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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) |

Created by:
mma129