I. Properties of gases

1. compressible and expandable
2. low density
3. 1 mole of gas at STP= 22.4L
A. STP= standard temp and pressure

II. Kinetic Molecular Theory - explains ideal gases

Ideal gases are theoretical
1. Volume of individual gas particles is zero compared to the total volume of gas*
2. Ideal gas is in constant, random motion, move in straight lines, until they collide with something.
3. There are no attractive or repulsive forces between particles. There for they are elastic.*
4. Energy is proportional to temperature
5. KE=1/2mv2

III. Four variables describe a gas

    1. volume (usually in L)
    2. temperature (in Kelvin: Celsius + 273)
    3. pressure (mmHg, atm, torr)
    4. 1. force/unit of area: force of the moving particles. Each collision exerts a force.
2. Measure
D. Number of moles
E. STP: P = 1 atm, 760 mmHg, 760 torr, 101.3 kPa

IV. Avogadro’s Principle: equal volumes of gases under the same conditions have equal numbers of molecules.. no matter what the gases

•. Always talk of gases in volume not mass volume of gas at STP is called molar volume
    1. What conditions cause a gas to condense
      1. Two ideal gas assumptions not true
    1. volume of gas is negligible
    2. no attractive forces between particles
2. If temp low and P is high: we see attractive forces
3. Vapor pressure increases temperature
    1. open container water vapor evaporates
    2. closed container evaporates until air in container is saturated with vapor (can’t hold any more)—then evaporates and condense equal each other
4. This is called equilibrium vapor pressure (changes due to temp)

B. Volatile substances have a high vapor pressure

1. At low temperatures, volatile substances still evaporate
2. Low molar mass (smell quickly) High molar mass (smell longer)

C. Phase Diagram relate temp and pressure to physical state

1. Sublimation is change from solid to gas
2. Triple point is where you find all three states in coexistence
*Attractive forces create surface tension.

Gas Laws

  1. Charles’ Law: V1/T1 = V2/T2
    1. As temperature rises, gas takes up more space.
    2. Must use Kelvin (273.16 + C)
    3. Pressure must be held constant
  1. Boyle’s Law: P1V1 = P2V2
    1. As volume decreases, pressure increases
    2. Temp. must be held constant
  1. Dalton’s law of Partial pressure
    1. the total pressure in a gas mixture is the sum of the partial pressures of the individual gases.
    2. Partial pressure is the pressure of the individual gas that makes up the total pressure.
    3. Ptotal = Pa + Pb + Pc
    4. Mole fractions used to find partial pressures
      a. mole fraction of gas A = moles of gas A/ total number of moles of gas
      b. then take that fraction and multiply it by total pressure to get partial P

Gas Diffusion

I. gas molecules disperse from regions of higher concentration to lower concentration.
II. Light molecules travel faster than heavier ones
III. Graham’s law of diffusion
a. it is related to the their molecular masses refer to lab

Gases

    1. Gases diffuse to fill their containers and the rate relates to molar mass
    2. Graham’s law: Va/Vb = Mb/Ma

Fitting the Gas Laws together

    1. PV=nRT; P=pressure; V=volume; n=moles; R=rate constant; T=temperature
    2. R=8.314 L-Pa/mol0K or .0821Loatm/moloK
    3. there will be one variable to solve for
    4. you can also use a general gas law P1V1/T1 = P2V2/T2

VI. Use the general gas law if the system is changing. I.E. they give two T, P, or V.

Example: A 500g block of dry ice (solid CO2) vaporizes to a gas at room temperature. Calculate the volume of gas produced at 25 degree Celsius and 975 kPa.

Determine what you are solving for. Decide which R you are using, and then plug in the variables into the equation.