C142A Ch 5 2011 lecture notes

C142A Ch 5 2011 lecture notes - Chapter 5: Gases 5.1 5.2...

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Chapter 5: Gases 5.1 Early Experiments 5.2 The gas laws of Boyle, Charles, and Avogadro 5.3 The Ideal Gas Law 5.4 Gas Stiochiometry 5.5 Dalton’s Law of Partial Pressures 5.6 The Kinetic molecular Theory of Gases 5.7 Effusion and Diffusion 5.8 Collisions of Gas Particles with the Container Walls 5.9 Intermolecular Collisions 5.10 Real Gases 5.11 Chemistry in the Atmosphere
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• Helium He 4.0 • Neon Ne 20.2 • Argon Ar 39.9 • Hydrogen H 2 2.0 • Nitrogen N 2 28.0 • Nitrogen Monoxide NO 30.0 • Oxygen O 2 32.0 • Hydrogen Chloride HCl 36.5 • Ozone O 3 48.0 • Ammonia NH 3 17.0 •M e t h a n e C H 4 16.0 Substances that are Gases under Normal Conditions Substance Formula M (g/mol)
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Important Characteristics of Gases 1) Gases are highly compressible An external force compresses the gas sample and decreases its volume, removing the external force allows the gas volume to increase. 2) Gases are thermally expandable When a gas sample is heated, its volume increases, and when it is cooled its volume decreases. 3) Gases have low viscosity Gases flow much easier than liquids or solids. 4) Most Gases have low densities Gas densities are on the order of grams per liter whereas liquids and solids are grams per cubic cm, 1000 times greater. 5) Gases are infinitely miscible Gases mix in any proportion such as in air, a mixture of many gases.
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Pressure = force per unit area (Definition!) Gas samples exert a force pushing against any surface they contact, due to the impact of molecules striking the surface. Per unit area, it’s the same for all surfaces in or walls containing the same gas sample, unless the gas is flowing somewhere collectively. That pressure increases with molecular velocity and mass, and with the number of molecules per unit volume (number density). A sheet suspended within a gas sample gets same force on both sides, so it does not move. Average molecular velocity increases with temperature but decreases with molecular mass, in such a way that the mass effect cancels. This leaves pressure nearly proportional to temperature and to number density, but independent of type gas (or molecular mass): P = R(T)(n/V), where R is the gas constant, n = # moles = N/N A Pressure = P = Force Area
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Pressure of the Atmosphere • Called “Atmospheric pressure,” or the force exerted on earth’s surface by the gases in air • The force per unit area of these gases: Pressure = Force Area
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The Magdeburg sphere experiment
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Pressure of the Atmosphere • Called “Atmospheric pressure,” or the force exerted on earth’s surface by the gases in air • The force per unit area of these gases. • Can be measured using a barometer: Pressure = P = Force Area
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Figure 5.1: A Torricellian barometer. Density of Mercury = 13.6 g/cm 3 760 mm column of 1 cm 2 area (76 cm) weight = volume x density = (76 cm x 1 cm 2 ) x 13.6 g/cm 3 = 1030 g = 1.03 kg = 2.28 lbs P = force / area = weight / area = 2.28 lbs / cm 2 = 14.7 lbs / in 2 = 1.00 atm.
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Construct a Barometer using Water!
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This note was uploaded on 01/18/2012 for the course CHEM 142A taught by Professor Campbell during the Fall '11 term at University of Washington.

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C142A Ch 5 2011 lecture notes - Chapter 5: Gases 5.1 5.2...

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