1. Hybridization
Covalent bonds are formed by the sharing of electrons; orbitals overlap to allow for this sharing.
The mixing of two or more atomic orbitals of an atom forms hybrid orbitals. This process is called hybridization.
Predicting hybridization is easy – just count the total number of steric numbers (domains).
The AP test does not cover hybridization past 4 domains.

2. The total number of steric numbers (also known as substituents - bonding plus non-bonding groups) is equal to the number of atomic orbitals that participate in the hybrid orbital.
# of substituents
(steric numbers)
Hybridization
Example Molecule 2 sp
CO2 3
sp2
CH3 4
sp3
CH4, NH3, H2O 5
sp3d
PCl5, I3- 6
sp3d2
SF6

3. The sp hybrid orbital
In sp hybridization, only one p orbital is mixed with the s orbital
Example: BeF Steric Number: 2
Electron configuration of Be: 1s22s2
Electron configuration of F:1s22s22p5
                                                              

4. In the ground state, Be has no unpaired electrons – so how can the Be atom form a covalent bond with a fluorine?
Be obtains an unpaired electron by moving one electron from the 2s orbital to the 2p orbital resulting in two unpaired electrons, one in a 2s orbital and another in a 2p orbital Be F

5. Be F The Be atom can now form two covalent bonds with fluorine atoms
Although we would not expect these bonds to be identical (one is in a 2s electron orbital, the other is in a 2p electron orbital), the structure of BeF2 is linear and the bond lengths are identical
The 2s and 2p electrons produced a "hybrid" orbital for both electrons Be F

6. The sp2 hybrid orbital
In sp2 hybridization, two p orbitals are mixed with the s orbital to generate three new hybrids
Example: BF3

7. The sp3 hybrid orbital
In sp3 hybridization, all three p orbitals are mixed with the s orbital to generate four new hybrids
Example: CH4

9. Sigma (σ) and Pi Bonds (π)
Sigma bond: The first bond made with any other atom
Made from hybridized orbitals
s-s, s-p, or p-p head-on overlap between nucleus
Allows for free rotation
Pi bond: Any 2nd or 3rd bond made with any other atom
Made from leftover p orbitals
parallel, sideways p-p overlap, nucleus above or below overlap
Weaker bond than sigma
Fixed rotation
Single bond
1 sigma bond
Double bond
1 sigma bond and 1 pi bond
Triple bond
1 sigma bond and 2 pi bonds

10. Sigma (σ) and Pi Bonds (π)
How many σ and π bonds are in the acetic acid (vinegar) molecule CH3COOH? C H O
σ bonds = 6
+ 1 = 7
π bonds = 1

11. Given the structural formula for propyne: 1
Given the structural formula for propyne:     1.  Indicate the hybridization of each carbon atom in the structure above.   2.  Indicate the total number of sigma (σ) and pi (π) bonds in the molecule.  

12. Bond length and strength
The more electrons that are involved in bonding, the shorter the bond length and the stronger the bond (meaning higher bond energy).
Problem: Is the bond length between the two carbon atoms shorter in C2H6, C2H4, or C2H2. Why?
Problem: Use the bonding model to account for the fact that all the bond lengths in SO3 are identical and are shorter than a sulfur-oxygen single bond.

13. Problem
A. Draw the Lewis electron-dot structures for CO32-, CO2 and CO, including resonance structures where appropriate.
Which of the three species has the shortest C-O bond length? Explain the reason for your answer.
Account for the fact that the carbon-oxygen bond length in CO32– is greater than the carbon-oxygen bond length in CO2.

14. Warm Up
Draw the following structures according to shape. Use arrows to indicate bond polarity. Arrows point to the more electronegative atom. Name shape and bond angle.
1. CH BF CO HCN 5. CH2Cl CH2O 7. NH3

15. Bond Polarity and Dipole Moments
Polar Molecules: Molecules with a somewhat negative end and a somewhat positive end (a dipole moment)
Use an arrow to represent a dipole moment.
Point to the negative charge center with the tail of the arrow indicating the positive center of charge.
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16. Bond Polarity and Dipole Moments

17. No Net Dipole Moment (Dipoles Cancel)
Bond Polarity and Dipole Moments
No Net Dipole Moment (Dipoles Cancel)
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18. Polar Molecules Polar molecules have a permanent dipole.
Polar molecules line up in the presence of an electric field; nonpolar molecules do not.

19. Molecules that have a polar bond may or may not show a dipole moment – the shape of the molecule must be considered.

21. Polar vs. Nonpolar Molecules
Nonpolar molecules have a symmetrical charge distribution
Diatomic molecules with the same atoms are nonpolar.
Ex: Cl2
Linear, tetrahedral, trigonal planar shapes must have the same peripheral atoms to be nonpolar
Ex: CH4 BF3 CO2

22. Polar molecules have an asymmetrical charge distribution.
Polar vs. Nonpolar Molecules
Polar molecules have an asymmetrical charge distribution.
Always polar: trigonal pyramidal and bent
Ex: NH3, H2O
Diatomic molecules with different atoms are polar.
Ex: HCl
Linear, tetrahedral, and trigonal planar shapes have different peripheral atoms to be polar
Ex: HCN, CH2Cl2, CH2O

23. Polar vs. Nonpolar Molecules
The melting points of polar substances are higher than the melting points of non-polar substances with similar sizes. Concept Check: Which of the following would have the higher boiling point?

25. CONCEPT CHECK!
True or false: A molecule that has polar bonds will always be polar. -If true, explain why. -If false, provide a counter-example. Answer: False, a molecule may have polar bonds (like CO2) but the individual dipoles might cancel out so that the net dipole moment is zero.
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26. Draw the Lewis structure for CO2. Does CO2 contain polar bonds?
Let’s Think About It
Draw the Lewis structure for CO2.
Does CO2 contain polar bonds?
Is the molecule polar or nonpolar overall? Why?
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27. CONCEPT CHECK!
True or false: Lone pairs make a molecule polar. -If true, explain why. -If false, provide a counter-example. Answer: False, lone pairs do not always make a molecule polar. They might be arranged so that they are symmetrically distributed to minimize repulsions, such as XeF4.
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28. Intramolecular vs. Intermolecular Forces
Intramolecular forces: The forces within individual molecules holding it together
(ex: covalent and ionic bonds)
Intermolecular forces – weak interactions between molecules
Intermolecular forces can cause a condensed state of matter (liquids and solids).
Intermolecular forces are stronger in solids, weaker in liquids, and nearly absent in gases

29. Three types of Intermolecular forces (collectively called van der Waals forces)
London Dispersion
Dipole-Dipole
Hydrogen Bonding
Johannes van der Waals

30. London Dispersion
The weakest IM force
Present in all molecules; the only type of IM force present in non-polar substances and Nobel gases.
Caused by instantaneous dipoles
Random movement of electrons can create a momentary nonsymmetrical distribution of charge even in nonpolar molecules

31. London Dispersion
Instantaneous dipoles can induce a short-lived dipole in a neighboring molecule

32. Compare the boiling points of the noble gases:
London Dispersion
Dispersion force increases as the number of electrons in the molecule increases. (higher molecular weight, more electrons.) (Ex: CCl4 experiences greater London forces than CH4)
Concept Check: Which Nobel gas would you predict to have the lowest boiling point? Why?
Compare the boiling points of the noble gases:
helium °C experience fewer dispersion forces
neon °C
argon °C
krypton °C
xenon °C
radon °C experience more dispersion forces

33. Dipole-Dipole
Dipole-Dipole Forces
Attraction between molecules with dipole moments (molecules that have a permanent dipole)
Occur in addition to dispersion forces
Molecules orient themselves according to their poles
Maximizes (+,-) interactions
Minimizes (+,+) and (-,-) interactions

34. Strength increases as polarity increases.
Stronger than dispersion, but yet only 1% the strength of ionic bonds.

35. Hydrogen Bonding
Special type of Dipole-Dipole
The strongest of the Intermolecular forces.

36. Hydrogen is bound to a highly electronegative atom (F, O, N) with a lone pair
Important in bonding of molecules such as water and DNA
Ex: NH3, H2O, HF can hydrogen bond.
Ex: HCN does not, why?