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VSEPR model for geometry of a molecule or an ion

1. Sketch the Lewis structure.2. Count the total number of electron domains

around the central atom.a) One domain the bond of each atom attached to the

central atom (C.A.)b) One domain for each unshared pair of electrons

(nonbonding pair) on C.A.3. Describe the molecular geometry in terms of

the angular arrangement that maximizes the distance between bonding domains.

4. A double or triple bond is one domain.

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Effect of Nonbonding Electrons

• Unshared pairs of electrons will exert great repulsive forces on adjacent bonding domains.

• (Lone pairs push away other atoms.)• Nonbinding electrons will distort other bond

angles –they are compressed, or smaller than when each domain is a bonding one.

• Multiple bonds have a similar effect on adjacent domains.

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11 GeometriesBonding Domains

Nonbonding Domains

Molecular Geometry

2 0 or 3 Linear

3 0 Trigonal planar

2 1 or 2 Bent

4 0 Tetrahedral

3 1 Trigonal pyramidal

5 0 Trigonal bipyramidal

4 1 Seesaw

3 2 T-shaped

6 0 Octahedral

5 1 Square pyramidal

4 2 Square planar

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Hybrids

# bonds on Central Atom + # unshared electron pairs on Central Atom = # domains needed

Domains needed

Hybrid

1 none

2 sp

3 sp2

4 sp3

5 sp3d

6 sp3d2

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Covalent Bonding and Orbital Overlap

• Lewis structures and VSEPR theory gives us the shape of the molecule and the location of electrons in a molecule.

• They do not explain why a chemical bond forms.• How do we account then, for molecular shape in

terms of quantum mechanics? That is, which orbitals are involved in bonding?

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Valence-bond theory

• A covalent bond forms when the orbitals on two atoms overlap.– The shared region of space between the

orbitals is called the orbital overlap.– There are two electrons (usually one from

each atom) of opposite spin in the orbital overlap.

• As two nuclei approached each other their atomic orbitals overlap.

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• As the amount of overlap increases, the potential energy of the system decreases.

• At some distance the minimum energy is reached.– The minimum energy corresponds to the

bonding distance (or bond length).

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• As the two atoms get closer their nuclei begin to repel and the energy increases.

• At the bonding distance the attractive forces between nuclei and electrons just balance the repulsive forces (nucleus-nucleus, electron-electron).

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Multiple Bonds

• In the covalent bonds we have seen so far the electron density has been concentrated symmetrically about the internuclear axis.

• Sigma bonds: electron density lies on the axis between the nuclei.– All single bonds are Sigma bonds

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What about overlap in multiple bonds?

• Pi bonds: electron density lies above and below the plane of the nuclei.– A double bond consists of one sigma bond

and one pi bond.– A triple bond has one sigma and two pi bonds.

• The p orbitals involved in pi bonding come from unhybridized orbitals.

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Example

• Example: ethylene, C2H4 has :

– One sigma bond and one pi bond;– Both C atoms sp2 hybridized;– Both C atoms with trigonal planar electron pair

and molecular geometries

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Example

• Example: acetylene, C2H2:

• Electron-domain geometry of each C is linear.• Therefore, the C atoms are sp hybridized• The sp hybrid orbitals form the C-C and C-H

sigma bonds• There are two unhybridized p orbitals on each C

atom.

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Example, continued

• Both unhybridized p orbitals form the two pi bonds– One pi bond is above and below the plane of the

nuclei;– One pi bond is in front of and behind the plane of the

nuclei.

• When triple bonds form (e.g., N2) one pi bond is always above and below and the other is in front of and behind the plane of the nuclei.

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Delocalized pi bonding

• So far all the bonds we have encountered have been localized between two nuclei.

• In the case of benzene– There are six C-C sigma bonds and six C-H sigma

bonds– Each C atom is sp2 hybridized– One hybridized p orbital on each C atom, resulting in

six unhybridized carbon p orbitals in a ring.

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Benzene

• In benzene there are two options for the three pi bonds:– Localized between C atoms or– Delocalized over the entire ring (i.e., the pi electrons

are shared by all six C atoms).

• Experimentally, all C-C bonds are the same length in benzene.– Therefore, all C-C bonds are of the same type.

(Recall that single bonds are longer than double bonds.)

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General Conclusion

• Every pair of bonded atoms shares one or more pairs of electrons.

• The sharing of two electrons between atoms on the same axis as the nuclei results in sigma bond.

• Sigma bonds are always localized in the region between two bonded atoms.

• If two atoms share more than one pair of electrons, the additional pair form Pi bonds.

• When resonance structures are possible, delocalization of the electrons is also possible

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Molecular Orbitals

• Some aspects of bonding are not explained by Lewis structures, VSEPR theory, or hybridization

• For example:– Why does O2 interact with a magnetic field?

– Why are some molecules colored?

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Molecular orbital (MO) theory

• For these molecules we use molecular orbital theory

• Just as electrons in atoms are found in atomic orbitals, electrons in molecules are found in molecular orbitals.

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Molecular orbitals

• Molecular orbitals– Some characteristics are similar to those of atomic

orbitals:• Each contains a maximum of two electrons with opposite

spins.• Each has a definite energy• Electron density distribution can be visualized with contour

diagrams.

– However, unlike atomic orbitals, molecular orbitals are associated with an entire molecule.

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Bond order

• Bond order = ½(bonding electrons – antibonding electrons– Bond order = 1 for single bond– Bond order = 2 for double bond.– Bond order = 3 for triple bond.

• Fractional bonds orders are possible

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Example: Bond Order

• Example H2 molecule

– H2 has two bonding electrons.

– Bond order for H2 is:

– 1/2 (bonding electrons – antibonding electrons)= ½(2-0)=1

– Therefore H2 has a single bond.

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Example 2: Bond Order

• Consider the species He2

– He2 has two bonding electrons and two antibonding electrons.

– Bond order for He2 is: ½(bonding electrons – antibonding electrons)

– =1/2(2-2) =0

– Therefore He2 is not a stable molecule

– MO theory correctly predicts that hydrogen forms a diatomic molecule but that helium does not!

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Paramagnetism and Diamagnetism

• Paramagnetism occurs when one or more UNPAIRED electrons are attracted into a magnetic field. – Ex. B2: B-B 2s22p1 -- 2p12s2

– More unpaired electrons = Stronger attraction force

• Diamagnetism occurs when NO UNPAIRED electrons are weakly repelled from a magnetic field. – Ex. C2: C-C 2s22p2 -- 2p22s2