example O2 molecule
“g” = symmetric to C axis
“u” = anti-symmetric
Diatomic molecules
Consider the EN of each atom – the higher the EN,
the lower is the energy of the orbitals !
The highest...
Example CO
HOMO
LUMO
“lone
pair” on C
Chemical Reactivity
Important are the HOMO and LUMO (“frontier orbitals”)
http://www.meta-synthesis.com/webbook/12_lab/lab...
GroupOrbitals
Construction of Group Orbitals – example H2O
Interaction 1: in-phase H orbitals
Interaction 2: out-of-phase H orbitals
Indicate different MO types: (bonding, non-bonding. anti-bonding)
Combination of 3 H
orbitals to
3 group orbitals
BH3 molecule
Compare HOMO/LUMO to BH3 !
=> what is an acid / base ?
Homework (3)
Number Molecule
1 CN
2 CN(-)
3 BC
4 BN
5 BO
6 BF
7 CF
8 NO
9 NO (+)
10 NO (-)
Number Molecule
11 NF
12 OF
13 ...
Molecular structure and bonding
Molecular structure and bonding
Molecular structure and bonding
Molecular structure and bonding
Molecular structure and bonding
Molecular structure and bonding
Molecular structure and bonding
Molecular structure and bonding
Molecular structure and bonding
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Molecular structure and bonding

How can we describe the bonding between atoms forming molecules ? We start from valence bond theory and come to molecular orbital theory.

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Molecular structure and bonding

  1. 1. Molecular Structure and Bonding Dr.Christoph Phayao University June 2014
  2. 2. Part 1 What is a chemical bond ?
  3. 3. Ionic Bond Normally between a metal and a non-metal: They exchange electrons and become ions (charged atoms) which attract each other by electrostatic force. A pair of ions does not stay alone but form crystals
  4. 4. Covalent Bond Two non-metals share (valence) electrons: (Remark: Transition metals can form covalent bonds also !)
  5. 5. Polar Covalent Bond Two non-metals share electrons unevenly because of electronegativity difference. Electrons are closer to one atom than the other. This results on partially negative and positive charges on the atoms
  6. 6. Metallic Bond Metal atoms share all their valence electrons, which freely move between all atoms which form a network. Therefore all metals can conduct electricity and look shiny
  7. 7. Bond Polarity
  8. 8. Polar Bonds Uneven sharing of electrons due to differences in Electronegativity The “pull” an atom has for electrons
  9. 9. Electronegativity Trends
  10. 10. Common Electronegativites Highest value, set to 4
  11. 11. Polar Molecules Electrons are not equally shared in a bond, which can lead to a dipolmoment of the whole molecule
  12. 12. Polar Bonds and Geometry
  13. 13. Which bond type ? (exception: Transition metals !)
  14. 14. Electron counting
  15. 15. Formal Charge Split all bonds in the middle => “real” charge on atoms (2) Octet Rule Count all bonding electrons for one atom => 8 is most stable (3) Oxidation Number Give all bonding electrons to the more electronegative atom
  16. 16. Special Cases “Extended octet” Especially P and S can use d- orbitals to make more than 3 resp. 2 bonds ! 6 VE: Especially common for B and Al !
  17. 17. Part 2: Valence Bond Theory (VB) “Valence Electrons are located in bonds and lone pairs”
  18. 18. Sigma bonds
  19. 19. Pi Bonds
  20. 20. “Resonance” Write the resonance formula for OZONE ! Does the molecule have a charge ?
  21. 21. Important exception: Carbon Monoxide !
  22. 22. Homework(2) Draw Lewis structure(s) and find formal charges (all atoms) and hybridization (central atom) in: 1. NO3 (-) 2. PO4 (3-) 3. CH3 Cl 4. CH2Cl2 5. SO2 6. SO3 7. CO3 (2-) 8. H2O2 9. N2O 10. Cl O2
  23. 23. ***** Break *****
  24. 24. Prentice Hall © 2003 Chapter 9 • Atomic orbitals can mix or hybridize in order to adopt an appropriate geometry for bonding. • Hybridization is determined by the electron domain geometry. sp Hybrid Orbitals • Consider the BeF2 molecule (experimentally known to exist): Hybrid Orbitals
  25. 25. Figure 11.2 The sp hybrid orbitals in gaseous BeCl2. atomic orbitals hybrid orbitals orbital box diagrams
  26. 26. Figure 11.2 The sp hybrid orbitals in gaseous BeCl2(continued). orbital box diagrams with orbital contours
  27. 27. Figure 11.3 The sp2 hybrid orbitals in BF3.
  28. 28. sp2 and sp3 Hybrid Orbitals
  29. 29. Figure 11.4 The sp3 hybrid orbitals in CH4.
  30. 30. Figure 11.5 The sp3 hybrid orbitals in NH3.
  31. 31. Figure 11.5 continued The sp3 hybrid orbitals in H2O.
  32. 32. Including d-orbitals 3d orbitals can be filled as well => Al acts as Lewis acid => P and S have “hypervalence”
  33. 33. Figure 11.6 The sp3d hybrid orbitals in PCl5.
  34. 34. Figure 11.7 The sp3d2 hybrid orbitals in SF6.
  35. 35. Acid or Base ? Compare AlCl3 and PCl3 ? Which acts as acid and which as base – and why ? Why is FeCl3 a strong Lewis acid ?
  36. 36. SOLUTION: PROBLEM: Describe the types of bonds and orbitals in acetone, (CH3)2CO. PLAN: Use the Lewis structures to ascertain the arrangement of groups and shape at each central atom. Postulate the hybrid orbitals taking note of the multiple bonds and their orbital overlaps. H3C C CH3 O sp3 hybridized sp3 hybridized C C C O H H HHH H sp2 hybridized bonds bond C C C O sp3 sp3 sp3 sp3 sp3 sp3 sp3 sp3 sp2 sp2 sp2 sp2 sp2sp2 H H H H H H
  37. 37. Tasks • Draw the Lewis Structures and the Hybrid Orbitals for Ethane, Ethene and Ethyne (mark the hybrid orbitals) • Which hybridization has the central atom in: H2O, O2, NH3, NH4+, N in pyridine, O in THF, S in SOCl2, C in HCHO compared to CO
  38. 38. Chemical Reactivity From the hybrid orbitals we can estimate if a molecule acts as Lewis acid or base (if there is an electrophilic or nucleophilic center) Consider the “empty” pz orbital of C in HCHO vs. the “filled” sp orbital of C in CO -> in the first case, it acts as Lewis acid, in the second as base !
  39. 39. ***** Break *****
  40. 40. VSEPR
  41. 41. VSEPR Theory Clip: http://www.youtube.com/watch?v=nxebQZUVvTg
  42. 42. MO Theory
  43. 43. The Central Themes of MO Theory A molecule is viewed on a quantum mechanical level as a collection of nuclei surrounded by delocalized molecular orbitals. Atomic wave functions are summed to obtain molecular wave functions. If wave functions reinforce each other, a bonding MO is formed (region of high electron density exists between the nuclei). If wave functions cancel each other, an antibonding MO is formed (a node of zero electron density occurs between the nuclei).
  44. 44. Amplitudes of wave functions added Figure 11.14 An analogy between light waves and atomic wave functions. Amplitudes of wave functions subtracted.
  45. 45. Prentice Hall © 2003 Chapter 9 Molecular Orbitals • Molecular orbitals: • each contain a maximum of two electrons • have definite energies • can be visualized with contour diagrams • are distributed over the whole molecule (not only in between 2 atoms) • When two AOs overlap, two MOs form.
  46. 46. Prentice Hall © 2003 Chapter 9 Molecular Orbitals The Hydrogen Molecule
  47. 47. Prentice Hall © 2003 Chapter 9 Figure 11.15 The MO diagram for H2. Energy MO of H2 *1s 1s AO of H 1s AO of H 1s H2 bond order = 1/2(2-0) = 1 Filling molecular orbitals with electrons follows the same concept as filling atomic orbitals.
  48. 48. Prentice Hall © 2003 Chapter 9 Electron Configurations and Molecular Properties • Two types of magnetic behavior: • paramagnetism (unpaired electrons in molecule): strong attraction between magnetic field and molecule; • diamagnetism (no unpaired electrons in molecule): weak repulsion between magnetic field and molecule. • Magnetic behavior is detected by determining the mass of a sample in the presence and absence of magnetic field:
  49. 49. Diatomic molecules
  50. 50. The energy level is the lower, the higher the EN of the atom is !
  51. 51. Naming of MO’s: example O2 molecule “g” = symmetric to C axis “u” = anti-symmetric
  52. 52. Diatomic molecules Consider the EN of each atom – the higher the EN, the lower is the energy of the orbitals ! The highest filled MO is called “HOMO”, the lowest unoccupied MO “LUMO” -> check example CO http://firstyear.chem.usyd.edu.au/calculators/ mo_diagrams.shtml
  53. 53. Example CO HOMO LUMO “lone pair” on C
  54. 54. Chemical Reactivity Important are the HOMO and LUMO (“frontier orbitals”) http://www.meta-synthesis.com/webbook/12_lab/lab.html
  55. 55. GroupOrbitals
  56. 56. Construction of Group Orbitals – example H2O
  57. 57. Interaction 1: in-phase H orbitals
  58. 58. Interaction 2: out-of-phase H orbitals
  59. 59. Indicate different MO types: (bonding, non-bonding. anti-bonding)
  60. 60. Combination of 3 H orbitals to 3 group orbitals BH3 molecule
  61. 61. Compare HOMO/LUMO to BH3 ! => what is an acid / base ?
  62. 62. Homework (3) Number Molecule 1 CN 2 CN(-) 3 BC 4 BN 5 BO 6 BF 7 CF 8 NO 9 NO (+) 10 NO (-) Number Molecule 11 NF 12 OF 13 CH4 14 BH3 15 SbF6 16 XeF2 17 XeF4 18 XeF6 http://firstyear.chem.usyd.edu.au/calculators/mo_diagrams.shtml

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