1. AP Chem
Turn in “Naming Practice WS” if you did not turn it in to the sub on Friday.
Pull up “Phet Molecule Shapes” on your laptops and complete the firs 2 pages.
Unit 1 Test Thurs 10/12

2. Why the disparity of bond angles?
Lone pairs repel electrons in bonds, pushing them away

3. Forming Bonds
Bond= a region that forms when electrons from different atoms interact with each other
The attraction between 2 or more atoms allows for the formation of a compound.
Only valence electrons participate in bonding
Octet Rule
Atoms bond in order to get a full valence shell, or 8 valence electrons
Exceptions:
Hydrogen is happy with 2 valence electrons
Boron is happy with 6 valence electrons

4. Covalent Bonds
A covalent compound is one that is made up of 2 or more nonmetals
A covalent bond involves the sharing of electrons
Lewis Structures
Show just the valence electrons

5. Remember: The group (vertical column) the element is in gives you information about the number of valence electrons it has!

6. “HONC Rule”
Bonds each atom needs to form in order to achieve a full valence shell:
Hydrogen: 1 bond
Oxygen: 2 bonds
Nitrogen: 3 bonds
Carbon: 4 bonds
Halogens: 1 bond
1 pair of electrons shared between atoms = single bond
2 pairs of electrons shared between atoms = double bond
3 pairs of electrons shared between atoms = triple bond

7. Exceptions to the Octet Rule: Expanded Octet
Central atom ends up with more than an octet. Occurs when there are too many electrons to fit in.
Central atom with an expanded octet MUST have an atomic number larger than 10 (beyond neon). In the discussion of hybridization, we see that when there are 5 or 6 “groups” on the central atom, d-orbitals are involved.
Extra electrons should be first placed on the outside atoms.
After the outside atoms have fulfilled the Octet Rule, and there are still extra electrons, start with placing them as lone pairs on the central atom. If the central atom has a positive charge should you move a lone pair from the outside atoms to share.

8. Hybrid Orbitals
To explain molecular geometries, we can assume that the atomic orbitals on an atom (usually the central atom) mix to form new orbitals called hybrid orbitals.

9. Ex: Hybrid Orbitals- Boron
The 2s orbital and two of the 2p orbitals are hybridized together to generate three new hybrid orbitals called sp2 orbitals that can each overlap with the orbitals of another atom and form covalent bonds.

10. Ex: Hybrid Orbitals—Carbon
The 2s orbital and the 2p orbitals are hybridized together to generate four new hybrid orbitals called sp3 orbitals that can each overlap with the orbitals of another atom and form covalent bonds.

11. Ex: Hybrid Orbitals—Be
Orbital diagram for a ground state beryllium atom:
The Be atom in its ground state cannot bond with the fluorine atoms because it has no unpaired electrons. The Be atom could form two bonds by promoting one of the 2s electrons to a 2p orbital:
The 2s orbital and the 2p orbital are hybridized together to generate two new hybrid orbitals called sp orbitals. These sp hybrid orbitals can each overlap with the orbitals of another atom to form covalent bonds.

12. Hybridization
Hybridization: combining/mixing of electron orbitals when atoms bond.
The hybridization is determined by the # things around the central atom
The shape is determined by the lone pairs on the central atom

13. Common Hybridizations
# Sets
Orbital Hybridization 2 sp 3
sp2 4
sp3 5
sp3d 6
sp3d2
Note: “sets” refer to bonded atoms as well as lone pairs on the central atom

14. VSEPR Theory Valence Shell Electron Pair Repulsion A=Central Atom
X=Bonded Atoms
E=electron pairs (non-bonding) on the central atom

16. Larger Molecules VSEPR can be applied to larger molecules 4 3 4
Trigonal planar
tetrahedral
bent
109.5
120
< 109.5

17. Bond Types Sigma Bond (σ) a bond formed between two elements
orbitals come together head on

18. Bond Types Pi Bond (π) a bond formed above and below elements
orbitals come together side by side

19. Bond Types
Single bonds
composed of one σ bond

20. Bond Types
Double bonds
composed of one σ bond and one π bond

21. Bond Types
Triple Bond
formed by one σ bond and two π bonds

22. Energy Associated with Bonding
Atoms with a full valence shell are stable (“happy”) and are in a low energy state.
Breaking Bonds (NOT spontaneous): Energy is absorbed
Go from low energy (“happy” atoms) to high energy (“unhappy” atoms)
Ripping two atoms apart requires energy= is endothermic

23. Energy Associated with Bonding
SPONTANEOUS Bond formation = Energy is released
Go from high energy (“unhappy” atoms) to low energy (“happy” atoms)
Forming a bond is exothermic

24. Calculating Bond Energy

25. Covalent Bonds-Formation
Atoms bonded together exist at the lowest energy possible

26. Covalent Bonds-Formation
Bond Energy
The energy needed to form/break a bond
Dependent upon two atoms in bond as well as other atoms within the molecule, therefore average is used

27. Covalent Bonds-Formation
Bond length
Average distance between two bonded atoms to give the lowest energy possible
The longer the bond length the weaker the energy
The more pi bonds the stronger the bond (more energy)
The stronger the bonds between atoms (ie single bond vs double bond) the shorter the bond length

28. Examples
Draw the Lewis structures for C2H6 and C2H2. Indicate which has the longest C-C bond and which has the strongest C-C bond.
1. the average C–H bond energy is 413 kJ/mol
considering that there are 4 C-H bonds, the total energy of CH4 molecule is 4*413 =1652 kJ/mol
The average O-H bond energy is 467 kJ/mol
there are 2 O-H bonds in H2O, the total energy is 2*467 = 934 kJ/mol