1. Formation of Covalent Bonds
Localized Electron (LE) Model-electron pairs are still localized around specific atoms, but orbitals around central atom are modified
Formation of Covalent Bonds
Two different theories which attempt to explain covalent molecular/ionic structure/shape
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Molecular Orbital (MO) Model-all electrons in molecule are combined into set of molecular orbitals which describe bonding in entire molecule

2. Why hybrid orbitals?
VSEPR model does good job at predicting molecular shape, despite fact that it has no obvious relationship to filling and shapes of atomic orbitals
Based on shapes and orientations of 2s/2p orbitals on carbon atom, not obvious why CH4 molecule should have tetrahedral geometry
Used to reconcile covalent bonds formed from overlap of atomic orbitals with molecular geometries from VSEPR model
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3. Hybridization/Hybrid orbitals
Hybridization (orbitals in covalently bonded atoms)
Mixing of 2/more atomic orbitals of similar energies on same atom to produce new orbitals of equal energies
Hybrid orbitals
Valence bond theory creates hybrid orbitals that are linear combinations of s/p orbitals in valence shell (d if necessary)
# atomic orbits = # hybrid orbitals
Each hybrid orbital equivalent to others but large lobes point in different directions
Atomic orbitals not used to make hybrids unaffected
Mixtures of atomic orbitals with intermediate energy
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4. One 2s electron promoted to empty (2p) orbital
No unpaired electrons
One 2s electron promoted to empty (2p) orbital
2 occupied orbitals blend to form 2 sp hybrid orbitals/2 remaining p orbitals
unchanged
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5. 2 sp hybrids point in opposite directions at 180o to each other
sp has 50% s/50% p character
2 sp hybrids point in opposite directions at 180o to each other
Require energy to promote 2s  2p orbital
Large lobe of hybrid orbital can be directed at other atoms better than unhybridized atomic orbital
Overlap more strongly/stronger bonds result
Energy released by bond formation offsets energy expended to promote electrons
Each sp hybrid involved in s bond/remaining p orbitals forms 2p bonds (All contain single unpaired electron)
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6. +
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7. 2s electron promoted to 2p orbital
2s electron promoted to 2p orbital
Have 3 unpaired electrons that form 3 sp2 orbitals
Creates 3 identical orbitals of intermediate energy/length
Leaves one unhybridized p orbital
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8. Large lobes of orbitals lie in plane at angles of 120o and point toward corners of triangle+
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9. Four sp3 orbitals identical in shape
Promotion of 2s electron to 2p orbital results in valence shell with 4 unpaired electrons in four sp3 hybrid orbitals.
Four sp3 orbitals identical in shape
tetrahedral
4 orbitals form one 2s/three 2p orbitals (s1p3)
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10.
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11. D-orbital hybridization
Central atoms located in Period 3 and above can use empty d orbitals to receive promoted s electron
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12. dsp3 Hybridization 5 effective pairs around central atom
Trigonal bipyramidal shape
Lobes have bond angles of 90o & 120O
PCl5 example
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13. Each dsp3 orbital also has small lobe not shown in diagram
Phosphorus uses set of 5 dsp3 orbitals to share electron pairs with sp3 orbitals on 5 chlorine atoms
Other sp3 orbitals on each chlorine atom hold lone pairs
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14. d2sp3 Hybridization Six effective pairs around central atom
Octahedral structure
Lobes have angles of 90o
SF6 example
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15. Atomic Hybrid
orbital set orbital set_
s/p 2 sp
s/p/p 3 sp2
s/p/p/p 4 sp3
s/p/p/p/d 5 sp3d
s/p/p/p/d/d 6 sp3d2
Examples
BeF2, HgCl2
BF3, SO3
CH4, NH3, H O, NH4+
PF5, SF4, BrF3, SbCl52-
SF6, ClF5, XeF4, PF4-
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16. Trigonal planar Bent sp2 2 1 Tetrahedron sp3 4 0
Basic Derived Hybrid Bonding Nonbonding
Structure structure e- pairs e- pairs
Linear sp 2 0
Trigonal planar sp2 3 0
Trigonal planar Bent sp2 2 1
Tetrahedron sp3 4 0
Tetrahedron Triangular pyramid sp3 3 1
Tetrahedron Bent sp3 2 2
Trigonal bipyramid sp3d 5 0
Trigonal bipyramid Distorted tetrahedron sp3d 4 1
Trigonal bipyramid T-shape sp3d 3 2
Trigonal bipyramid Linear sp3d 2 3
Octahedron sp3d2 6 0
Octahedron Square pyramid sp3d2 5 1
Octahedron Square planar sp3d2 4 2
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17. Strength of sigma bonds p-p > p-s > s-s
(s/single /δ)
(Head-to-head overlap)
Lobes of bonding orbital point toward each other.
Overlap of two S orbitals to form sigma bond (green)
Overlap of two P orbitals to form sigma bond (green)
Strength of sigma bonds
p-p > p-s > s-s
p-orbitals allow overlap to greater extent as compared to p-s which is larger as compared to s-s overlap
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18. Maximum electron density lies along bond (electron pair shared in area centered on line connecting nuclei)
Line joining 2 nuclei passes through middle of overlap region (between nuclei)
Maximum overlap forms strongest-possible sigma bond
Atoms arrange themselves to give greatest-possible orbital overlap
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19. (p/)
(perpendicular)
Axes parallel to each other but perpendicular to internuclear axis
Occupies space above/below internuclear axis (imaginary line connecting nuclei of two atoms)
Electron density zero along bond
Atomic orbitals interact above and below nuclei
Formed only in addition to sigma bond
Always present in molecules with double/triple bonds
Occur only w/sp or sp2 hybridization present on central atom, but not sp3
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20. Significantly less overlap between component p-orbitals due to parallel orientation
Weaker than sigma bonds-electrons farther from nucleus, so more reactive
Pi bonds are superimposed on sigma bonds so they simply modify dimensions of molecule.
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21. Multiple Bonds Double Bonds Triple Bonds
Consist of 1 s bond (overlap of 2 sp orbitals) and 1 p bond (overlap of 2 p orbitals)
s bond where electron pair located directly between atoms
p bond where shared pair occupies space above and below s bond
Consist of 1 s / 2 p bonds
Side-to-side overlap makes p bond electrons more reactive
Electron density no longer located on internuclear axis (one electron cloud above and one below)
Bond is weaker-p orbitals do not overlap as much in p bond as s bond
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22. Consists of sigma (green)/2 pi bonds (red)
Double bonds
Triple bonds
Consists of sigma (green)/2 pi bonds (red)
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23. sp hybridized nitrogen atom s bond in N2 molecule
2 p bonds in N2 are formed when electron pairs are shared between two sets of parallel p orbitals
Total bonding picture for N2
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24. Carbon is unique sp3 hybrid orbitals = single bonds
sp2 hybrid orbitals = double bonds
sp hybrid orbitals = triple bonds
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25. Carbon atom of methane (CH4)
Made up of 4 C-H sigma (σ) bonds
Each hybrid sp3 orbitals of carbon undergoes end-on overlap with s-orbitals of H atoms
Tetrahedral geometry
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26. Carbon atoms of ethyne (acetylene - C2H2)
2 C-H σ bonds, 1 C-C σ bond, 2 C-C π bonds
Δ bonds (gray) are linear in arrangement
Unhybridized p-orbitals (green/purple) interact with each other laterally, resulting in bond formation
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27. Carbon atoms of ethene (ethylene - C2H4)
4 C-H σ bonds, 1 C-C σ bond, 1 C-C π bond
Hybrid orbital overlap end-on with s-orbitals of H atoms (δ bond in gray)
Unhybridized p-orbitals (purple) at right angles to plane of hybrids, overlap laterally (π bond-double)
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28. Formaldehyde has following Lewis structure:
Describe it bonding in terms of appropriate hybridized/unhybridized orbitals
VSEPR predicts trigonal planar geometry which suggests sp2 hybrid orbitals on C
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29. Acetonitrile molecule
Predict bond angles around each C
Approximately 109° around left C and 180° on right C
Give hybridization of each C
sp3, sp
Determine total number of δ/π bonds
5 δ bonds and 2 π bonds
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30. hybrid orbitals in the CO2 molecule
orbitals of sp hybridized carbon atom
orbital arrangement for sp2 hybridized oxygen atom
(a) orbitals in carbon dioxide-carbon-oxygen double bonds each consist of one s bond and one p bond. (b) Lewis structure for carbon dioxide
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31. Determine the total number of sigma and pi bonds in each of the following:
Using the simple Lewis structure, also determine the hybridization for each:
CH3Cl
4 δ, 0 Π, sp3
PH3
2 δ, 2 Π, sp
H2S
3 δ, 0 Π, sp3
CO32-
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32. SO32- 2 δ, 0 Π, sp3 CS2 3 δ, 1 Π, sp2 SiF4 NO3- 4 δ, 0 Π, sp3 PO43-
ClO4-
4 δ, 0 Π, sp3
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33. VSEPR did not explain why bonds exist between atoms
Pair of electrons attracted to both atomic nuclei
Bond is formed
As extent of overlap increases, strength of bond increases
Electronic energy drops as atoms approach each other
Begin to increase again when they become too close
Optimum distance (observed bond distance) at which total energy is at minimum
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34. Basic s, px, py, and pz orbitals unsatisfactory for two reasons
Orbitals not directed in particular direction-tend to spread out in all directions
Geometry of orbitals rarely consistent with molecular geometry
Could not adequately explain fact that some molecules contain two equivalent bonds with bond order between that of single and double bonds
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35. Atomic orbitals explain bonding/ account for molecular geometries
Mathematical descriptions of where electrons most likely found
Obtained by solving Schrödinger equation
As angular momentum (ml) and energy of electron increases, it tends to reside in differently shaped orbitals
Orbitals corresponding to three lowest energy states (s, p, and d, respectively)
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36. Localized Electron Model (Valence Bond Theory)
Describes structure of covalent bonds (how bonding occurs)
Atoms in molecule bond together w/shared electrons
Lewis structure shows valence electron arrangement
Use VSEPR model to predict molecular geometry
Atoms use atomic orbitals to share electrons/hold lone pairs (new set of hybridized orbitals can form)
Lone pairs: electron pairs localized on atom
Bonding pairs: electron pairs in space between atoms
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37. Combine Lewis’s notion of electron-pair bonds with atomic orbitals
Lewis theory: covalent bonding occurs when atoms share electrons (concentrates electron density between nuclei)
Valence-bond theory: buildup of electron density between two nuclei occurs when valence atomic orbital of one atom shares space (overlaps) with that of another atom
Shortcomings of LE Model
Electrons not actually localized
Does not deal effectively w/molecules containing unpaired electrons
Gives no direct information about bond energies
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38. Direct overlap (electron sharing) of two atomic orbitals
Local view of bonding (how adjacent atoms share electrons)
How 2 electrons of opposite spin share space between nuclei/form bond
Paired electrons localized in specific internuclear spaces between bonded atoms or remain unshared (lone pairs)
As bond is formed, paired electrons spread out over molecule to form final electron cloud surrounding nuclei
Description confirmed by many chemistry/physics experiments, including actual Scanning Tunneling Microscope picture of p-orbital
Electronic structure/geometry is best compromise between maximum overlap (electron-nucleus attraction) and repulsion (electron-electron/nucleus-nucleus)
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39. δ bonds have cylindrical symmetry
Formed between pair of atoms within molecule
Increased electron density on internuclear axis
Rotation around bond does not change overlap of contributing atomic orbitals
Lower energy than π bonds
Bonds (δ-head on/π-sideways) made by overlap of atomic (s,p) or hybridized (sp2) orbitals
Formation of bonding orbital accompanied by formation of antibonding orbitals (δ*/π*) which remain unoccupied and does not contribute to structure of molecule
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40. π bonds are not cylindrically symmetric
May cover more than 2 nuclei (resonance)
Increased electron density above/below internuclear axis (not on axis itself)
Rotation breaks bond
Electrons not always shared equally (EN)
Skeletal structures (Lewis structures) correspond to valence-bond model
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41. Rules for drawing reasonable resonance structures
Resonance structures represent molecules not adequately described by single structure, because electrons shared by more than 2 nuclei
Rules for drawing reasonable resonance structures
All resonance structures must be valid Lewis structures
In all possible resonance structures atomic nuclei must not change their positions
All atoms must not change their hybridization
Only electron distribution may be changed
All resonance structures must have same # unpaired electrons
All atoms involved in resonance (electron sharing)/atoms directly bonded to them must lie in (or nearly in) same plane
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42. Homework:
Read 9.1, pp
Q pp , #12, 14, 16, 21, 22, 27, 28
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43. Molecular Orbital (MO) Theory (Model)
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44. Molecular Orbital Theory
Explains distributions (organization of valence electrons) and energy of electrons in molecules
Molecule is similar to atom (have distinct energy levels that electrons can populate)
Takes global view of bonding (all electrons in molecule are needed to describe how bonding occurs)
Useful for describing properties of compounds
Bond energies, electron cloud distribution, and magnetic properties
Solution to VB problems by creating new set of orbitals w/intermediate between those of basic orbitals used to construct them
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45. Basic principles of MO Theory
2 atomic orbitals w/similar energies overlap to form 2 molecular orbitals
Electrons from each element participating in bond occupy new molecular orbitals
MOs delocalized over many atoms (don’t directly correspond to specific bond as in VB theory)
No hybridization (all available atomic orbitals mixed into multiple combination
Molecular orbitals have different energies depending on type of overlap
Bonding orbitals (lower energy than corresponding AO)
Nonbonding orbitals (same energy as corresponding AO)
Antibonding orbitals (higher energy than corresponding AO)
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46. Molecular orbitals (MOs) made of fractions of atomic orbitals
All atoms in molecule provide atomic orbitals to make MOs, but not all atomic orbitals must participate in all MOs
MO filled by all available electrons (2 per orbital), starting with lowest energy MO orbital
π bonds perpendicular to δ bonds, so can’t mix
Reflects orbital geometry of molecule
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47. Principles for formation of MOs
# MOs formed = # atomic orbitals combined
Atomic orbitals combine best with other atomic orbitals of similar energy
How effectively 2 atomic orbitals combine proportional to their overlap
As overlap increases, energy of bonding MO lowered, energy of antibonding MO raised
Lower energy molecular orbitals fill first
Each MO accommodates 2 electrons w/opposite spins (Pauli exclusion principle)
When MOs of same energy populated, Hund’s rule is followed (equal energy orbitals ½ filled before pairing up)
Electron in antibonding orbital “cancels” its corresponding electron in bonding orbital
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48. In atoms, electrons occupy atomic orbitals,
In molecules they occupy molecular orbitals which surround molecule
Two atomic orbitals combine to form two molecular orbitals, one bonding (s) and one antibonding (s*)
Each line in diagram represents orbital
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49. Bonding Orbital
MOs w/lower energy than its corresponding original atomic orbitals
Promotes formation of stable bond
High electron density along internuclear axis
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50. Antibonding Orbital
MOs w/higher energy than its corresponding atomic orbitals
Destabilizes (negative impact) formation of bond
Much lower electron density along internuclear axis
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51. Antibonding molecular orbital is designated as π* molecular orbital
When two atomic orbitals overlap in side-by-side fashion, molecular orbitals called π molecular orbitals
Antibonding molecular orbital is designated as π* molecular orbital
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52. Bond Order (BO) Describes nature of bond formed by molecular orbitals
Refers to average number of bonds that atom makes in all of its bonds to other atoms
Bond order above 0 considered stable because it has excess of bonding electrons
If bond order = 0 (BE = ABE), species does not exist
Larger bond order = Greater bond strength = Greater bond energy = Shorter bond length
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53. Pg , Sample 9.6
For the species O2, O2+, and O2-, give the electron configuration and the bond order for each. Which has the strongest bond?
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54. Paramagnetism: (oxygen-2 unpaired e’s)
Magnetism can be induced in some nonmagnetic materials when in presence of magnetic field
Paramagnetism: (oxygen-2 unpaired e’s)
Unpaired electrons
Attracted to induced magnetic field
Much stronger than diamagnetism
Diamagnetism
Paired electrons
Repelled from induced magnetic field
Much weaker than paramagnetism
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55. Calculate bond order of NO+ ion ½(10 - 4) = 3
NO+ ion has total of = 14 electrons to place in molecular orbitals as follows
Calculate bond order of NO+ ion
½(10 - 4) = 3 
Is the NO+ ion diamagnetic or paramagnetic?  
0 unpaired electrons so ion is diamagnetic 
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56. Hydrogen Atom (H) 1 bonding electron/0 antibonding electrons
Stable (lower energy, greater stability)
Bond order of ½
One unpaired electron-paramagnetic
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57. Hydrogen Molecule (H2)
2 bonding electrons in δ1s molecular orbital /0 antibonding electrons
Stable
Bond order of 1
Electrons in line w/2 nuclei, so are s molecular orbitals
No unpaired electrons-diamagnetic
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58. Does He2 exist?
He2 has four electrons, two in s1s orbital and two in s*1s orbital
Bond order of 0, so He2 does not exist
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59. Bonding in Homonuclear Diatomic Molecules (composed of two identical atoms)
In order to participate in molecular orbitals, atomic orbitals must overlap in space
Larger bond order is favored
When molecular orbitals are formed from p orbitals, s orbitals are favored over p orbitals (stronger)
Electrons are closer to nucleus = lower energy
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60. Molecular orbital energy-level diagrams, bond orders, bond energies, and bond lengths for diatomic molecules
Note that for O2 and F2, 2p orbital lower in energy than the π2p orbitals
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61. Bonding in Heteronuclear Diatomic Molecules (different atoms)
For atoms adjacent to each other in periodic table
Use molecular orbital diagrams for homonuclear molecules
Significantly different atoms
Each molecule must be examined individually
Use only electrons that are going to be involved in bonding
No universally accepted molecular orbital energy order
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62. N2 NO bond order = 3 bond order = 2.5
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63. Outcomes of MO Model Strengths Disadvantages
Correctly predicts relative bond strength and magnetism of simple diatomic molecules
Accounts for bond polarity
Correctly portrays electrons as being delocalized in polyatomic molecules
Disadvantages
Difficult to apply quantitatively to polyatomic molecules
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64. Combining Localized Electron and Molecular Orbital Model
Resonance
Attempt to draw localized electrons in structure in which electrons not localized
s (δ) bonds can be described using localized electron model
p (π) bonds (delocalized) must be described using molecular orbital model
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65. Benzene s bonds (C - H and C - C) are sp2 hybridized
Localized model
p bonds result of remaining p orbitals above/ below plane of benzene ring
Delocalizing gives stability
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66. Homework: Read 9.2-9.5, pp. 426-441 Q pp. 443-444, #32, 33, 38, 40, 46
Do 1 additional exercise and 1 challenge problem
Submit quizzes by to me:
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