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**Orbitals and Covalent Bond**

Chapter 9 Orbitals and Covalent Bond

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Molecular Orbitals The overlap of atomic orbitals from separate atoms makes molecular orbitals Each molecular orbital has room for two electrons Two types of MO Sigma ( ) between atoms Pi ( ) above and below atoms

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**Sigma bonding orbitals**

From s orbitals on separate atoms + + + + + + Sigma bonding molecular orbital s orbital s orbital

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**Sigma bonding orbitals**

From p orbitals on separate atoms p orbital p orbital Sigma bonding molecular orbital

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**Pi bonding molecular orbital**

Pi bonding orbitals p orbitals on separate atoms Pi bonding molecular orbital

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**Sigma and pi bonds All single bonds are sigma bonds**

A double bond is one sigma and one pi bond A triple bond is one sigma and two pi bonds.

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**Atomic Orbitals Don’t Work**

to explain molecular geometry. In methane, CH4 , the shape is tetrahedral. The valence electrons of carbon should be two in s, and two in p. the p orbitals would have to be at right angles. The atomic orbitals change when making a molecule

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Hybridization We blend the s and p orbitals of the valence electrons and end up with the tetrahedral geometry. We combine one s orbital and 3 p orbitals. sp3 hybridization has tetrahedral geometry.

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In terms of energy 2p Hybridization sp3 Energy 2s

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**How we get to hybridization**

We know the geometry from experiment. We know the orbitals of the atom hybridizing atomic orbitals can explain the geometry. So if the geometry requires a tetrahedral shape, it is sp3 hybridized This includes bent and trigonal pyramidal molecules because one of the sp3 lobes holds the lone pair.

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**sp2 hybridization C2H4 Double bond acts as one pair. trigonal planar**

Have to end up with three blended orbitals. Use one s and two p orbitals to make sp2 orbitals. Leaves one p orbital perpendicular.

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In terms of energy 2p 2p Hybridization sp2 Energy 2s

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**Where is the P orbital? Perpendicular**

The overlap of orbitals makes a sigma bond (s bond)

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**Two types of Bonds Sigma bonds from overlap of orbitals.**

Between the atoms. Pi bond (p bond) above and below atoms Between adjacent p orbitals. The two bonds of a double bond.

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H H C C H H

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**sp2 hybridization When three things come off atom. trigonal planar**

120º One p bond, s + lp =3

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**What about two When two things come off. One s and one p hybridize.**

linear

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**sp hybridization End up with two lobes 180º apart.**

p orbitals are at right angles Makes room for two p bonds and two sigma bonds. A triple bond or two double bonds.

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In terms of energy 2p 2p sp Hybridization Energy 2s

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CO2 C can make two s and two p O can make one s and one p O C O

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N2

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N2

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Breaking the octet PCl5 The model predicts that we must use the d orbitals. dsp3 hybridization There is some controversy about how involved the d orbitals are.

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**dsp3 Trigonal bipyrimidal can only s bond. can’t p bond.**

basic shape for five things.

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**PCl5 Can’t tell the hybridization of Cl**

Assume sp3 to minimize repulsion of electron pairs.

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d2sp3 gets us to six things around Octahedral Only σ bond

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**Molecular Orbital Model**

Localized Model we have learned explains much about bonding. It doesn’t deal well with the ideal of resonance, unpaired electrons, and bond energy. The MO model is a parallel of the atomic orbital, using quantum mechanics. Each MO can hold two electrons with opposite spins Square of wave function tells probability

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**What do you get? Solve the equations for H2 HA HB get two orbitals**

MO2 = 1sA - 1sB MO1 = 1sA + 1sB

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**The Molecular Orbital Model**

The molecular orbitals are centered on a line through the nuclei MO1 the greatest probability is between the nuclei MO2 it is on either side of the nuclei this shape is called a sigma molecular orbital

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**The Molecular Orbital Model**

In the molecule only the molecular orbitals exist, the atomic orbitals are gone MO1 is lower in energy than the 1s orbitals they came from. This favors molecule formation Called an bonding orbital MO2 is higher in energy This goes against bonding antibonding orbital

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**The Molecular Orbital Model**

Energy 1s 1s MO1

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**The Molecular Orbital Model**

We use labels to indicate shapes, and whether the MO’s are bonding or antibonding. MO1 = s1s MO2 = s1s* (* indicates antibonding) Can write them the same way as atomic orbitals H2 = s1s2

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**The Molecular Orbital Model**

Each MO can hold two electrons, but they must have opposite spins Orbitals are conserved. The number of molecular orbitals must equal the number atomic orbitals that are used to make them.

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H2- s1s* Energy 1s 1s s1s

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Bond Order The difference between the number of bonding electrons and the number of antibonding electrons divided by two

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**Only outer orbitals bond**

The 1s orbital is much smaller than the 2s orbital When only the 2s orbitals are involved in bonding Don’t use the s1s or s1s* for Li2 Li2 = (s2s)2 In order to participate in bonds the orbitals must overlap in space.

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**Bonding in Homonuclear Diatomic Molecules**

Need to use Homonuclear so that we know the relative energies. Li2- (s2s)2 (s2s*)1 Be2 (s2s)2 (s2s*)2 What about the p orbitals? How do they form orbitals? Remember that orbitals must be conserved.

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B2

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B2 s2p* s2p p2p* p2p

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**Expected Energy Diagram**

s2p* p2p* p2p* 2p 2p p2p p2p s2p Energy s2s* 2s 2s s2s

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B2 2p 2p Energy 2s 2s

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**B2 (s2s)2(s2s*)2 (s2p)2 Bond order = (4-2) / 2 Should be stable.**

This assumes there is no interaction between the s and p orbitals. Hard to believe since they overlap proof comes from magnetism.

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**Magnetism Magnetism has to do with electrons.**

Remember that spin is how an electron reacts to a magnetic field Paramagnetism attracted by a magnet. associated with unpaired electrons. Diamagnetism repelled by a magnet. associated with paired electrons. B2 is paramagnetic.

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Magnetism The energies of of the p2p and the s2p are reversed by p and s interacting The s2s and the s2s* are no longer equally spaced. Here’s what it looks like.

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**Correct energy diagram**

s2p* p2p* p2p* 2p s2p 2p p2p p2p s2s* 2s 2s s2s

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B2 s2p* p2p* 2p 2p s2p p2p s2s* 2s 2s s2s

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**Patterns As bond order increases, bond energy increases.**

As bond order increases, bond length decreases. Supports basis of MO model. There is not a direct correlation of bond order to bond energy. O2 is known to be paramagnetic. Movie.

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**Magnetism Ferromagnetic strongly attracted**

Paramagnetic weakly attracted Liquid Oxygen Diamagnetic weakly repelled Graphite Water Frog

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Examples C2 N2 O2 F2 P2

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**Heteronuclear Diatomic Species**

Simple type has them in the same energy level, so can use the orbitals we already know. Slight energy differences. NO

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NO 2p 2p 2s 2s

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You try NO+ CN- What if they come from completely different orbitals and energy? HF Simplify first by assuming that F only uses one if its 2p orbitals. F holds onto its electrons, so they have low energy

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s* 1s 2p s

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**Consequences Paramagnetic**

Since 2p is lower in energy, favored by electrons. Electrons spend time closer to fluorine. Compatible with polarity and electronegativity.

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**Names sp orbitals are called the Localized electron model**

s and p Molecular orbital model Localized is good for geometry, doesn’t deal well with resonance. seeing s bonds as localized works well It is the p bonds in the resonance structures that can move.

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p delocalized bonding C6H6 H H

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C2H6

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NO3-

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