Week 4-5 of notes
Week 4-5 of notes CHEM 1040 - 003
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CHEM 1040 - 003
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This 10 page Class Notes was uploaded by Olivia Hammond on Saturday February 27, 2016. The Class Notes belongs to CHEM 1040 - 003 at Auburn University taught by Ria Astrid Yngard in Spring 2016. Since its upload, it has received 33 views. For similar materials see Fundamental Chemistry II in Chemistry at Auburn University.
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Date Created: 02/27/16
CHAPTER 15 Dynamic Equilibrium - At equilibrium: the rate of the forward reaction = the reate of the reverse reaction. - Concentrations of products and reactants are constant and not changing over time. The Reaction Quotient and Equilibrium Constant - In General: - Reaction Quotient: the concentration of you products over the concentration of your reactants, with each concentration raised to a power that corresponds to the coefﬁcient in a balanced chemical equation; this ratio changes over time due to changing concentrations of the reactants and products. - Equilibrium Constant: similar to the reaction quotient except this is only when the system is at equilibrium; this ratio does not change over time - Because each concentration is in the units of mol/L both Kc and Qc are unitless. Problem 1 [worked example 15.1 from book and corresponds to problems in slides] Magnitude of Kc - Kc Large: If Kc is large, the reaction tends to “lie to the right” or the reaction will “favor products” - Kc Small: If Kc is small, the reaction tends to “lie to the left” or the reaction will “favor reactants.” - Kc in between: If Kc is neither small nor large, the equilibrium mixture will contain comparable amounts of both reactants and products; therefore, neither one is favored over the other. Qc and Kc The value of Qc changes as the reaction progresses. When the system reaches equilibrium Qc is equal to the equilibrium constant. You start off with only NO2. Over time, NO2O4 increases, but if you were to continue the reaction, the reaction will continue until it reaches equilibrium. 2 2 @equilibrium: Kc = [NO2] Before equilibrium: Qc = [NO2] [NO2O4] [NO2O4] Heterogeneous Equilibria - Heterogeneous Equilibria: the species in a reversible reaction are in different phases (ex: dissociating chalk) - The concentrations of solids and pure liquids do not appear in the equilibrium expression. [because C is in the solid phase it does not appear in the equilibrium constant] Manipulating Equilibrium Expressions 1. If the equation is reversed, invert the equilibrium constant. 2. If the coefﬁcients in the reaction are multiplied by a factor, Kc1, is raised to a power equal to that same factor. 3. The equilibrium constant for a reaction made up of two or more reactions is the product of the equilibrium constants for the individual reactions. Problem 2 [corresponds to the problems in the slides] Gaseous Equilibria Kp = pressure at equilibrium Pa, Pb = partial pressures at equilibrium Problem 3 [corresponds with problems on slides] Chemical Equilibrium and Free Energy A reaction with a very large K (>>>1) has a (-) ∆G —> spontaneous in the forward direction A reaction with a very small K (<<<1) has a (+) ∆G —> this will not occur in the forward direction spontaneously Chemical Equilibrium and Q Q is a ratio of the system before it reaches equilibrium If Qc is greater than Kc; products are favored If Qc is smaller than Kc; reactants are favored So, in the case of the example Qc = 111 which is greater than Kc; therefore, there is a lot of HI and little H2 and I2. For Qcc to equal to Kc, H2 and I2 must increase and HI must be consumed to reach equilibrium. If Q > K, the reaction procedes from right to left until equilibrium is reached If Q < K, the reaction procedes from left to right until equilibrium is reached. If Q = K, the reaction is a equilibrium Relationship between ∆G and ∆G Reaction not at equilibrium: ∆G = ∆G + RT(lnQ) - When not at equilibrium the spontaneity of a reaction is dependent coupon the concentration of the reactants and products (Q) Problem 4 [in correlation with the problem given in slides] Relationship between∆G and K o Reaction at equilibrium: ∆G= 0 and Q = K Problem 5 [in correlation with problem on slides] Calculating Equilibrium Concentrations Use of initial reactant concentrations to determine equilibrium concentrations: 1 Use initial concentrations to calculate the reaction quotient, Q, and compare Q to K to determine the direction in which the reaction will proceed. 2 Deﬁne x as the amount of a particular species consumed, and use the stoichiometry of the reaction to deﬁne (in terms of x) the amount of other species consumed or produced. 3 For each species in the equilibrium, add the change in concentration to the initial concentration to get the equilibrium concentration. 4 Use the equilibrium concentrations and the equilibrium expression to solve for x. 5 Using the calculated value of x, determine the concentrations of all species at equilibrium. 6 Check your work by plugging the calculated equilibrium concentrations into the equilibrium expression. The result should be very close to the c stated in the problem. Problem 6 [in correlation with problem on slides] Problem 7 [in correlation with problems on slides] Le Chatelier’s Principle - When a stress is applied to a system at equilibrium, the system, will respond by shifting in the direction that minimizes the effect of the stress. - It will try to minimize the change that is done on the system - An equilibrium that shifts to the right is one in which more products are produced by the forward reaction. - An equilibrium that shifts to the left is one in which more reactants are produced by the reverse reaction. - stress: - addition/removal of a reactant or product - pressure/volume change - temperature change Addition/removal Effect - Removal of NH3: shifts right —-> the reaction will try and make more NH3 to compensate for the NH3 that was removed from the equation - Addition of NH3: shifts left <— the reaction will want to use up the excess NH3 that has been added to the equation - Addition of N2: shifts right —> the reaction will want to use up the excess N2 that has been added to the equation - Removal H2: shifts left <— the reaction will try and make more H2 to compensate for the H2 that has been removed Inert Gas: a gas that does not react or inﬂuence the other gases in the reactants or products Volume/Pressure Change Effect - When volume is decreased, the equilibrium is driven toward the side with the smallest number of moles of gas Pressure increases when volume decreases: shift to right —> “favor products” Pressure decreases when volume increases: shift to left —> “favor reactants” - Addition of inert gas that is neither reactant nor product has no effect on the vapor or pressure at equilibrium Temperature Change In an exothermic reaction you must treat heat as a product In an endothermic reaction you must treat heat as a reactant. - In an exothermic reaction: An increase in temperature causes the reaction to shift to the left because it is an addition of energy. A decrease in temperature causes the reaction to shift to the right because it is a release of energy. - In an endothermic reaction: An increase in temperature causes the reaction to shift to the right; a decrease in temperature causes the reaction to shift to the left. Problem 8 [corresponds with the problems given in slides]
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