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Chem notes week 14

by: Andrea Scota

Chem notes week 14 CHE 106 - M001

Andrea Scota
GPA 3.7

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Chem notes for week 14
General Chemistry Lecture I
R. Doyle
Class Notes
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This 3 page Class Notes was uploaded by Andrea Scota on Saturday December 12, 2015. The Class Notes belongs to CHE 106 - M001 at Syracuse University taught by R. Doyle in Fall 2015. Since its upload, it has received 33 views. For similar materials see General Chemistry Lecture I in Chemistry at Syracuse University.

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Date Created: 12/12/15
Pink- mentioned in class Chem Notes Week 14 TEXTBOOK CHAPTER 10 (contd. Sections 10.5-10.9) Gases Further Applications of the Ideal-Gas Equation (10.5)  Using the ideal gas law we can relate the density of a gas to it’s molar mass: o  = dRT/P o Equation tells us that the density of a gas depends of pressure, molar mass, and temperature. The higher the molar mass and pressure, the denser the gas.  When have equal molar masses of two gases at the same pressure but different temperatures, the hotter gas is less dense than the cooler one (hotter gas rises)  We can use the ideal gas equation to solve problems involving gases and reactants or products in chemical reactions Gas Mixtures and Partial Pressures (10.6)  The total pressure of a mixture of gases equals the sum of the pressures that each would exert if it were present alone  Partial pressure: the pressure exerted by particular component of a mixture of gases  In gas mixtures, the total pressure is the sum of the partial pressures that each gas would exert if it were present alone under the same conditions (Dalton’s Law of partial pressures) o The pressure due to each component is additive  The partial pressure of a component of a mixture is equal to its mole fraction times the total pressure: o P =lX P i t o Mole fraction (X ) iI the ratio of moles of one component of a mixture to the total moles of all components The Kinetic-Molecular Theory of Gases (10.7)  The kinetic-molecular theory of gases accounts for the properties of an ideal gas in terms of a set of statements about the nature of gases. Can be summarized by: o Gases consist of large numbers of molecules that are in continuous, random motion. o The combined volume of all the molecules of the gas is negligible relative to the total volume in which the gas is contained o Attractive and repulsive forces between gas molecules are negligible o Energy can be transferred between molecules during collisions but, as long as temperature remains constant, the average kinetic energy of the molecules does not change with time o The average kinetic energy of the molecules is proportional to the absolute temperature. At any given temperature, the molecules of all gases have the same average kinetic energy  The pressure of gas is caused by collisions of the molecules with the walls of the container; magnitude determined by how often and home forcefully they strike the walls. o P is the average of the forces over a given area P=F/A  Absolute temperature is a measure of the average kinetic energy of its molecules  The individual molecules of a gas do not all have the same kinetic energy at a given instant. Their speeds are distributed over a wide range; the distribution varies with the molar mass of the gas and with temperature  The root-mean-square (rms) speed, u , is the speed of a molecule rms possessing kinetic energy identical to the average kinetic energy of the sample. It varies with proportion tot the square root of the absolute temperature and inversely with the square root of the molar mass: o U = 3RT/M. rms o The most probable speed of a gas molecule is given by u mp = 2RT/M Molecular Effusion and Diffusion (10.8)  Effusion: when gas escapes through a tiny hole  Diffusion: the spread of one substance throughout a space or throughout a second substance o Related to the speeds at which molecules move  The rate at which a gas undergoes effusion is inversely proportional to the square root of its molar mass (Grahm’s Law)  Because moving molecules can undergo frequent collisions with one another, the mean free path, the mean distance traveled between collisions, is short.  Collisions between molecules limit the rate at which a gas molecule can diffuse  For diffusion and effusion is faster for lower mass molecules than for higher mass ones Real Gases: Deviations from Ideal Behavior  Departures from ideal behavior increase in magnitude as pressure increases and as temperature decreases  Real gases depart from ideal behavior because: o The molecules posses finite volume o The molecules experience attractive forces for one another o These two effects make the volume of real gases larger and their pressures smaller than those of an ideal gas  Van-der Waals equation is an equation of state for gases, which modifies the ideal-gas equation to account for intristic molecular volume and intermolecular forces


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