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Published byAlisa McNally Modified over 2 years ago

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Gases are easily compressed Gases can expand Large amount of space between particles

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Compressing a gas causes pressure to increase 3 Factors Affect Gas Pressure Volume of the container (V) Measured in Liters Amount of gas (n) Measured in the number of moles Temperature Measured in Kelvins (T)

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Amount of gas affects on pressure Increase in number of particle Increases in number of collisions Increase in Pressure As the amount of gas changes, the pressure changes directly Ex. When the number of moles is doubled the pressure double. Also works when the moles decreases

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Gases will flow to the area with lower pressure. Ex: Deflating balloon Gas will leave the balloon into the surroundings Vacuum sealed container Gas will rush into the container

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Volumes effects on pressure Decrease in volume of the container Increase in the number of collisions Increase in pressure Volume inversely affects pressure A volume decrease causes a pressure increase If the volume is halved, the pressure is doubled. Squeezing a pack of Ketchup

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Temperatures affects on pressure Increase in Temperature Increase in the speed of the gas particles Increase in the number of collisions Increase in pressure Temperature directly affects pressure If the temperature double, so will the pressure Ex: Frozen balloon

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Boyles Law For a given mass of gas at constant temperature, the volume of a gas varies inversely with pressure.

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Boyles Law P ×V=P×V Pressure in units of kilopascals (kPa) Volume in units of liters (L)

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Boyles Law Ex: The volume of a raft has an initial volume of 1.2 liters and an initial pressure of 87kPa. If the final volume was 2.9 liters what was the final pressure ? P =87kPa V=1.2L P=? V=2.9 P ×V=P×V ( P ×V)/ V =P (87x1.2)/2.9= 36 kPa

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Charles Law: As the temperature of an enclosed gas increases, the volume increases, if the pressure is constant.

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Charles Law: V /T =V /T Volume in units of Liters (L) Temperature in units of Kelvins (K)

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Charles Law Ex: The volume of an inflated balloon at 24 ˚ C has a volume of 4 liters. The balloon is then moved to a room with a temperature of 58 ˚ C. What is the Volume? T = 24˚C+ 273 = 292 K T= 58˚C = 331 K V /T =V /T V = (VxT)/ T (4.00L x 331 K)/ 297 K= 4.46L

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Gay-Lussacs Law As the temperature of an enclosed gas increases, the pressure increases, if the volume is constant. P/T=P/T Pressure: kPa Temperature: K

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Gay-Lussacs Law Ex: An aerosol can is stored at 25˚C and has a pressure of 103 kPa. If the f=can is heated to 928˚C, what is the final pressure? T=25˚C + 273= 298 K T= 928˚C + 273= 1201 K P/T=P/T T(P/T)=P 103kPa x (1201 K / 298 K)= 415 kPa Or 4.15 x 10² kPa

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The Combined Gas Law Calculates for situations where the amount of gas is constant. (P x V)/T= (P x V)/ T P: kPa V: L T: K

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Combined Gas Law Ex: The volume of a balloon is 30.0 L at 313 K and 153 kPa pressure. What would the volume be at standard temperature and pressure (STP) STP: 273K and kPa (P x V)/T= (P x V)/ T V= (V x P x T)/ (P x T) (30L x 53kPa x 273K)/(101.3kPa x 313K)= 39.5 L

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Gas law

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Ideal Gas Law was created to calculate the number of moles of a contained gas. Symbol for number of moles is n Number of moles is directly proportional to volume

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1mol of every gas is 22.4 L at STP The Ideal Gas Constant R=(P x V)/(T x n) At STP : (101.3 x 22.4)/(273 x 1)= 8.31(L · kPa)/(K · mol) R= 8.31

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Common Equation PV= nRT P: kPa V: L n: mol R: (L · kPa)/(K · mol) T: K

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Converting mass to moles. 12 grams of CO how many moles? 12grams/(MMof C + 2 x the MMof O) 12/(12+2·16)=.2727 mols

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What is the pressure of 113g of Xenon gas at 187˚C, held in a 1.2L container? nRT/V=P 113g/131.3MM of Xe=.8606 mols T= 187˚C+ 273= 460k (.8606 ·8.31·460)/1.2= kPa

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Real Gases vs. Ideal Gases: Ideal gases follow the gas laws at all pressures and temperatures. Real gases can not be described by the gas laws at certain temps or pressures.

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An ideal gas has particles with NO volume, and there are no attractions between the particles.

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A real gas has particles that have volume, and there are interactions between particles. Real gases differ most from ideal gases at high pressures and low temperatures.

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Partial pressure: The pressure that one gas contributes to the total pressure. Daltons Law : P total = P1 + P2 + P3 + …

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101.3 kPa

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On top of Mt. Everest, air pressure is kPa. Since oxygen is 21% of air, the pressure of oxygen is 7 kPa. You need kPa Oxygen to live, so must Have compressed O 2.

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Diffusion: The tendency of molecules to spread out evenly. Effusion: A gas escapes through a tiny hole. Lower the molar mass of a gas = Faster Effusion and Diffusion.

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Grahams Law: RateA / Rate B = (Molar mass B / Molar mass A)^.5

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