1. Section 5 Molecular Geometry
Chapter 6
VSEPR Theory
As shown at right, diatomic molecules, like those of (a) hydrogen, H2, and (b) hydrogen chloride, HCl, can only be linear because they consist of only two atoms.
To predict the geometries of more-complicated molecules, one must consider the locations of all electron pairs surrounding the bonding atoms. This is the basis of VSEPR theory.

2. Lewis Structures show bonding and lone pairs, but do not denote shape.
Section 5 Molecular Geometry
Chapter 6
Lewis Structures show bonding and lone pairs, but do not denote shape.
However, we use Lewis Structures to help us determine shapes.
Here we see some common shapes for molecules with two or three atoms connected to a central atom.

3. What Determines the Shape of a Molecule?
Section 5 Molecular Geometry
Chapter 6
What Determines the Shape of a Molecule?
Simply put, electron pairs, whether they be bonding or nonbonding, repel each other.
By assuming the electron pairs are placed as far as possible from each other, we can predict the shape of the molecule.
This is the Valence-Shell Electron-Pair Repulsion (VSEPR) model.

4. Chapter 6
Electron Domains
We can refer to the directions to which electrons point as electron domains. This is true whether there is one or more electron pairs pointing in that direction.
The central atom in this molecule, A, has four electron domains.

5. Valence-Shell Electron-Pair Repulsion (VSEPR) Model
Chapter 6
Valence-Shell Electron-Pair Repulsion (VSEPR) Model
“The best arrangement of a given number of electron domains is the one that minimizes the repulsions among them.”
(The balloon analogy in the figure to the left demonstrates the maximum distances, which minimize repulsions.)

6. Electron-Domain Geometries
Chapter 6
The Table shows the electron-domain geometries for two through six electron domains around a central atom.
To determine the electron-domain geometry, count the total number of lone pairs, single, double, and triple bonds on the central atom.

7. Molecular Geometries Chapter 6
Once you have determined the electron-domain geometry, use the arrangement of the bonded atoms to determine the molecular geometry.
Tables 9.2 and 9.3 show the potential molecular geometries. We will look at each electron domain to see what molecular geometries are possible.

8. Linear Electron Domain
Chapter 6
In the linear domain, there is only one molecular geometry: linear.
NOTE: If there are only two atoms in the molecule, the molecule will be linear no matter what the electron domain is.

9. Trigonal Planar Electron Domain
Chapter 6
There are two molecular geometries:
trigonal planar, if all electron domains are bonding, and
bent, if one of the domains is a nonbonding pair.

10. Tetrahedral Electron Domain
Chapter 6
There are three molecular geometries:
tetrahedral, if all are bonding pairs,
trigonal pyramidal, if one is a nonbonding pair, and
bent, if there are two nonbonding pairs.

11. Nonbonding Pairs and Bond Angle
Chapter 6
Nonbonding pairs are physically larger than bonding pairs.
Therefore, their repulsions are greater; this tends to compress bond angles.

12. Multiple Bonds and Bond Angles
Chapter 6
Double and triple bonds have larger electron domains than single bonds.
They exert a greater repulsive force than single bonds, making their bond angles greater.

13. Expanding beyond the Octet Rule
Chapter 6
Expanding beyond the Octet Rule
Remember that some elements can break the octet rule and make more than four bonds (or have more than four electron domains).
The result is two more possible electron domains: five = trigonal bipyramidal; six = octahedral (as was seen in the slide on electron-domain geometries).

14. Trigonal Bipyramidal Electron Domain
Chapter 6
Trigonal Bipyramidal Electron Domain
There are two distinct positions in this geometry:
Axial
Equatorial
Lone pairs occupy equatorial positions.

15. Trigonal Bipyramidal Electron Domain
Chapter 6
Trigonal Bipyramidal Electron Domain
There are four distinct molecular geometries in this domain:
Trigonal bipyramidal
Seesaw
T-shaped
Linear

16. Octahedral Electron Domain
Chapter 6
Octahedral Electron Domain
All positions are equivalent in the octahedral domain.
There are three molecular geometries:
Octahedral
Square pyramidal
Square planar

17. VSEPR Theory, continued
Section 5 Molecular Geometry
Chapter 6
VSEPR Theory, continued
Sample Problem E
Use VSEPR theory to predict the molecular geometry of boron trichloride, BCl3.

18. VSEPR Theory, continued
Section 5 Molecular Geometry
Chapter 6
VSEPR Theory, continued
Sample Problem E Solution
Boron trichloride is an AB3 type of molecule.
Its geometry should therefore be trigonal-planar.

19. VSEPR and Lone Electron Pairs
Section 5 Molecular Geometry
Chapter 6
VSEPR and Lone Electron Pairs
Click below to watch the Visual Concept.
Visual Concept

20. VSEPR Theory, continued
Section 5 Molecular Geometry
Chapter 6
VSEPR Theory, continued
Sample Problem F
Use VSEPR theory to predict the shape of a molecule of carbon dioxide, CO2.
Use VSEPR theory to predict the shape of a chlorate ion,

21. VSEPR Theory, continued
Section 5 Molecular Geometry
Chapter 6
VSEPR Theory, continued
Sample Problem F Solution
Draw the Lewis structure of carbon dioxide.
There are two carbon-oxygen double bonds and no unshared electron pairs on the carbon atom.
This is an AB2 molecule, which is
linear.

22. VSEPR Theory, continued
Section 5 Molecular Geometry
Chapter 6
VSEPR Theory, continued
Sample Problem F Solution, continued
Draw the Lewis structure of the chlorate ion.
There are three oxygen atoms bonded to the central chlorine atom, which has an unshared electron pair.
This is an AB3E molecule, which is
trigonal-pyramidal.

23. VSEPR and Hybrid Orbitals
Section 5 Molecular Geometry
Chapter 6
VSEPR and Hybrid Orbitals
VSEPR predicts shapes of molecules very well.
How does that fit with orbitals?
Let’s use H2O as an example:
If we draw the best Lewis structure to assign VSEPR, it becomes bent.
If we look at oxygen, its electron configuration is 1s22s22p4. If it shares two electrons to fill its valence shell, they should be in 2p.
Wouldn’t that make the angle 90°?
Why is it 104.5°?

24. Chapter 6 Hybrid Orbitals
Section 5 Molecular Geometry
Chapter 6
Hybrid Orbitals
Hybrid orbitals form by “mixing” of atomic orbitals to create new orbitals of equal energy, called degenerate orbitals. When two orbitals “mix” they create two orbitals; when three orbitals mix, they create three orbitals; etc.

25. Chapter 6 Be—sp hybridization
Section 5 Molecular Geometry
Chapter 6
Be—sp hybridization
When we look at the orbital diagram for beryllium (Be), we see that there are only paired electrons in full sub-levels.
Be makes electron deficient compounds with two bonds for Be. Why? sp hybridization (mixing of one s orbital and one p orbital)

26. These sp hybrid orbitals have two lobes like a p orbital.
Section 5 Molecular Geometry
Chapter 6
sp Orbitals
Mixing the s and p orbitals yields two degenerate orbitals that are hybrids of the two orbitals.
These sp hybrid orbitals have two lobes like a p orbital.
One of the lobes is larger and more rounded, as is the s orbital.

27. Position of sp Orbitals
Section 5 Molecular Geometry
Chapter 6
Position of sp Orbitals
These two degenerate orbitals would align themselves 180 from each other.
This is consistent with the observed geometry of Be compounds (like BeF2) and VSEPR: linear.

28. Boron—Three Electron Domains Gives sp2 Hybridization
Section 5 Molecular Geometry
Chapter 6
Boron—Three Electron Domains Gives sp2 Hybridization
Using a similar model for boron leads to three degenerate sp2 orbitals.

29. Carbon: sp3 Hybridization
Section 5 Molecular Geometry
Chapter 6
Carbon: sp3 Hybridization
With carbon, we get four degenerate sp3 orbitals.

30. What Happens with Water?
Section 5 Molecular Geometry
Chapter 6
What Happens with Water?
We started this discussion with H2O and the angle question: Why is it 104.5° instead of 90°? Oxygen has two bonds and two lone pairs—four electron domains. The result is sp3 hybridization!

31. Chapter 6 Hybrid Orbitals Section 5 Molecular Geometry
Click below to watch the Visual Concept.
Visual Concept

32. Section 5 Molecular Geometry
Chapter 6

33. Schematic Representations of the Three States of Matter
Section 5 Molecular Geometry
Chapter 6
Schematic Representations of the Three States of Matter
Hfus= 6.02kJ/mole Hvap = 40.7kJ/mole Hrxn to H and O = 934 kJ/mole
Copyright © Cengage Learning. All rights reserved

34. Types of Intermolecular Force
Section 5 Molecular Geometry
Chapter 6
Types of Intermolecular Force
Strongest to weakest forces:
dipole–dipole forces
Dipole-induced dipole
ion–dipole forces
hydrogen bonding (a special dipole–dipole force)
dispersion forces (or London dispersion forces)
Hfus= 6.02kJ/mole Hvap = 40.7kJ/mole Hrxn to H and O = 934 kJ/mole
Copyright © Cengage Learning. All rights reserved

35. Intermolecular Forces
Section 5 Molecular Geometry
Chapter 6
Intermolecular Forces
The forces of attraction between molecules are known as intermolecular forces.
The boiling point of a liquid is a good measure of the intermolecular forces between its molecules: the higher the boiling point, the stronger the forces between the molecules.
Intermolecular forces vary in strength but are generally weaker than bonds between atoms within molecules, ions in ionic compounds, or metal atoms in solid metals.
Boiling points for ionic compounds and metals tend to be much higher than those for molecular substances: forces between molecules are weaker than those between metal atoms or ions.

36. Comparing Ionic and Molecular Substances
Section 5 Molecular Geometry
Chapter 6
Comparing Ionic and Molecular Substances

37. Intermolecular Forces, continued
Section 5 Molecular Geometry
Chapter 6
Intermolecular Forces, continued
The strongest intermolecular forces exist between polar molecules.
Because of their uneven charge distribution, polar molecules have dipoles. A dipole is created by equal but opposite charges that are separated by a short distance.
The direction of a dipole is from the dipole’s positive pole to its negative pole.

38. Intermolecular Forces, continued
Section 5 Molecular Geometry
Chapter 6
Intermolecular Forces, continued
A dipole is represented by an arrow with its head pointing toward the negative pole and a crossed tail at the positive pole. The dipole created by a hydrogen chloride molecule is indicated as follows:

39. Section 5 Molecular Geometry
Chapter 6

40. Intermolecular Forces, continued
Section 5 Molecular Geometry
Chapter 6
Intermolecular Forces, continued
The negative region in one polar molecule attracts the positive region in adjacent molecules. So the molecules all attract each other from opposite sides.
Such forces of attraction between polar molecules are known as dipole-dipole forces.

41. Intermolecular Forces, continued
Section 5 Molecular Geometry
Chapter 6
Intermolecular Forces, continued
Dipole-dipole forces act at short range, only between nearby molecules. Rapidly becomes weaker as distance between molecules increases.
Dipole-dipole forces explain, for example the difference between the boiling points of iodine chloride, I–Cl (97°C), and bromine, Br–Br (59°C).

42. Comparing Dipole-Dipole Forces
Section 5 Molecular Geometry
Chapter 6
Comparing Dipole-Dipole Forces

43. Chapter 6 Dipole-Dipole Forces Section 5 Molecular Geometry
Click below to watch the Visual Concept.
Visual Concept

44. Intermolecular Forces, continued
Section 5 Molecular Geometry
Chapter 6
Intermolecular Forces, continued
A polar molecule can induce a dipole in a nonpolar molecule by temporarily attracting its electrons.
The result is a short-range intermolecular force that is somewhat weaker than the dipole-dipole force.
Induced dipoles account for the fact that a nonpolar molecule, oxygen, O2, is able to dissolve in water, a polar molecule.

45. Dipole-induced Dipole
Section 5 Molecular Geometry
Chapter 6
Dipole-induced Dipole

46. Dipole-Induced Dipole Interaction
Section 5 Molecular Geometry
Chapter 6
Dipole-Induced Dipole Interaction
Click below to watch the Visual Concept.
Visual Concept

47. Ion–Dipole Interactions
Section 5 Molecular Geometry
Chapter 6
Ion–Dipole Interactions
Ion–dipole interactions are found in solutions of ions.
The strength of these forces is what makes it possible for ionic substances to dissolve in polar solvents.

48. Intermolecular Forces, continued
Section 5 Molecular Geometry
Chapter 6
Intermolecular Forces, continued
Some hydrogen-containing compounds have unusually high boiling points. This is explained by a particularly strong type of dipole-dipole force.
In compounds containing H–F, H–O, or H–N bonds, the large electronegativity differences between hydrogen atoms and the atoms they are bonded to make their bonds highly polar.
This gives the hydrogen atom a strong positive charge.

49. Intermolecular Forces, continued
Section 5 Molecular Geometry
Chapter 6
Intermolecular Forces, continued
The small size of the hydrogen atom allows the atom to come very close to an unshared pair of electrons in an adjacent molecule.
The intermolecular force in which a hydrogen atom that is bonded to a highly electronegative atom is attracted to an unshared pair of electrons of an electronegative atom in a nearby molecule is known as hydrogen bonding.

50. Intermolecular Forces
Section 5 Molecular Geometry
Chapter 6
Intermolecular Forces
Hydrogen bonds are usually represented by dotted lines connecting the hydrogen-bonded hydrogen to the unshared electron pair of the electronegative atom to which it is attracted.
An excellent example of hydrogen bonding is that which occurs between water molecules. The strong hydrogen bonding between water molecules accounts for many of water’s characteristic properties.

51. Intermolecular Forces
Section 5 Molecular Geometry
Chapter 6
Intermolecular Forces

52. Visual Concepts
Chapter 6
Hydrogen Bonding

53. Intermolecular Forces, continued
Section 5 Molecular Geometry
Chapter 6
Intermolecular Forces, continued
London Dispersion Forces
Even noble gas atoms and nonpolar molecules can experience weak intermolecular attraction.
In any atom or molecule—polar or nonpolar—the electrons are in continuous motion.
As a result, at any instant the electron distribution may be uneven. A momentary uneven charge can create a positive pole at one end of an atom of molecule and a negative pole at the other.

54. Intermolecular Forces, continued
Section 5 Molecular Geometry
Chapter 6
Intermolecular Forces, continued
London Dispersion Forces, continued
This temporary dipole can then induce a dipole in an adjacent atom or molecule. The two are held together for an instant by the weak attraction between temporary dipoles.
The intermolecular attractions resulting from the constant motion of electrons and the creation of instantaneous dipoles are called London dispersion forces.

55. Occurs in all molecules, including nonpolar ones.
Section 5 Molecular Geometry
Chapter 6
Occurs in all molecules, including nonpolar ones.
The figure below shows how a nonpolar particle (in this case a helium atom) can be temporarily polarized to allow dispersion force to form.
The tendency of an electron cloud to distort is called its polarizability.

56. later was a professor at Duke
Section 5 Molecular Geometry
Chapter 6
Fritz London first proposed their existence in 1930.
Fritz London
Born in Germany,
later was a professor at Duke

57. London Dispersion Force
Section 5 Molecular Geometry
Chapter 6
London Dispersion Force
Click below to watch the Visual Concept.
Visual Concept