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22 Cards in this Set
- Front
- Back
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Standard Temperature and Pressure
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1 atm and 0 degrees C
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mean free path
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disance travelled by a gas molecule between collisions
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properties of gas mixtures
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-unlike liquids all gases will mix with each other regardless of polarity and form homogenous mixture
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Kinetic Moleclar theory of an ideal gas
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-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 |
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Ideal Gas Law
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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 |
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standard molar volume
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22.4 liters
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partial pressure of a gas
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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) |
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Average Translational KE formula for a gas
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KEavg=3/2 RT
-just average for mole of gas, molecule chosen at random might have any KE |
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RMS velocities of two gases in homogenous mixture
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v1/v2 = squareroot M2/ squareroot M1
-notice v1 on top and m1 on bottom |
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Effusion
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-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 |
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Chemical Kinetics
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how fast equillibrium will be achieved
-for a reaction need a collision and for a collision molecules need to reach activation energy |
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Arrhenius equation
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-gives rate constant K
-greater the temperature greater the K |
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Equations for reaction rates
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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 |
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Recognizing reaction orders
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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 |
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Reversible Reaction rates
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slowest step = rate determining step
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When slow step is second step
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-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] |
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Catalysts
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-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 |
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heterogenous catalyst
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-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 |
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homogeneous catalyst
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-same phase as reactants and products
-often aqueous acids or bases act as homogenous catalysts |
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Chemical Equilibrium
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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 |
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law of mass action
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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 |
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Reaction Quotient
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-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 Q<K: increase products decrease reactants: rightward shift Q>K increase reactants decrease products: leftward shift |
