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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 - PowerPoint PPT Presentation
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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.