Final Review

# Final Review - Exam 1 Gases o States of matter solid liquid...

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Exam 1 Gases o States of matter- solid, liquid, gas o Gas ± material with no definite shape or volume ± move more readily ± N2 is the most common gas (air) ± Air: N 2 , O 2 , H 2 O, Ar, He, CO 2 , O 3 , CH 4 o What kind of substances are gases? ± Mostly small molecules ± Some free elements H 2 , O 2 , Ar, Noble gases, N 2 , F 2 , Cl 2 , Br 2 , Hg ± Smells are molecules in their gas state o Microscopic characteristics ± Particles are randomly moving ± Particles will collide- elastic collisions (don’t lose energy) ± Particles exert pressure by bouncing off walls ± Simple relationship between moles, pressure, volume, and temperature (PV=nRT) o Microscopic Properties ± Gases are compressible ± Lower density than solids of liquids ± Completely mix- no “immiscible” gas o Units of Pressure ± Pressure = F/A (Force/Area) SI units of Area = m 2 SI units of Force = N N/m 2 = 1Pa Atmospheric Pressure at sea level = 101 kPa = 1atm 760 torr = 1atm = 760 mm Hg ± Relationship of physical properties of gas V P (inverse proportion) o V α 1/P o V = k(1/P) where k=constant o VP = k o Boyle’s Law ± V 1 P 1 = V 2 P 2 V T o V α T o (V 1 /T 1 ) = k = (V 2 /T 2 ) ± No change in pressure o Charles’ Law ± (V 1 /T 1 ) = (V 2 /T 2 ) V n (n = # of moles) o V α n o Avogardro’s Law

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± (V 1 /n 1 ) = (V 2 /n 2 ) No change in P or T The number of moles does not change by substance! PV=nRT o R = ideal gas constant = 0.082057 Latm/molK o See Figure 1 o Density of a Gas ± Much lower than liquids or solids ± g/L or g/cm 3 ± D = PM/RT M = molar mass o Dalton’s Law of Partial Pressure (1801) ± Total pressure of the mixtue of gases is equal to the sum of their individual pressures ± Implies that P depends on the total number of moles, not the chemical nature ± P tot = P a + P b + P c +…. . ± P tot = n tot RT/V If you have two gases, n tot = n a + n b P tot = (n a + n b )RT/V o Kinetic Molecular Theory of Gases ± Gas molecules are separated by large spaces ± Random motion with collisions ± Average kinetic energy of gas particles is proportionate to the gas’s temperature ± U = (3RT/M) Where U = average of the square of speeds and M = molar mass Average speed depends on T and molar mass RMS (root mean square) o Gas Diffusion ± Gradual mixing of 2 gases randomly ± Graham (1832)- at a constant pressure and temperature, the rate of diffusion of gases is inversely proportional to the square of M ± r 1 /r 2 = (M 1 /M 2 ) r = rate M = molar mass ± Lighter gas diffuses more quickly o Gas Effusion ± Process by which gas escapes out of a tiny opening in a container ± Graham’s law works here too o Ideal Gas Assumptions ± no attractions between molecules (“stickiness”) ± assume that the molecules have no volume ± This is where the real gas law comes in Intermolecular Forces o Forces that hold molecules together
o Not bonds! Only forces/attractions

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## This note was uploaded on 07/25/2008 for the course LBS 172 taught by Professor Laduca during the Spring '08 term at Michigan State University.

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Final Review - Exam 1 Gases o States of matter solid liquid...

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