Chem1AFinal

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victimsofadown
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Chem1AFinal
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2011-05-22 22:20:35
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Chem1AFinal
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  1. λ, ν, c, h
    • λ = Wavelength - distance between successive crests (cm, m, nm)
    • ν = Frequency - number of cycles in a given amount of time (1/s = Hz)
    • c = Speed of Light - 3.00x108m/s
    • h = Planck's constant - 6.626x10-34J·s
  2. Coulombs Law
    • E = (constant) x q1q2/r
    • q = charge of ions
    • r = size of atom
  3. Vaporization Energy
    • ΔHvap
    • Amount of energy required to separate a molecule from the liquid (boil)
    • ΔHvap increase = stronger intermolecular forces
    • Larger surface area = faster rate of evaporation
  4. Vapor pressure
    • Pressure in a closed container from molecules that have vaprorized from liquid within (they can't escape, so they create pressure)
    • Higher vapor pressure = weaker intermolecular forces (more easily vaporized = more molecules in air = more vapor pressure)
  5. Viscosity
    • Resistance of liquid to flow
    • 1 poise = 1P = 1g/cm x seconds
    • H2O is 1 cP at room temperature
    • Larger viscosity = larger intermolecular forces
    • more spherical molecular shape = decreased viscosity (less surface to surface contact)
    • increased temperature = decreased viscosity (increase to average kinetic energy, easier to overcome flow)
  6. Phase diagrams
    • Regions represent states, lines represent state changes
    • Critical point: the furthest point on a vapor pressure curve
    • Triple point: temperature/pressure where all 3 states exist simulatenously
    • Normal BP, FP, etc = pressure @ 1 atm.
  7. Supercritical fluid
    • Substances past the critical point on a phase diagram
    • State that has some liquid properties and some gas properties
    • Made from sealed liquid being heated, increased pressure, increased vapor density, decreased liquid density... essentially they all merge.
  8. Henry's Law
    • Sgas=KHPgas
    • S = solubility, KH = Henry's Law constant (different for each gas), P = pressure
    • As pressure increases, gas solubility increases
    • As pressure decreases, gas solubility decreases
  9. Raoult's Law
    • Psolvent in sol=Xsolvent x P0
    • P = pressure, X = mole fraction, P0 = pressure of pure solvent
    • Vapor pressure calculation
  10. Boiling point elevation
    • ΔTb = m x Kb x i
    • m = molality, Kb = boiling point constant (°C/m),
    • i = van't hoff factor
    • B.p. sol higher than B.p. pure solvent
  11. Osmotic pressure
    • Π = MRTi
    • M = molarity, R = .0821 atm L / mol K, T = temp in K, i = van't hoff factor
    • amount of pressure needed to prevent osmotic flow
  12. van't hoff factor
    • i
    • amount of dissociation that occurs in solution
  13. 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
  14. 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
  15. 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-
  16. The only 3 free radical compounds (exceptions to lone pair rule)
    NO, NO2, and ClO2
  17. How do you check a Lewis Dot Diagram?
    • Correct # of e-?
    • Octet rule?
    • Σ Formal Charges = charge of molecule?
  18. 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
  19. Trends in bond energy (energy needed to break a bond)
    • More e- shared by atoms = stronger covalent bond
    • Shorter covalent bond = stronger covalent bond
  20. 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
  21. 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
  22. How do lone pairs affect bond angle?
    Lone pair "take up more space" and decrease the bond angle between atoms
  23. 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
  24. 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)
  25. 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
  26. 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
  27. 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
  28. How does Molecular orbit theory explain bonding?
    Electrons belong to whole molecule, orbitals belong to whole molecule (delocalization)
  29. 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
  30. 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)
  31. 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
  32. What is bond order?
    • (Bonding electrons - antibonding electrons) / 2
    • Only use valence electrons
    • higher bond order = stronger/shorter bonds
    • fractions possible
  33. 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
  34. 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
  35. 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
  36. E, W, q, C, Cs, ΔH,
    • E = Energy - anything that has the capacity to do work (Quantity that an object can posess)
    • W = Work - a force acting over a distance (way an onject can exchange E w/ other objects [in and out])
    • q = Heat - flow of energy caused by a difference in temperature (way an onject can exchange E w/ other objects [in and out])
    • C = Heat capacity - amount of energy required to change an objects temperature by 1 degree C
    • Cs = Specific heat - amount of energy required to raise the temperature of 1g substance by 1 °C
    • ΔH = total energy of a system (internal E + work) / determines endo/exothermic [ΔH = q (@ constant P)]
  37. Bomb calorimeter details vs Coffee cup calorimeter
    • Bomb - Constant volume = -CcalorimeterxΔT
    • Coffee - Constant pressure = -Cs x Masss x ΔTs
  38. Standard Conditions (thermodynamics) + symbol
    • symbol = °
    • gas = 1 atm pressure
    • liquid/solid = 1 atm pressure, 25°C
    • solution = 1M concentration
  39. Standard Enthalpy of formation (symbol + meaning)?
    • ΔHf° - enthalpy change for the reaction forming exactly one mole of a pure compound from its constituent elements at standard conditions
    • ΔHf° of elements is always - KJ/mol

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