molecular geometry and bonding theories - university of illinois

1

Molecular Geometry and Molecular Geometry and Bonding TheoriesBonding Theories

Chapter 9

2

In this chapterIn this chapter• Molecular shapes• Valence Shell Electron Pair Repulsion model.• Molecular shapes and polarity.• Valence bond theory.• Hybridization.

3

• Lewis structures give atomic connectivity: they tell us which atoms are physically connected to which.

• The shape of a molecule is determined by its bond angles.

• Consider CCl4: experimentally we find all Cl-C-Clbond angles are 109.5°.– Therefore, the molecule cannot be planar.– All Cl atoms are located at the vertices of a tetrahedron

with the C at its center.

Molecular ShapesMolecular Shapes

2

4


5

• The Valence Shell Electron Pair Repulsion (VSEPR) theory proposed by Gillespie and Nyholm is a reliable method for predicting shapes of covalent molecules and polyatomic ions.

• The VSEPR theory is based on the idea that bond and lone pair electrons in the valence shell of an element repel each other and move as far apart as possible.

• VSEPR is extremely successful in predicting the structures of molecules and ions of the main group elements, but is generally useless for transition metal containing compounds.


6

There are five fundamental geometries for molecular shape:

AB2

AB6AB5

AB4AB3

3

7

PredictingPredicting the Shapethe Shape• Predict the shape of CH2Cl2• CH2Cl2 has one central atom and 4 other atoms

around it. Draw the Lewis structure. • Referring to the possible arrangements we can

have for such a molecule the shape can be predicted.

8

• When considering the geometry about the central atom, we consider all electrons (lone pairs and bonding pairs).

• When naming the molecular geometry, we focus only on the positions of the atoms.


9

• To determine the shape of a molecule, need to distinguish between lone pairs (or non-bonding pairs, those not in a bond) of electrons and bonding pairs (those found between two atoms).

• The electron domain geometry is defined by the positions in 3D space of ALL electron pairs (bonding + non-bonding). This becomes especially important when the central atom has lone pairs.

• The electrons adopt an arrangement in space to minimize e--e- repulsion.


4

10

11

12

Lone pairs on Central AtomLone pairs on Central Atom

5

13

14

• The electron pair geometry is decided only by looking at the positions of the electrons.

• The molecular geometry is decided only by the positions of atoms, lone pairs are ignored in the molecular geometry.

• Lone pairs are considered to be “fat”, while bonding pairs are considered “skinny”.

• The increased volume of lone pairs makes bond pairs squeeze together, reducing bond angles.

Lone Pairs & Bond AnglesLone Pairs & Bond Angles

15


6

16

• Experimentally, the H-X-H bond angle decreases on moving from C to N to O:

104.5O107O

NHH

HC

H

HHH109.5O

OHH


• Since the electrons in a bonding pair are attracted by two nuclei they do not repel as much as lone pairs.

• The bond angle decreases as the number of lone pairs on the central atom increases.

• Typically the relative strength of repulsion is:

17

• Similarly, electrons in multiple bonds repel more than electrons in single bonds.

• A multiple bond is counted as one electron domain.• The Lewis structure of COCl2 (carbonyl chloride)

indicates the molecular and the electron pair geometry to be trigonal-planar.

C OCl

Cl111.4o

124.3o

Multiple Bonds & Bond AnglesMultiple Bonds & Bond Angles

18

• Atoms that have expanded octets have AB5 (trigonal bipyramidal) or AB6 (octahedral) electron pair geometries.

• For trigonal bipyramidal structures there is a plane containing three electrons pairs. The fourth and fifth electron pairs are located above and below this plane.

• For octahedral structures, there is a plane containing four electron pairs. Similarly, the fifth and sixth electron pairs are located above and below this plane.

Molecules with Expanded Valence Shells

7

19

20

21

• To minimize e--e- repulsion, lone pairs are always placed in equatorial positions.

Molecules with Expanded Valence Molecules with Expanded Valence ShellsShells

• In an octahedron, all positions are considered to be equivalent and there are no axial or equatorial positions.

• If the molecule has any lone pairs they are placed above or below the plane.

8

22

• In acetic acid, CH3COOH, there are three central atoms.

• We assign the geometry about each central atom separately.

Shapes of Large MoleculesShapes of Large Molecules

23

Determining Molecular ShapesDetermining Molecular Shapes• Draw the Lewis structure.• Determine the total number of bonding pairs

(count any multiple bonds as 1 pair) around the central atom.

• Pick the appropriate electron pair geometry and then choose the molecular shape that matches the total # of single-bond pairs and lone pairs.

• Predict the bond angles, remembering that the lone pairs occupy more volume than bonding pairs.

24

• Difference in electronegativity between two atoms makes the bond between them polar.

• It is possible for a molecule to contain polar bonds, but not be polar.

• For example, the bond dipoles in CO2 cancel each other because CO2 is linear.

Molecular Shape & PolarityMolecular Shape & Polarity

9

25

• In water, the molecule is not linear and the bond dipoles do not cancel each other.

• Therefore, water is a polar molecule.

Molecular Shape & PolarityMolecular Shape & Polarity

26

Molecular Shape and Molecular Molecular Shape and Molecular PolarityPolarity

27

• Why does a bond form?• How do we account for molecular shape in

terms of quantum mechanics, meaning in terms of orbital shapes?

• What are the orbitals that are involved in bonding?

• Valence Bond Theory proposed by LinusPauling provides these answers.

• Provides a qualitative visual picture of molecular structure and bonding.

Questions not answered by VSEPRQuestions not answered by VSEPR

10

28

Valence Bond (VB) TheoryValence Bond (VB) Theory• Orbitals overlap to form a bond between

atoms.• Two electrons of opposite spins can be

accommodated in the overlapping orbitals. Usually one electron is supplied by each of the 2 bonded atoms.

• Overlapping orbitals result in a high probability of the 2 electrons being located in the region that is influenced by both nuclei. Both electrons are simultaneously attracted by both nuclei.

29

30

Orbital OverlapOrbital Overlap• Orbital overlap produces a single bond that is known

as a sigma (σ) bond.• The electron density of the sigma bond is greatest

along its axis.

The 2s and 2p electrons that are not involved in the bond are the lone pairs on the F atoms.

11

31

Hybrid OrbitalsHybrid Orbitals• Consider the CH4 molecule, which has a tetrahedral shape.• C has the electron configuration 1s22s22p2 and H has 1s1.

Carbon has only 2 unpaired electrons.• The 2p orbitals in C are at right angles to each other, the 1s

orbital in H is spherical.• According to Pauling’s theory of orbital hybridization, a new

set of orbitals called hybrid orbitals can be created by combining the appropriate # of s, p, d and f orbitals.

• This combination results in hybrid orbitals with a spatial orientation that matches the geometry of the molecule.

•• The number of hybrid orbitals is the same as the number of The number of hybrid orbitals is the same as the number of atomic orbitals used in their constructionatomic orbitals used in their construction..

32

Hybridization in MethaneHybridization in Methane• By mixing the (one) 2s and (three) 2p valence

orbitals in the C atom a new set of hybrid orbitals called the (four) sp3 orbitals is created.

• The four sp3 hybrid orbitals are tetrahedral in shape.

• All four sp3 orbitals have identical shapes and are equivalent in energy.

• Each H atom bonds to the C atom by orbital overlap between the 1s orbital and the sp3

hybrid orbitals.

33

spsp2 2 Hybrid OrbitalsHybrid Orbitals• sp2 hybrid orbitals are formed by the mixing of

one s and two p orbitals. • The large lobes of sp2 hybrid orbitals lie in a

trigonal plane and the angles between them are 120°.

• All molecules with trigonal planar electron pair geometries have sp2 orbitals on the central atom.

• One unhybridized p orbital is left as a result of the formation of sp2 orbitals.

12

34

35

• Formed by the combination of one s and one p orbital.

• The lobes of sp hybrid orbitals are 180º apart, meaning the geometry is linear.

• All molecules with linear geometry will have sp hybrid orbitals on the central atom.

• The formation of sp hybrid orbitals leaves 2 unhybridized p orbitals on the central atom.

sp Hybrid Orbitalssp Hybrid Orbitals

36

13

37

• Since there are only three p-orbitals, trigonal bipyramidal and octahedral electron pair geometries must involve d-orbitals.

• Trigonal bipyramidal electron pair geometries require sp3d hybridization.

• Octahedral electron pair geometries require sp3d2 hybridization.

• The electron pair geometry from VSEPR theory always determines the hybridization.

spsp33d and spd and sp33dd22 Hybrid OrbitalsHybrid Orbitals

38

39

14

40

Deciding HybridizationDeciding Hybridization• Decide the hybridization of the central atom

for SF3+ and I3- ions.

F S F

F

+

I

I

I

-

41

Multiple BondsMultiple Bonds• σ-bonds: electron density lies on

the axis between the nuclei. All single bonds are σ-bonds.

• π-bonds: electron density lies above and below the plane of the nuclei.

• A double bond consists of one σ-bond and one π-bond.

• A triple bond has one σ-bond and two π-bonds.

• Often, the p-orbitals involved in π-bonding come from unhybridized orbitals.

42

Multiple BondsMultiple Bonds• Ethylene, C2H4, has:

– one σ- and one π-bond– both C atoms are sp2 hybridized– both C atoms have trigonal planar electron pair and

molecular geometries– The formation of a π bond requires that the molecule be

planar.

15

43

44

Multiple BondsMultiple BondsConsider acetylene, C2H2–The electron pair geometry of each C is linear & the C atoms are sp hybridized;

–The sp hybrid orbitals form the C-C and C-H σ-bonds;–There are two unhybridized p-orbitals which form the two π-bonds.

–one π-bond is above and below the plane of the nuclei;–one π-bond is in front and behind the plane of the nuclei.

45

Delocalized Delocalized ππ BondingBonding• So far all the bonds we have encountered are

localized between two nuclei.• In the case of benzene

– there are 6 C-C σ bonds, 6 C-H σ bonds,– each C atom is sp2 hybridized,– and there are 6 unhybridized p orbitals on each C atom.

16

46

Delocalized Delocalized ππ BondingBonding

• In benzene there are two options for the 3 π bonds– localized between C atoms or– delocalized over the entire ring (i.e. the π

electrons are shared by all 6 C atoms).• Experimentally, all C-C bonds are the same length in

benzene.• Therefore, all C-C bonds are of the same type (recall

single bonds are longer than double bonds).

47

Multiple BondsMultiple Bonds• Every two atoms share at least 2 electrons.• Two electrons between atoms on the same

axis as the nuclei are σ bonds.• σ-bonds are always localized.• If two atoms share more than one pair of

electrons, the second and third pair form π-bonds.

• When resonance structures are possible, delocalization is also possible.

molecular geometry and bonding theories - university of illinois

Documents