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| PV=nRT (R = universal gas constant, 0.0821) |
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Temperature: 273K
Pressure: 1 atm
V of one mole of gas: 22.4 L |
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| Equal volumes of gases at the same temperature and pressure contain the same number of molecules. |
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| Kinetic Molecular Theory #1 |
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| Gases consist of large numbers of molecules that are in continuous random motion. |
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| Kinetic Molecular Theory #2 |
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| The volume of all the molecules of gas is negligible compared to the total volume in which the gas is contained (because the molecules are so small) |
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| Kinetic Molecular Theory #3 |
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| Gas molecules don't attract or repel each other. |
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| Kinetic Molecular Theory #4 |
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| The collisions of gas molecules are elastic (they don't lose any energy or charge or anythann) |
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| Kinetic Molecular Theory #5 |
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| The average kinetic energy of the gas molecules is proportional to the absolute temperature, and at any given temperature, all molecules have the same average kinetic energy. |
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| What causes pressure of gas? |
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| The collisions of gas particles against walls of the container. The pressure is determined by how often and with what force the molecules hit. |
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| KEave=3/2(RT) (R=8.31) *Only need T |
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| Root mean square velocity |
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| uRMS=radical((3RT)/M) (R=8.31, M=molar mass in kg) |
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| The escape of gas molecules through a tiny hole into evacuated space. |
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| Spread of one substance through a space of a second substance. |
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| Graham's Law/rate of effusion |
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| R1/R2=radical(m2/m1) (m = molar mass) |
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| Gas molecules have an average speed, but there is a range of speeds for each individual molecule. As the temperature increases, there are less molecules with the average speed and a greater range of speeds. |
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| Average Kinetic Energy is the same, but particles have farther to travel before colliding with the container; less collisions; decreased pressure. |
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| Average Kinetic Energy increases, increase in urms; more collisions; increased P. |
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| Makes the P higher to adjust for attractive forced |
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| Makes the V lower to make up for the molecules volume...wait what? |
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