By Robert Liu and Julie Baldassano

Chapter 8: Bonding-General Concepts
Ionic bonds: bonds between two ions, usually with a very large electronegativity difference. Example: CaCl2
Covalent bonds: bonds in which electrons are shared between the two nuclei. Example: H2
Polar Covalent bonds: bonds in which electrons are shared, but unequally. Example: H2O

Chapter 8: Electronegativity
The ability of an atom in a molecule to attract shared electrons to itself. Generally increases across a period and decreases going down a group

Chapter 8: Bond Polarity
Molecules with a center of negative charge and a center of positive charge are known as dipolar.
Any molecule with a polar bond will have dipole moments, making the entire molecule polar. However, symmetrical arrangements will cancel out dipole moments.

Chapter 8: Ions
Ionic compounds are formed from a nonmetal ion and a metallic ion. (MgCl2, Cacl2)
Cations are always smaller in radius than their parent atoms. Anions, always bigger.
With isoelectronic ions(those with the same amount of valence electrons), ionic radius decreases as nuclear charge increases.

Chapter 8: Lattice Energy
As charge increases, the lattice energy increases. Mg2+>Na+.
As distance between the two nuclei in an ion decrease, lattice energy increases.

Chapter 8: Ionic Character
Electronegativity and % ionic character have a positive relationship.

Chapter 8: Bond Energy
Simple Arithmetic. Take H2 +F2=2HF for example.
Bonds broken-bonds formed=(triangle)H

Chapter 8: The LEBM
1. Description of valence electron arrangement using lewis dot structures.
2. Predicting geometry using VSEPR
3. Description of the type of atomic orbitals used to share electrons

Chapter 8: Lewis Structures
Use valence electrons only.
Hydrogen obeys duet rule, everything else typically obeys octet rule.
One dot= one electron, one line=a pair of electrons.

Chapter 8: Exceptions
Boron tends to form compounds in which it is electron deficient.
Third period elements can exceed the octet rule using their open d orbitals
(citation needed)
Odd numbers of electrons (NO), oxygen has 6, nitrogen has 5. 6+5=11 o_O? Results in 5 around nitrogen, 8 around oxygen.

Chapter 8: Resonance
Occurs when more than one lewis structure can be drawn for a molecule.

Chapter 8: VSEPR
Minimize repulsions by spacing atoms as far away as possible from each other.
Bond angles: Linear=180
Trigonal Planar=120
Tetrahedral=109.5
Trigonal Pyramidal=<120 ~107
Bent=<120 ~104.5

Chapter 8 - Quiz
1) The geometry of the SO
3 molecule is best described as
(A) trigonal planar
(B) trigonal pyramidal
(C) square pyramidal
(D) bent
(E) tetrahedral
2) For which of the following molecules are resonance structures necessary to describe the bonding satisfactorily?
(A) H
2
S
(B) SO
2
(C) CO
2
(D) OF
2
(E) PF
3
3) The electron-dot structure (Lewis structure) for which of the following molecules would have two unshared pairs of electrons on the central atom?
(A) H
2
S
(B) NH
3
(C) CH
4
(D) HCN
(E) CO
2
1) A
2) B
3) A

Chapter 9 - Hybridization
Hybridization: the mixing of native atomic orbitals to form special orbitals for bonding
Molecules using localized electron model:
1) Draw Lewis structure
2) Place electron pairs using VSEPR model
3) Match effective electron pairs to hybrid orbitals

Chapter 9 - sp sp:
• linear shape
•
2 effective pairs
•
180 degree angles

Chapter 9 - sp 2 sp 2 :
• trigonal planar shape
•
3 effective pairs
•
120 degree angles

Chapter 9 - sp 3 sp 3 :
• tetrahedral shape
•
4 effective pairs
•
109.5 degree angles

Chapter 9 - dsp 3 dsp 3 :
• trigonal bipyramidal shape
•
5 effective pairs
•
90 and 120 degree angles

Chapter 9 - d 2 sp 3 d 2 sp 3 :
• octahedral shape
•
6 effective pairs
•
90 degree angles

Chapter 9 σ and π bonds
σ: electrons shared in area on a line between the atoms, localized bonding
π: a shared electron pair occupies the space above and below the line between the atoms, delocalized bonding

Chapter 9 - MO Model
Molecular orbital model:
-Pros: no resonance, effective for molecules with unpaired electrons, bond energy information
Antibonding (MO2) are higher energy than bonding (MO1) orbitals
Bond order = # of bonding electrons - # of antibonding electrons

Chapter 9 - MO Model
- Start at lowest energy level
- Each orbital can hold two electrons
(each with different spins)
- Use number of valence electrons to fill
- Degenerate = B2, C2, N2 so pi2p is lower energy than sigma2p, flipping the lines

Chapter 9 - Magnetism & Trend
Paramagnetic: attracted to magnetic field
- unpaired electrons
Diamagnetism: repelled from magnetic field
-paired electrons
- High bond order = high bond energy = short bond length

Chapter 9 Quiz
1) SH
2
•
•
•
•
•
3) NH
•
2
1-
•
•
•
• sp sp 2 sp 3 sp 3 d sp 3 d 2 sp sp 2 sp 3 sp 3 d sp 3 d 2
2) HCN
•
•
•
•
•
4) SF
4
•
•
•
•
• sp sp 2 sp 3 sp 3 d sp 3 d 2 sp sp 2 sp 3 sp 3 d sp 3 d 2
5) Using molecular orbital theory, determine the magnetism of O
2
O
2
− .
and aO
2 is paramagnetic; O
2
− is diamagnetic.
bO
2 is diamagnetic; O
2
− is paramagnetic.
cBoth O
2 and O
2
− are diamagnetic.
dBoth O
2 and O
2
− are paramagnetic.

Chapter 9 Quiz Answers
1) SH
2 is sp 3
2) HCN is sp
3) NH
2
1is sp
4) SF
4 is sp 3 d
5) d- Both O
2
3 and O
2
− are paramagnetic.