CHEM&121 Exam 3: Chapter 7

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tgherasim
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297117
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CHEM&121 Exam 3: Chapter 7
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2015-02-28 17:55:07
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chemistry
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chapter 6
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  1. Properties of Gases
    • A gas consists of particles - atoms/molecules - that move randomly & rapidly.
    • The size of the particles are small compared to the space between them.
    • Because the space is large, there are no attractive forces between them.
    • The kinetic energy increases with increasing temp.
    • When gas particle collide, they rebound. When they hit a wall, they exert pressure.
  2. Gas Pressure
    • Pressure (P) is the force (F) exerted per unit area (A). 
    • 1 atm = 760 mmHg = 760 torr = 14.7 psi = 101,325 Pa
  3. Boyle's Law
    • Pressure and Volume
    • For a fixed amount of gas at constant temperature, the pressure and volume of a gas are INVERSELY related (when one increases, the other decreases).
    • The same number of gas particles occupies 1/2 the volume and exerts 2X the pressure.
    • Equation: P1V1=P2V2
  4. Charles's Law
    • Volume and Temperature
    • For a fixed amount of gas at constant pressure, the volume of a gas is PROPORTIONAL to its Kelvin temp (if one increases, the other will also).
    • Equation for C to K: C + 273
    • Equation
  5. Gay-Lussac's Law
    • Pressure and Temperature
    • For a fixed amount of gas at constant volume, the pressure of a gas is PROPORTIONAL to its Kelvin temp (if one increases, the other will also).
    • Equation
  6. Combined Gas Law
    • Shows the relationship of Pressure, Volume, and Temperature for a constant number of moles.
    • Equation
  7. Avogadro's Law
    • Volume and Moles
    • When the pressure and temperature are held constant, the volume of a gas is PROPORTIONAL to the number of moles present (when one increases, the other will also).
    • Equation: 
  8. STP Conditions
    • Pressure: 1 atm (760 mmHg)
    • Temp: 273 K (0 degrees C)
    • 1 mole of any gas has the same volume as 22.4 L (Standard Molar Volume).
  9. Ideal Gas Law
    • Pressure, Volume, Temp, Universal gas Constant (R), and Moles
    • Equation: PV=nRT
    • R=0.0821
  10. Dalton's Law & Partial Pressures
    • The total pressure (Ptotal) of a gas mixture is the sum of the partial pressures of its component gas.
    • Equation for Total: Ptotal = A + B + C
    • Equation for Partial: % of Gas = Decimal, THEN Decimal x Ptotal = Partial Pressure
  11. London Dispersion Forces
    • Weakest of all the intermoleculars due to momentary changes in electron density in a molecule.
    • Includes all molecules and atoms
    • The larger the molecule, the larger the attractive force between two molecules and the stronger the intermolecular force.
    • Examples: CH4, H2CO, H2O
  12. Dipole-Dipole Force
    • Strongest after Dispersion.
    • Attractive forces between TWO POLAR molecules.
    • Examples: H2CO, H2O
  13. Hydrogen Bonding
    • Strongest after Dipole-Dipole
    • Molecules containing H bonded to F, O, or N
    • Examples: H2O
  14. Ion-Dipole Force
    • Strongest out of all the intermolecular forces.
    • Mixtures of ionic compounds and polar compounds.
  15. Boiling Point
    • The temperature at which a liquid is converted to a gas phase. 
    • The increase in strength of the intermolecular forces, the increase in boiling points. 
    • Endothermic: energy is absorbed
    • I.e. Hydrogen Bonding (H2O) would have a HIGHER boiling point than London Dispersion (CH4)
  16. Melting Point
    • The temperature at which a solid is converted to a liquid phase.
    • The increase in strength of the intermolecular forces, the increase in melting points.
    • Endothermic: energy is absorbed.

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