1. Unit 3: Chemical Bonding and Molecular Structure Cartoon courtesy of NearingZero.net

2. The Octet Rule Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level. Diatomic Fluorine

3. Hydrogen Chloride by the Octet Rule

4. Formation of Water by the Octet Rule

5. Comments About the Octet Rule 2nd row elements C, N, O, F observe the octet rule. 2nd row elements B and Be often have fewer than 8 electrons around themselves - they are very reactive. 3rd row and heavier elements CAN exceed the octet rule using empty valence d orbitals. When writing Lewis structures, satisfy octets first, then place electrons around elements having available d orbitals.

6. Multiple Covalent Bonds: Double bonds Two pairs of shared electrons

7. Multiple Covalent Bonds: Triple bonds Three pairs of shared electrons

8. Bond Dissociation Energy It is the energy required to break a bond. It gives us information about the strength of a bonding interaction.

9. Resonance Occurs when more than one valid Lewis structure can be written for a particular molecule. These are resonance structures. The actual structure is an average of the resonance structures.

10. Resonance in Ozone Neither structure is correct.

11. Hybridization The Blending of Orbitals

12. We have studied electron configuration notation and the sharing of electrons in the formation of covalent bonds. Methane is a simple natural gas. Its molecule has a carbon atom at the center with four hydrogen atoms covalently bonded around it. Lets look at a molecule of methane, CH 4.

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

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

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

16. 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. A Problem Arises

17. 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…? Unequal bond energy

18. 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. Unequal bond energy #2

19. 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? Unequal bond energy #3

20. The simple answer is, “No”. 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. Measurements show that all four bonds in methane are equal. Thus, we need a new explanation for the bonding in methane. Enter Hybridization

21. In the case of methane, they call the hybridization sp 3, 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. sp 3 Hybrid Orbitals

22. Here is another way to look at the sp 3 hybridization and energy profile… sp 3 Hybrid Orbitals

23. While sp 3 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. Notice that this produces two hybrid orbitals, while leaving two normal p orbitals sp Hybrid Orbitals

24. While sp 3 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. Notice that this produces two hybrid orbitals, while leaving two normal p orbitals sp Hybrid Orbitals

25. Another hybrid is the sp 2, which combines two orbitals from a p sublevel with one orbital from an s sublevel. Notice that one p orbital remains unchanged. sp 2 Hybrid Orbitals

26. VSEPR Model The structure around a given atom is determined principally by minimizing electron pair repulsions. (Valence Shell Electron Pair Repulsion)

27. Predicting a VSEPR Structure Draw Lewis structure. Put pairs as far apart as possible. Put pairs as far apart as possible. Determine positions of atoms from the way electron pairs are shared. Determine positions of atoms from the way electron pairs are shared. Determine the name of molecular structure from positions of the atoms. Determine the name of molecular structure from positions of the atoms.

28. VSEPR and the water molecule

29. VSEPR and the ammonia molecule

30. Polar-Covalent bonds Nonpolar-Covalent bonds Covalent Bonds  Electrons are unequally shared  Electronegativity difference between.3 and 1.7  Electrons are equally shared  Electronegativity difference of 0 to 0.3

31. Polarity A molecule, such as HF, that has a center of positive charge and a center of negative charge is said to be polar, or to have a dipole moment. A molecule, such as HF, that has a center of positive charge and a center of negative charge is said to be polar, or to have a dipole moment.

32. Relative magnitudes of forces The types of bonding forces vary in their strength as measured by average bond energy. Covalent bonds (400 kcal) Hydrogen bonding (12-16 kcal ) (Van der Waals Forces) Dipole-dipole interactions (2-0.5 kcal) London forces (less than 1 kcal) Strongest Weakest

33. Hydrogen Bonding Hydrogen bonding in Kevlar, a strong polymer used in bullet-proof vests. Bonding between hydrogen and more electronegative neighboring atoms such as oxygen and nitrogen

34. Hydrogen Bonding in Water

35. Hydrogen Bonding between Ammonia and Water

36. Dipole-Dipole Attractions Attraction between oppositely charged regions of neighboring molecules.

37. The water dipole

38. The ammonia dipole

39. London Dispersion Forces The temporary separations of charge that lead to the London force attractions are what attract one nonpolar molecule to its neighbors. Fritz London 1900-1954 London forces increase with the size of the molecules.

40. London Forces in Hydrocarbons