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Gas Laws
gas laws and stoichiometry
Term | Definition |
---|---|
Kinetic Molecular Theory pt. 1 | Gas molecules do not attract or repel each other |
Kinetic Molecular Theory pt. 2 | Gas particles are much smaller than the spaces between them |
Kinetic Molecular Theory pt. 3 | Gas particles are in constant, random motion |
Kinetic Molecular Theory pt. 4 | No kinetic energy is lost when gas particles collide with each other or with the walls of their container |
Kinetic Molecular Theory pt. 5 | All gases have the same kinetic energy at a given temperature |
Rate of effusion depends on | the mass of the particles |
In diffusion, lighter particles diffuse... | more rapidly than heavier particles |
Graham's Law of Effusion | Rate A / Rate B = sqrt(Rate B / Rate A) *compound A has the smallest molar mass |
Pressure formula | P= force / area |
A barometer is used the measure | atmospheric pressure |
Units of standard pressure | 101.3 kPa 1 atm 760 mmHg |
Dalton's Law of Partial Pressure | The total pressure of a mixture of gases is equal to the sum of the pressures of all the gases in the mixture. Ptotal = Px + Py + Pz |
Pressure is measured in these 3 units: | mmHg, kPa, atm |
Temperature is measure in | Kelvin (conversion from Celsius = C + 273) |
Volume is measured in | Liters |
Number of particles is solved by completing | gas stoichiometry |
Boyle's Law | P1V1 = P2V2 inverse relationship |
Charles' Law | V1 / T1 = V2 / T2 direct |
Gay-Lussac's Law | P1 / T1 = P2 /V2 direct |
Combined Gas Law | P1V1 / T1 = P2V2 / T2 direct |
Ideal Gas Law | PV = nRT p = pressure, v = volume, n = number of moles, R = given constant, T = temperature |
How many variations to the Ideal Gas Law? | 3 |
Variables | M = Molar mass m = mass D = density (g/L) |
Variations | M = mRT / PV M = DRT / P D = MP / RT |
If GRAMS are given, start with: | stoichiometry |
If LITERS are given, start with: | PV = nRT (or if @ STP, you may use the 22.4/mol conversion) |