Chapters 8 & 9 Review

Chapters 8 & 9 Review

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=180Trigonal Planar=120Tetrahedral=109.5Trigonal Pyramidal=<120 ~107Bent=<120 ~104.5

Chapter 8 - Quiz1) The geometry of the SO3 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) H2S

(B) SO2

(C) CO2

(D) OF2

(E) PF3

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) H2S(B) NH3

(C) CH4

(D) HCN(E) CO2

1) A2) B3) A

Chapter 9 - Hybridization

Hybridization: the mixing of native atomic orbitals to form special orbitals for bonding

Molecules using localizedelectron model:1) Draw Lewis structure2) Place electron pairs using VSEPR model3) Match effective electron pairs to hybrid orbitals

Chapter 9 - sp

sp:

• linear shape

• 2 effective pairs

• 180 degree angles

Chapter 9 - sp2

sp2:

• trigonal planar shape

• 3 effective pairs

• 120 degree angles

Chapter 9 - sp3

sp3:

• tetrahedral shape

• 4 effective pairs

• 109.5 degree angles

Chapter 9 - dsp3

dsp3:

• trigonal bipyramidal shape

• 5 effective pairs

• 90 and 120 degree angles

Chapter 9 - d2sp3

d2sp3:

• 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

2

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 Quiz5) Using molecular orbital theory,

determine the magnetism of O2 and O2

−.

a-O2 is paramagnetic; O2− is diamagnetic.

b-O2 is diamagnetic; O2− is paramagnetic.

c-Both O2 and O2− are diamagnetic.

d-Both O2 and O2− are paramagnetic.

2) HCN

• sp

• sp2

• sp3

• sp3d

• sp3d2

4) SF4

• sp

• sp2

• sp3

• sp3d

• sp3d2

1) SH2

• sp

• sp2

• sp3

• sp3d

• sp3d2

3) NH21-

• sp

• sp2

• sp3

• sp3d

• sp3d2

Chapter 9 Quiz Answers

1) SH2 is sp3

2) HCN is sp3) NH2

1- is sp3

4) SF4 is sp3d

5) d- Both O2 and O2− are paramagnetic.