Chapters 9 and 11 review
Chapter 9 Topics list:
Chemical Bond Theories
-Valence Bond Theory (uses Lewis Structures)
Bonds form using shared electrons between overlapping orbitals on adjacent atoms.
Orbitals arrange around the central atom to avoid each other.
-Molecular Orbital Theory: Uses MO Diagrams
Orbitals on atoms “mix” to make molecular orbitals, which go over 2 or more atoms.
Two electrons can be in an orbital.
Orbitals are either: bonding, antibonding, or nonbonding
-Bonds are either sigma or pi
-Orbitals of bonding atoms overlap directly between bonding atoms (sigma bonds)
-Atomic Orbitals change shape when they make a molecule
Hybrid Orbitals
-Atomic valence orbitals combine to form hybrid orbitals
-Hybrid orbitals go in the VSEPR electron geometry direction
-Formula-->Lewis Structure→#Structural Pairs-->Hybrids
Sigma vs. Pi Bonding
-hybrid orbitals
-H 1s orbitals terminal atom p orbitals
-pi bonding involves unhybridized p orbitals
Relationship Between Hybridization and Number of Possible Pi bonds
# Structural Pairs on the central atom | Hybridiztion | Unhybridized p Orbitals | Number of Possible p-p Pi Bonds |
2 | sp | two p | 2 |
3 | sp2 | one p | 1 |
4 | sp3 | None | 0 |
5 | sp3d | None | 0 |
6 | sp3d2 | None | 0 |
-Isomers: Molecules with same formula but different structure
-Conformers: Different temporary shapes of the same molecule
*Bond Rotations CAN happen around single bonds NOT double bonds
Chemical Bonding Theories
-Valence Bond Theory (uses Lewis Structures)
Bonds form using shared electrons between overlapping orbitals on adjacent atoms.
Orbitals arrange around the central atom to avoid each other.
-Molecular Orbital Theory: Uses MO Diagrams
Orbitals on atoms “mix” to make molecular orbitals, which go over 2 or more atoms.
Two electrons can be in an orbital.
Orbitals are either: bonding, antibonding, or nonbonding
-Bonding Orbitals: Electrons in these orbitals help hold atoms near each other
-Antibonding Orbitals: Electrons in these orbitals push atoms apart from each other
-Nonbonding Orbitals: Electron in these have no effect on bonding
Chapter 11 Sections 11.1-11.3
Phase changes
Fusion or melting: solid to liquid and energy is absorbed
Vaporization: liquid to gas and energy is absorbed
Sublimation: solid to gas and energy is absorbed
Freezing: liquid to solid and energy is released
Condensation: gas to liquid and energy is released
Deposition: gas to solid and energy is released
Properties of liquids
Enthalpy- energy required to vaporize a liquid.
Vapor pressure- the gas pressure exerted by a vaporizing liquid.
Boiling point- Temperature at which vapor pressure is greater than or equal to the external atmospheric pressure.
Normal boiling point- Temperature required for a liquid’s vapor pressure to reach 1 atm. (760 mm Hg ).
Surface tension- A liquid’s surface to naturally resist change.
Viscosity- A measure of a liquid’s resistance to flow.
Dynamic equilibrium- Rate of going from liquid to gas and back to liquid is equal.
Volatility- Tendency for liquids to turn to gases. (Weak IMF=more volatile).
Note: Long molecules flow poorly. They have a high viscosity.
Intermolecular Forces
Intermolecular forces dictate whether a molecule holds together well or not.
Examples of molecules that do hold together well include:
-Water -Salt
-Steal -Diamond
Note: Solids are better at holding together than any other phase.
Gases are poor at holding together.
-Molecules with greater intermolecular forces have larger enthalpy of vaporization.
-Greater IMFs lead to lower vapor pressure.
Equations to know:
R = Gas constant (8.314 J/KxMol)
T = Temp. in Kelvin
lnP = (-deltaH of vaporization/RT) + C
Straight line version: lnP = (-deltaH of vaporization/R)(1/T)+C
Two point version: ln(P2/P1) = (-deltaH of vaporization/R)((1/T2)-(1/T1))
Surface Tension
Cohesive forces- Forces between same molecules.
Adhesive forces- Forces between different molecules.
Types of Intermolecular Forces
Dipole-Dipole- Two polar molecules are attracted to each others opposite charges.
Induced Dipole-Induced Dipole (Dispersion)(Vanderwaals)- Electrons from one molecule repel electrons from the other molecule causing the molecule to become partially positive, allowing the other molecule to bond to it.
The tendency to do this is called polarizability. The trend on the periodic table for this is down and to the left.
Dipole-Induced Dipole- Similar to the previous type, a polar bond comes in contact with a non-polar bond and causes it to become partially charged.
Strength of Dipole forces increases as electronegativity difference between atoms in a molecule and also a lack of symmetry.
Strength of Induced-Dipole forces increases as size of the molecule increases.
H-Bonding > Dipole-Dipole > Induced-Dipole
H-Bonds are so strong because:
