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22 Cards in this Set

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 Standard Temperature and Pressure 1 atm and 0 degrees C mean free path disance travelled by a gas molecule between collisions properties of gas mixtures -unlike liquids all gases will mix with each other regardless of polarity and form homogenous mixture Kinetic Moleclar theory of an ideal gas -gas molecules have zero volume -gas molecules exert no forces other than repulsive forces due to collisions -gas molecules make completely elastic collisions (KE is conserved) -average KE of molecules is directly proportional to the temperature Ideal Gas Law PV=nRT Pressure Volume=#moles x universal gas constant x temp kelvin -when gas expands in container it does work on surroundings and pressure decreases due to loss in KE and increase in volume and temperature decreases -all gases that behave ideally will have same volume temperature, n, pressure standard molar volume 22.4 liters partial pressure of a gas total pressure of the gaseous mixture times the mole fraction of the particular gas -partial pressures of gases add up to total pressure (daltons law) Average Translational KE formula for a gas KEavg=3/2 RT -just average for mole of gas, molecule chosen at random might have any KE RMS velocities of two gases in homogenous mixture v1/v2 = squareroot M2/ squareroot M1 -notice v1 on top and m1 on bottom Effusion -spreading of gas from high to low pressure through a pinhole Rate1/Rate2= SquarerootM2/SquarerootM1 -similar to grahams law -can be used to estimate diffuseion rates Chemical Kinetics how fast equillibrium will be achieved -for a reaction need a collision and for a collision molecules need to reach activation energy Arrhenius equation -gives rate constant K -greater the temperature greater the K Equations for reaction rates rate forward= k[A]^a[B]^b k= rate constant for forward reactoin A, B= molecules in equation a,b= stoichiometric # of moles used as exponents in elementary reactions total order of reaction = a + b Recognizing reaction orders first order reaction: rate = k[A] A decreases exponentially graph of ln[A] vs. t = straight line with slope -k Second Order 2A-> products rate=k[A]squared plot 1/[A] gives straight line with slope k third order rate = k[A]cubed plot 1/(2[A]squared) straight line with slope k Reversible Reaction rates slowest step = rate determining step When slow step is second step -assume first step reaches equillibrium -for rate set forward reaction rate = to the reverse reaction rate for first step fast ex. 1. NO + Br2 -> NOBr2 slow 2. NOBr2 + NO -> 2NOBr k1[NO][Br2]=k-1[NOBr2] Catalysts -increase rate of reaction -lower Ea -does not alter equillibrium -concentration of catalysts are usually small compared to reactants and products, if it is large the rate changes little heterogenous catalyst -different phase than the reactants and products, usually solids -reactant may physically adsorb (van der wals) or chemically adsorb (covalent bond) to surface of catalyst -binding is exothermic homogeneous catalyst -same phase as reactants and products -often aqueous acids or bases act as homogenous catalysts Chemical Equilibrium Forward rate=reverse rates -no change in concentration of products -net rate= 0 -point of greatest entropy -set forward rates= to reverse rates A+B-> C+D rate forward=rate reverse kf[A]a[B]b=kr[C]c[D]d algebraic manipulation: kf/kr= [C]c[D]d/[A]a[B]b another name for kf/kr= Keq :law of mass action Keq= [C]c[D]d/[A]a[B]b law of mass action Keq= [C]c[D]d/[A]a[B]b - kf is reciprocal for kr -for Keq K is capitalized -depends only on temperature -solids or pure liquids are not used in equation Reaction Quotient -for reactions not at equillibrium Q=C]c[D]d/[A]a[B]b or products raised to coefficients/reactants raised to coeeficients can have any positive value -will always change towards Keq Q=K : equillibrium QK increase reactants decrease products: leftward shift