Chapter 12

# Chapter 12 - Chapter 12 Study Guide ...

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Unformatted text preview: Chapter 12 Study Guide  Reaction Rate: the rate of the reaction, how quickly a reaction will go to  completion (Example: Think of it as MPH for chemicals)  Collision Theory: theory that states in order for a reaction to occur, there must  be a collision of molecules, molecules must have enough energy to overcome  the “energy barrier”, and molecules must be in proper orientation  Activation Energy: Minimum amount of energy needed to overcome the  energy barrier  (Ex: The min.  Amount of work  You need to do to  Pass a class.)  Things that affect reaction  rates:  Concentration:  increasing concentration increases reaction rates, because a  larger number of collisions occur.  Surface area:  increasing surface area of a solid increases reaction rates  because as surface area increases, the number of atoms exposed increases as  well, allowing more collisions.  Temperature:  a temperature increase increases the average kinetic energy.  A  increase in kinetic energy causes more collisions to occur, and also increases  the fraction of effective collisions.      Reason fraction of effective collisions increases:  the increase in kinetic  energy allows more molecules to break the energy barrier.  In other words,  more molecules meet the activation energy requirement.  Fraction of effective collisions:  (the number of effective collisions)/(total  collisions)  Catalyst: something that alters the pathway in which a reaction occurs without  itself being consumed in the reaction. (Example: Think about this. You need to  travel a 10 mile road which curves. A catalyst would be like a shortcut which  cuts 6 miles off the original distance but gets you to the same place.)    Intermediate:  something that forms temporarily during a reaction,  accompanying a catalyst.    Ex:  Step 1) O3(g) + Cl(g) ­> ClO(g) + 02    Step 2)  ClO(g) + O3(g) ­> Cl(g) + 2O2(g)  Overall: 2O3(g) + Cl(g) + ClO(g) ­> ClO(g) + Cl(g) + 3O2(g)  Net:  2O3(g)  ­­>  3O2(g)  ClO is the intermediate in this reaction, and Cl is the catalyst.  Cl is in both  reactants and products, so it is not consumed in the reaction.  ClO is the  intermediate because it is temporarily created during the reaction and then  consumed in Step 2.  Chemical Equilibrium  Chemical equilibrium: state in a chemical reaction where concentrations of  the reactants and products remain constant.  Products are converted to  reactants at an equal rate as reactants are converted to products. (Example:  The mall with two floors. You have the same # of people on both floors.  However, two people decide to leave the 2nd floor while two people come up  there. As long as the number is the same on both floors, equilibrium is  maintained. It doesn’t have to be the same people.)  Equilibrium Constant Expression: aA +bB   cC + dD    [C]c [D]d  Keq =  ______________    [A]a [B]b  Keq > 1  Products are favored (lies to the right)  Keq < 1 Reactants are favored (lies to the left)  Keq = 1 Similar amounts of reactants and products (lies in middle)  Homogeneous equilibrium:  equilibrium in which reactants and products are  in the same physical state (solid, gas, liquid)  Heterogeneous equilibrium:  equilibrium in which reactants and products are  in different physical states.  In heterogeneous equilibriums, solids and liquids are not included in the  equilibrium constant equation.    CO(g) + H20(l)    CO2(g) + H2(g)  Keq = [H2][CO2]  / [CO]  H20 is not included because it is not a gas or aqueous and because it is a pure  substance.  Le Chatelier’s Principle  Le Chatelier’s Principle: if a system at equilibrium is disrupted, it shifts to  establish a new equilibrium. (Example: Let’s say you have a scale where one of  the sides is heavier than the other. You want to equal out the sides so you take  some weight from the heavier side and move it to the lighter side. That’s what  happens chemically in equilibrium except you don’t have to do anything.)      A(g) + B(g)  C(g) + D(g)    Equilibrium Shifts due to Concentration Changes  Add reactant, equilibrium shifts right  Add Product, equilibrium shifts left  Remove Reactant, equilibrium shifts left  Remove product, equilibrium shifts right    Equilibrium Shifts Due to Volume Change  Relative Number of  Gaseous Molecules  in Balanced  Equation  Reactants <  products  Reactants >  products  Reactants =  products  Example  Increase Volume  Decrease Volume  N204(g)   2NO2(g)  N2(g) + 3H2    2NH3(g)  2NO(g)   N2(g) +  O2(g)  Shift right  Shift left  No Shift  Shift Left  Shift right  No Shift    Effects of Temperature Changes on the Position of Equilibrium  Type Of Reaction  Endothermic  Reaction  Exothermic  Reaction  Equation  Heat + A + B  C +  D  A + B  C + D +  Heat  Increase  Temperature  Shift right.  Keq  increases.  Shift left.  Keq  decreases.  Decrease  Temperature  Shift left.  Keq  decreases.  Shift right.  Keq  increases.  Treat heat like another reactant or product.   Increasing heat on one side will shift the  equilibrium to the other, and removing it on one side will shift equilibrium toward that  side.    ...
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