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Che 107 Dr Blue Exam 2 Study Guide

by: Alena Comley

Che 107 Dr Blue Exam 2 Study Guide Che 107

Marketplace > University of Kentucky > Chemistry > Che 107 > Che 107 Dr Blue Exam 2 Study Guide
Alena Comley
GPA 3.5

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This covers chapters 13, 14 and 15.2-15.5. rate, rate laws, equilibrium, ICE tables
General Chemistry II
Dr. Blue
Study Guide
rate, Rate Laws, Equilibrium, ICE tables
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This 6 page Study Guide was uploaded by Alena Comley on Sunday March 6, 2016. The Study Guide belongs to Che 107 at University of Kentucky taught by Dr. Blue in Spring 2016. Since its upload, it has received 27 views. For similar materials see General Chemistry II in Chemistry at University of Kentucky.


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Date Created: 03/06/16
Che 107 Exam 2 Study Guide 13.2 Rate=Change in concentration/ change in time −1 ∆ A ] −1 ∆[B] 1 ∆[C] 1 ∆[D] aA+bBcC+dD = = = = a ∆ t b ∆ t c ∆t d ∆t Reaction rates change with concentration, time, temperature, and presence of a catalyst 13.3 Rate law: relationship between the rate of reaction and concentration if reactants k- rate constant different for each reaction and set of conditions x y aA+bBcC+dD Rate Law: rate=k[A] [B] n2 rate2= k[A] rate1 k[ A]1 pick trial with bigger # for top Compare trials where only one reactant changes in concentration Overall order of reaction- sum of the orders of the reactants 13.4 Integrated rate law-dependence of concentration on time Rate law has units of M/s Half-life : time required for half of the reactant to be consumedderive half-lives 13.5 The effect of temperature on reaction rate −Ea Arrhenius Equation m n RT A=frequency factor; E = rate=k[A] [B] k=Ae a activation energy; R= gas constant (8.314 J/mol*K); T= temperature (K) Orientation factor- how do we position 2 molecules such that the collision leads to a reaction Collision frequency- how often do we get these molecules to collide As temperature ↑, the fraction of molecules with enough energy to surmount the activation energy barrier also ↑ −E a 1 k1 −E a 1 1 lnk= +ln Aln = ( − ) R T k2 R T2 T 1 As temp increases (T >2 )1 k 2 13.6 The rate-determining step is the slow step in the process 13.7 Catalysis: catalysts increase the rate of the reaction by lowering the activation energy Homogeneous- reactants and catalyst in the same phase Heterogeneous- reactants and catalyst in different phase Not consumed (overall)- react early in the reaction then are regenerated later 14.2 Dynamic equilibrium-rate of forward and reverse reactions are equal Concentrations of reactants and products remain constant (does not mean concentrations are equal) S =k P Condensation rate=vaporization rate gas H gas S, P are not equal but proportional (not all reactants consumed) 14.3 [C] D d aA+bB↔cC+dD K= a b Equilibrium constant A B K is dimensionless (no units) [products] /[reactants] [ ]- refers to concentration in Molarity 14.4 Concentration and Pressure Δn Relationship between K anP K C KP= K CRT) KP: must be in atm K C mol/L (M) R: L*atm/mol*K T:Kelvin Δn: change in moles of gas KPwill be equal to K Chen Δn=0 14.5 Solids and liquids are excluded from the equilibrium calculation 14.6 ICE tables I:initial concentration C:change in concentration E:equilibrium concentration K changes with temperature 14.7 The reaction quotient Calculate Q the same way as K When: Q<K reaction runs to the right (towards products) Q>K reaction runs to the left (towards reactants) Q=K no change (equilibrium) 14.8 Finding equilibrium concentrations Know K and all but one concentration/pressure Know K and initial concentrations/pressures −b± b√−4 ac Use quadratic formula x= 2a 14.9 Le Châtlier’s Principle When a chemical system at equilibrium is disturbed, the system shifts in a direction to minimize the disturbance Disturbances that affect equilibrium Changing the concentration of a reactant or product Changing the total pressure Changing the volume of the container Addition of inert gas Changing the temperature 15.2 Acids and Bases Acids: Sour taste Can dissolve many metals Neutralize bases Turn litmus red Bases: Bitter Slippery feel Neutralize acids Turn litmus blue 15.3 Arrhenius- limited to aqueous solutions Bronsted- Lowry- proton transferable Lewis- electron pair transfer Conjugate acid-base pairs: (1) share a common ion, and (2) differ by a single H + Amphoteric: substance that can act as an acid or a base 15.4 Acid strength and K a Strong acids are strong electrolytes and weak acids are weak electrolytes KNOW 6 STRONG ACIDS Hydrochloric Acid (HCl) Nitric Acid (HNO )3 Hydrobromic Acid (HBr) Perchloric Acid (HClO )4 Hydriodic Acid (HI) Sulfuric Acid (H 2O )4 Most acids are weak The stronger the acid, the weaker the conjugate base 15.5 Autoionization of water and Ph [H3O ] can be expressed the same as [H ] + -14 + - At 25°C, K w1.00 x 10 Kw=[H ][OH ] K wncreases as temperature increases


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