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# MCAT Gen. Chem Ch. 8

Term | Definition |
---|---|

Gases Are The Least: | Dense phase of matter |

Gases Are Fluids, And Therefore: | They conform to the shapes of their containers |

Gases Are Easily: | Compressible |

Gas Systems Are Described By The Variables: | T, temp, P press., V volume, n number of moles |

1 atm = | 760 mmHg = 760 torr = 101.325 kPa |

Simple Mercury Barometer | Measures incident (atmospheric) pressure. As pressure increases, more mercury is forced into the column, increasing its height. As pressure decreases, mercury flows out of the column under its own weight, and decreases its height |

Standard Temperature And Pressure (STP) | 273 K (0C) and 1 atm |

Equimolar Amounts Of Two Gases Will Occupy The: | Same volume at the same temperature and pressure. At STP, one mole of an ideal gas occupies 22.4 L. |

Kinetic Molecular Theory | Attempts to explain the behavior of gas particles. It makes a number of assumptions about the gas particles. |

Assumptions Of The Kinetic Molecular Theory, Gas Particles: | Have negligible volume, do not have intermolecular attractions or repulsions, undergo random collisions with each other and the walls of the container |

Collisions Between Gas Particles Are: | Elastic |

Average Kinetic Energy Of Gas Particles Is: | Directly proportional to temperature. |

Diffusion | Spreading out of particles from high to low concentration |

Effusion | Movement of gas from one compartment to another through a small opening under pressure |

Real Gases Deviate From Idea Behavior Under: | High pressure (low volume) and low temperature conditions. |

At Moderately High Pressures, Low Volumes, Or Low Temperatures, Real Gases Will: | Occupy less volume than predicted by the ideal gas law because the particles have intermolecular attractions. |

At Extremely High Pressures, Low Volume, Or Low Temperatures, Real Gases Will | Occupy more volume than predicted by the ideal gas law because the particles occupy physical space. |

Eq. 8.1: Ideal Gas Law | PV = nRT. R = ideal gas constant 8.21 x 10^-2 L*atm / mol*K |

Eq. 8.2: Density Of A Gas | D = m/v = PM / RT. M = Molar Mass. R = ideal gas constant 8.21 x 10^-2 L*atm / mol*K. |

Eq. 8.3: Combined Gas Law | P1V1/T1 = P2V2/T2 |

Eq. 8.4: Avogadro's Principle | n/V = k or n1/V1 = n2/V2 |

Eq. 8.5: Boyle's Law | PV = k or P1V1 = P2V2. k = a constant |

Eq. 8.6: Charles's Law | |

Eq. 8.7: Guy-Lussac's Law | P/T = k OR P1/T1 = P2/T2 |

Eq. 8.8: Dalton's Law (Total Pressure From Partial Pressures) | PT = PA + PB + PC + ... |

Eq. 8.9: Dalton' Law (Partial Pressure From Total Pressure) | PA = XA*PT. XA = moles of gas A / total moles of gas |

Eq. 8.10: Henry's Law | [A] = kH * PA or [A]1/P1 = [A]2/P2 = kH. kH = Henry's constant that depends on the identity of the gas. |

Eq. 8.11: Average Kinetic Energy Of A Gas | KE = 1/2mv^2 = 3/2*ka*T. ka = Boltzmann constant, 1.38 x 10^-23 J/K. |

Eq. 8.12: Root-Mean-Square Speed | urms = Sq.root(3RT/M). R = ideal gas constant, 8.314 J/mol*K. T = temperature. M = molar mass. |

Eq. 8.13: Graham's Law | r1/r2 = Sq.root(M2/M1) |

Eq. 8.14: Van der Waals Equation Of State | (P + n^2*a / V^2) * (V-nb) = nRT. a = physical constant for attractive forces. b = physical constant for big particles. |