1. Chapter 6 Section 5 Molecular Geometry Modern Chemistry
Chemical Bonds
Chapter 6
Section 5 Molecular Geometry
Modern Chemistry

2. 6.5 Chemical Bonding & Molecular Geometry
Molecular geometry (shape) is integral to molecular function.
Think back to what you learned about enzymes and enzyme-substrate interactions in your biology class
If a molecule didn’t have the proper shape it couldn’t do the proper job
Trying to understand what influences chemical shape and, therefore, chemical behavior is an important goal in chemistry

3. 6.5 Chemical Bonding & Molecular Geometry
What factors, including shape, help determine the chemical and physical properties of molecular compounds?
Type of bonding between atoms
Three dimensional arrangement of a molecule’s atoms in space (= shape)
Molecular polarity

4. 6.5 Chemical Bonding & Molecular Geometry
What theories or models are used to explain molecular shape?
Valence Shell Electron Pair Repulsion Theory
helps explain the bond angles between atoms
Hybrid Orbital Theory
describes the orbitals that contain the valence electrons of a molecule’s atoms

5. 6.5 Chemical Bonding & Molecular Geometry
Objectives
Explain VSEPR theory
Predict the shapes of molecules or polyatomic ions using VSEPR theory
Explain how the shapes of molecules are accounted for by hybridization theory
Cont’d next slide …

6. 6.5 Chemical Bonding & Molecular Geometry
Objectives cont’d
Describe dipole-dipole forces, hydrogen bonding, induced dipoles, and London dispersion forces and their effects on properties such as boiling and melting points
Explain what determines molecular polarity

7. 10.1 VSEPR Model
The valence shell electron pair repulsion (VSEPR) model predicts the shapes of molecules from their Lewis structures
The main premise of the model is that electron pairs about an atom repel each other
A molecule has a shape that minimizes electrostatic repulsions between valence shell electron pairs
Minimum repulsion results when the electron pairs are as far apart as possible

8. 6.5 Chemical Bonding & Molecular Geometry
The following slides show molecules that have a central atom surrounded by either two, three, four, five, or six other atoms and their associated geometries/molecular shapes & bond angles
Please note that central atoms with five or six bonded atoms are exceptions to the octet rule
The only central atoms that you might see with five or six bonded atoms are phosphorous (P) or sulfur (S)
Note that in the notation used in the following slides A = the central atom and B = the # of atoms bonded to the central atom

9. 6.5 Chemical Bonding & Molecular Geometry
Central atom with two
bonded atoms = AB2
Shape
Linear
BeCl2
Central atom with three
bonded atoms = AB3
Trigonal Planar
BF3
Central atom with four
bonded atoms = AB4
Tetrahedral
CH4

10. 6.5 Chemical Bonding & Molecular Geometry
Central atom with five
bonded atoms = AB5
Trigonal
Bipyramidal
Examples
PF5
Central atom with six
bonded atoms = AB6
Octahedral
SF6

11. 6.5 Chemical Bonding & Molecular Geometry
The Lewis dot structures of some molecular compounds, indicate the presence of lone pairs of electrons in addition to bonded atoms
Ex: H2O NH3
H O: H N H
H H
Shorthand Notation: AB2E AB3E
a lone pair of electrons : :

12. 6.5 Chemical Bonding & Molecular Geometry
Note that in the shorthand notation, lone pairs are denoted by the letter E and the subscript tells how many lone pairs are present
Please note - only the lone pairs around the CENTRAL atom are counted!
When lone pairs of electrons are present, they influence the shape of the molecule although they are not part of the shape

13. 6.5 Chemical Bonding & Molecular Geometry
Therefore, water (H2O) is not linear as would be expected from an AB2 molecule. Rather it is bent in shape since it is an AB2E2 molecule
Bent
2 bonds, 2 lone pairs

14. 6.5 Chemical Bonding & Molecular Geometry
Similarly, ammonia (NH3) is not trigonal planar as would be expected from an AB3 molecule. Rather it is trigonal pyramidal in shape since it is an AB3E molecule
Trigonal pyramidal
3 bonds, 1 lone pair

15. 6.5 Chemical Bonding & Molecular Geometry
Lewis Dot Structure
Shorthand Notation
Molecular Geometry
(Shape)
Approximate
Bond Angles
Example
AB2
linear
180°
BeCl2
AB3
trigonal planar
120°
BH3
AB2E
bent or V-shaped
115° O3
AB4
tetrahedral
109°
CH4
AB3E
trigonal pyramidal
107°
NH3
AB2E2
105°
H2O
AB5
trigonal bipyramidal
180°/120°/90°
PCl5
AB6
octahedral
180°/90°
SF6

16. 6.5 Chemical Bonding & Molecular Geometry Molecular Polarity
Molecular geometry (shape) can influence whether a molecule is classified as polar
Polar molecules contain an unequal distribution of charge
Generally:
polar molecules interact with other polar molecules, and
nonpolar molecules react with other nonpolar molecules
Ex: sugar dissolves in water/oil does not

17. 6.5 Chemical Bonding & Molecular Geometry Molecular Polarity
Whether a molecule is polar or nonpolar depends on
the difference in electronegativities between bonded atoms, and
the overall shape of the molecule
The electronegativity differences or dipoles for each bonded pair can be represented by arrows pointing toward the more electronegative atom. Longer arrows represent more polar bonds
H F H Br

18. Electronegativity: The ability of an atom in a molecule to attract shared electrons to itself.

19. 6.5 Chemical Bonding & Molecular Geometry Molecular Polarity
The dipoles in CO2 cancel because the linear shape orients the equal magnitude bond dipoles in exactly opposite directions.
The CO2 molecule is, therefore, nonpolar

20. 6.5 Chemical Bonding & Molecular Geometry Molecular Polarity
The bond dipoles do not cancel in COSe; they are oriented in the same direction and are of unequal length. They do not cancel in OF2 because the V-shape of the molecule does not orient them in opposite directions.

21. 6.5 Chemical Bonding & Molecular Geometry Molecular Polarity
The bond dipoles in BCl3 and CCl4 cancel because of the regular shape and equal magnitude.

22. 6.5 Chemical Bonding & Molecular Geometry Molecular Polarity
The bond dipoles in BCl2F and CHCl3 do not cancel because they are not of the same magnitude.

23. 6.5 Chemical Bonding & Molecular Geometry Molecular Polarity
Based on the given examples:
Note that a molecule with polar bonds can be nonpolar if the geometry causes the bond polarities to sum to zero.
Note that molecules are nonpolar when there are no lone pairs on the central atom and all of the atoms bonded to the central atom are identical
Note that molecules with lone pairs of electrons on a central atom are generally polar, but there are exceptions

24. 6.5 Chemical Bonding & Molecular Geometry Molecular Polarity
Some problem solving…..
Predict which of the following molecules are polar and which are non-polar after drawing their Lewis dot structures.
HCN
PF3
SiBr4
1) polar 2) polar 3) nonpolar

25. 6.5 Chemical Bonding & Molecular Geometry
Intermolecular Forces
In addition to the chemical bonds that exist within a molecule, there are also interactions between molecules
The interactions or “forces of attraction” between molecules are called intermolecular forces (IMFs) and include
Hydrogen Bonding
Dipole-Dipole Forces
London Dispersion Forces

26. 6.5 Chemical Bonding & Molecular Geometry Intermolecular Forces
The number and strength of the intermolecular forces affect the properties of the substance.
The stronger the IMF’s are, the higher the melting and boiling points of a given substance
The stronger the IMF’s, the more energy is required to melt, evaporate or boil.
IMF’s are broken to go from solid  liquid and from liquid  gas.
Breaking IMF’s requires energy.

27. 6.5 Chemical Bonding & Molecular Geometry Intermolecular Forces
To highlight the difference between intramolecular bonds (within molecules) and intermolecular bonds (between molecules) let’s consider what happens when water evaporates?
Which bonds remain intact?
Which bonds are broken?

29. 6.5 Chemical Bonding & Molecular Geometry Intermolecular Forces
The function of many biologically important molecules is determined by their intermolecular forces as well (think H-bonding in DNA)
Although there are several different types of intermolecular attractions, all intermolecular attractions depend on electrostatic interactions, the attraction between charges of opposite sign

30. 6.5 Chemical Bonding & Molecular Geometry Intermolecular Forces/London Dispersion Forces
Electrons move around the nuclei. They could momentarily all “gang up” on one side
This lop-sidedness of electrons creates a partial negative charge in one area and a partial positive charge in another.
All molecules have electrons. +
Positively charged nucleus -
Negatively charged electron + - + -
Electrons are fairly evenly dispersed. + -
As electrons move, they “gang up” on one side.

31. 6.5 Chemical Bonding & Molecular Geometry Intermolecular Forces/London Dispersion Forces
The instantaneous “lopsidedness” of electron density can induce a near neighbor to behave similarly.

32. 6.5 Chemical Bonding & Molecular Geometry Intermolecular Forces/London Dispersion Forces
All molecules have electrons…all molecules can have London Dispersion Forces
The more electrons that gang-up, the larger the partial negative charge.
The larger the molecule, the stronger the London Dispersion Forces
Larger molecules have more electrons
Larger molecules have stronger London Dispersion Forces than smaller molecules.

33. 6.5 Chemical Bonding & Molecular Geometry Intermolecular Forces/London Dispersion Forces
Electrons can gang-up and cause a non-polar molecule to be temporarily polar
The electrons will move again, returning the molecule back to non-polar
The polarity was temporary, therefore the molecule cannot always form LDF.
London Dispersion Forces are the weakest of the intermolecular forces because molecules can’t form them all the time.

34. 6.5 Chemical Bonding & Molecular Geometry Intermolecular Forces / Dipole-Dipole Forces
Dipole-dipole attractions exist due to the partial positive and negative charges on atoms produced by unequal sharing of electrons due to electronegativity differences in polar molecules
The larger the dipole moment of a molecule, the stronger the dipole-dipole attraction. - +

35. 6.5 Chemical Bonding & Molecular Geometry Intermolecular Forces / Dipole-Dipole Forces
Polar molecules always have a partial separation of charge.
Polar molecules always have the ability to form attractions with opposite charges
Dipole forces are stronger than London Dispersion Forces

36. 6.5 Chemical Bonding & Molecular Geometry Intermolecular Forces / Hydrogen Bonding
Hydrogen bonding is an extreme example of a dipole-dipole interaction
Hydrogen bonding occurs between a hydrogen atom that is bonded to a small, highly electronegative atoms such as nitrogen, oxygen or fluorine and a second nitrogen, oxygen or fluorine on a neighboring atom

37. 6.5 Chemical Bonding & Molecular Geometry Intermolecular Forces / Hydrogen Bonding
Many unusual properties of water can be explained due to extensive hydrogen bonding
Ice is less dense than liquid water and therefore floats (most solids are denser than their corresponding liquids)
Water has a high specific heat that modulates Earth’s daytime and nighttime temperatures
Water has a high heat of vaporization that allows for evaporative cooling through sweating
Water is a versatile solvent
Water can travel by capillary action against gravity

38. 6.5 Chemical Bonding & Molecular Geometry Intermolecular Forces
The types of bonding forces vary in their strength as measured by average bond energy.
Ionic or (>200 kcal/mol)
Covalent bonds
Strongest to
Weakest
Hydrogen bonding (12-16 kcal/mol )
Dipole-dipole interactions (2-0.5 kcal/mol)
London forces (less than 1 kcal/mol)

39. 6.5 Chemical Bonding & Molecular Geometry Intermolecular Forces
Although intermolecular forces such as hydrogen-bonding, dipole-dipole interactions, and London dispersion forces are much weaker than ionic or covalent bonds, they are still important forces that explain much chemical behavior

40. 6.5 Chemical Bonding & Molecular Geometry Hybrid Orbital Theory
Valence shell electron pair repulsion theory explains much about molecular geometry but not everything…
Let’s look at bonding in a molecule of CH4 to explain why hybrid orbital theory was developed

41. What Proof Exists for Hybridization?
We have studied electron configuration notation and
the sharing of electrons in the formation of covalent
bonds.
Lets look at a molecule of methane, CH4.
Methane is a simple natural gas. Its molecule has a
carbon atom at the center with four hydrogen atoms covalently bonded around it.

42. Carbon ground state configuration
What is the expected orbital notation of carbon
in its ground state?
Can you see a problem with this?
(Hint: How many unpaired electrons does this
carbon atom have available for bonding?)

43. Carbon’s Bonding Problem
You should conclude that carbon only has TWO electrons available for bonding. That is not enough!
How does carbon overcome this problem so that
it may form four bonds?

44. Carbon’s Empty Orbital
The first thought that chemists had was that carbon promotes one of its 2s electrons…
…to the empty 2p orbital.

45. However, they quickly recognized a problem with such
an arrangement…
Three of the carbon-hydrogen bonds would involve
an electron pair in which the carbon electron was a 2p, matched with the lone 1s electron from a hydrogen atom.

46. This would mean that three of the bonds in a methane
molecule would be identical, because they would involve
electron pairs of equal energy.
But what about the fourth bond…?

47. The fourth bond is between a 2s electron from the
carbon and the lone 1s hydrogen electron.
Such a bond would have slightly less energy than the other bonds in a methane molecule.

48. This bond would be slightly different in character than
the other three bonds in methane.
This difference would be measurable to a chemist
by determining the bond length and bond energy.
But is this what they observe?

49. The simple answer is, “No”.
Measurements show that
all four bonds in methane
are equal. Thus, we need
a new explanation for the
bonding in methane.
Chemists have proposed an explanation – they call it
Hybridization.
Hybridization is the combining of two or more orbitals
of nearly equal energy within the same atom into
orbitals of equal energy.

50. Hybridization - The Blending of Orbitals+ =
Poodle +
Cocker Spaniel =
Cockapoo + =
s orbital +
p orbital =
sp orbital

51. In the case of methane, they call the hybridization
sp3, meaning that an s orbital is combined with three
p orbitals to create four equal hybrid orbitals.
These new orbitals have slightly MORE energy than
the 2s orbital…
… and slightly LESS energy than the 2p orbitals.

52. sp3 Hybrid Orbitals
Here are some other ways to look at the sp3 hybridization
and energy profile…

53. sp Hybrid Orbitals While sp3 is the hybridization observed in methane,
there are other types of hybridization that atoms
undergo.
These include sp hybridization, in which one s orbital combines with a single p orbital.
This produces two hybrid orbitals, while leaving two normal p orbitals

54. sp2 Hybrid Orbitals
Another hybrid is the sp2, which combines two orbitals from a p sublevel with one orbital from an s sublevel.
One p orbital remains unchanged.

55. Summary of Hybridization and Molecular Geometry
Forms
Overall Structure
Hybridization of “A”
AX2
Linear sp
AX3, AX2E
Trigonal Planar
sp2
AX4, AX3E, AX2E2
Tetrahedral
sp3
A = central atom
X = atoms bonded to A
E = nonbonding electron pairs on A