1. Types of bonding, what type of element they're between, and how e- react
    • ionic (metals to nonmetals) - e- are transferred
    • covalent (nonmetals to nonmetals) - e- are shared
    • metallic (metal to metal) - e- are pooled
  2. Bond type by electronegativitiy
    • 0.0 = pure covalent
    • 0.1 - 0.4 = nonpolar covalent
    • 0.5 - 1.9 = polar covalent
    • 2.0 - 4.0 = ionic
  3. Octet rule exceptions
    • H, He have only 2 e-
    • Group IIIA elements may have 6 e- only
    • Elements in period 3+ may have 8, 10, 12, or 14 e-
  4. The only 3 free radical compounds (exceptions to lone pair rule)
    NO, NO2, and ClO2
  5. How do you check a Lewis Dot Diagram?
    • Correct # of e-?
    • Octet rule?
    • Σ Formal Charges = charge of molecule?
  6. Trends in bond length
    • More e- shared by atoms = shorter covalent bond
    • Bond length decreases from left to right across period
    • Bond length increases down the column
  7. Trends in bond energy (energy needed to break a bond)
    • More e- shared by atoms = stronger covalent bond
    • Shorter covalent bond = stronger covalent bond
  8. Using bond energies to estimate ΔHrxn
    ΔHrxn = ΣnΔH(bonds broken) - ΣnΔH(bonds formed)
  9. Evaluating resonance structures
    • Better structures have fewer formal charges
    • Better structures have smaller formal charges
    • Better structures have negative formal charges on more electronegative atoms
  10. Possible shapes of molecules with VSEPR formula and bond angle
    • AX2 - linear - 180˚
    • AX3 - trigonal planar - 120˚
    • AX2E - v shaped/bent
    • AX4 - tetrahedral - 109.5˚
    • AX3E - trigonal pyramidal
    • AX2E2 - v shaped/bent
    • AX5 - trigonal bipyramidal - 120˚ (equatorial) and 90˚ (axial to equatorial)
    • AX4E - irregular tetrahedral / seesaw
    • AX3E2 - T shaped
    • AX2E3 - linear
    • AX6 - octahedral - 90˚
    • AX5E - square pyramidal
    • AX4E2 - square planar
  11. How do lone pairs affect bond angle?
    Lone pair "take up more space" and decrease the bond angle between atoms
  12. How to draw 3D object on 2D plane
    • solid line - on the same plane as the paper
    • solid bar - in front of plane
    • dashed bar - behind plane
  13. Which shapes result in non-polar molecules (through vector cancellation) and which shapes result in polar molecules (uncancelled vectors)?
    • Nonpolar: linear, trigonal planar, tetrahedral, tigonal bipyramidal, octahedral, square planar
    • Polar: bent, trigonal pyramidal, seesaw, t-shaped, square pyramidal
  14. How does Valence bond theory explain bonding? List the hybrids, and their corelation to Lewis dot.
    • VBR states that a bond is the overlap of atomic (or hybrid) orbitals.
    • Hybrids are created based on the lewis dot structure, based on how many e- densitites the atom has.
    • 2 (sp), 3 (sp2), 4 (sp3), 5 (sp3d), 6 (sp3d2)
  15. How do you know how MANY hybrid orbitals to use?
    • The number of atomic orbitals combined = the number of hybrids formed
    • eg combining a 2s with a 2p gives 2 sp orbitals
  16. Describe the different types of bonds using greek letters, and how each overlaps in depth.
    • 1 single bond = 1 δ bond
    • 1 double bond = 1 δ bond and 1 π bond
    • 1 triple bond = 1 δ bond and 2 π bonds
    • δ bonds overlap once, along the axis of the bond using hybrid orbitals
    • π bonds overlap twice, perpendicular to axis using unhybridized p orbitals
  17. Explain the steps in proving a hybridization
    • Draw the lewis dot structure
    • Get the electron configuration for the element
    • Establish hybridization based on lewis dot
    • Draw energy levels using electron configuration
    • Create hybrids and fill in the electrons, make sure it matches (remember p orbitals are unhybridized in π bonds
  18. How does Molecular orbit theory explain bonding?
    Electrons belong to whole molecule, orbitals belong to whole molecule (delocalization)
  19. Differences between VBT, MO, and Lewis
    • VBT predicts many properties better than Lewis (bonding schemes, bond length, bond strengths, bond rigidity)
    • VB presumes electrons are localized, and does not account for delocalization
    • VB cannot predict perfectly (magnetic behavior)
    • MO can predict bond order, energies, magnetic properties
    • Both are used, but have strengths and weaknesses
  20. What forms a bonding molecular orbital? What are the symbols?
    • When the two wave functions combine constructively the resulting molecular orbital has less energy than the original atomic orbital
    • δ and π are bonding orbitals (most electon density between nuclei)
  21. What forms a antibonding molecular orbital? What are the symbols?
    • When the two wave functions combine deconstructively the resulting molecular orbital has more energy than the original atomic orbitals
    • δ* and π* are antibonding orbitals (most electon density outside nuclei)
    • nodes (spaces without electrons) between nuclei
  22. What is bond order?
    • (Bonding electrons - antibonding electrons) / 2
    • Only use valence electrons
    • higher bond order = stronger/shorter bonds
    • fractions possible
  23. MO paramagnetic vs diamagnetic
    • paramagnetic (attracted to magnets) if MO diagram has unpared electrons
    • diamagnetic (not attracted to magnets) if MO diagram has all electrons paired
  24. LUMO, HOMO, and what they are used for.
    • Lowest Unpaired Molecular Orbit
    • Highest Occupied Molecular Orbit
    • Difference is used to determine wavelength absorpotion by molecule
  25. Name of structure that ions form?
    Crystal Lattice
  26. What is lattice energy (ΔHlattice)? Formulaic definition? What formula/factors affect lattice energy?
    • Energy released when 1 mol solid crystal forms from ions in gas state
    • ΔHlattice = cation(g) + anion(g) -> 1 mol molecule(s)
    • Always exothermic
    • Depends on size/charge of ions [direct] and distance between ions [inverse]
    • E = C x q1q2/r
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