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Week 6 notes

by: Alexi Martin

Week 6 notes CHEM 1200

Alexi Martin
GPA 3.58

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About this Document

these notes cover part of chapter 15
Chemistry II
Dr. Alexander Ma
Class Notes
Chemistry, kinetics
25 ?




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This 5 page Class Notes was uploaded by Alexi Martin on Friday March 4, 2016. The Class Notes belongs to CHEM 1200 at Rensselaer Polytechnic Institute taught by Dr. Alexander Ma in Spring 2016. Since its upload, it has received 22 views. For similar materials see Chemistry II in Chemistry at Rensselaer Polytechnic Institute.


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Date Created: 03/04/16
Ch 15 Chemical Kinetics    Chemical Kinetics  ­is defined as the speed of the reaction: the reaction rate, depends on molarity, temperature,  pressure,etc.  ­transient condition  ­ability to control is important  Low T, synthesis of diamond and d/C  ­thin film can be deposited using CVD  ­non equilibrium lower temperature, kinetically dominated  ­graphite preferable over diamond because it forms at a lower temperature and lower pressure.  ­rate, graphite adding hydrogen  ­formation of diamond formed through adding hydrogen  Important Q in Kinetics  ­practical value  ­mechanism of reaction  Reaction Rate  ­ concentration decreases over time or increases over time  rate=∆concentration/∆time=∆product/∆time= ­∆reactant/∆product  ­the lower the ∆ the more accurate  ­rate as the limit approaches 0 ∆product/∆t   Reaction rate and stoichiometry  ­ change in number of molecules is a change for another  ­ change in concentration of each substance is multiplied by 1/coefficient  Average rate  ­ change in measured concentrations in any particular time period  Instantaneous Rate  ­slope of the curve or the first derivative of the function  example 1: ​­1/3x(∆I­/∆t)= ­⅓ (0.868­1)/10 =(4.4x10^­3)­2= ​ ­8.8x10^­3 M/s  example 2: ​(2.4x10^2 g/109.9 g)/.2 L  ­d[Br2]/dt=dBr/dt rate= ½ d[NOBr]/dt = t d[Br]/dt  ­½ (­10.92 M)/300 =d[Br2]/dt =​ 1.82x10^­2 M/s  Measuring Reaction Rate  1. reactions that are less than one hour should have continuous measuring, polarimetry (chiral  molecules), spectrophotometry(wavelength absorbed), total pressure.  2. reactions that are greater than one hour, sample at various times­ gas chromatography, draw  off aliquots  Factors affecting reaction rate   ­ small molecules react faster than large molecules  ­ gases react faster than liquids  ­ ions react faster than molecules  temperature​ ­ higher the temperature the higher the reaction rate        catalysts­ increase speed without being consumed, some catalysts speed up in the + direction,  some slow down in the negative direction, can be homogenous(same phase), or  heterogeneous( a different phase)  concentration​­ an increase in the molarity of the reactant  or the pressure of the reaction  increases the reaction rate  Rate Law  ­ must be determined experimentally  rate= k[A]^n[B]^m  Reaction order  ­sum of all the exponents of the reactants  example 3: ​ 8.6x10^­5=k[1.00x 10^­6][3x10^­6] ​k= 2.2x10^7 1/Ms  example 4: ​rate=k[acetaldehyde]^2  6.73x10^­6[1.75x10^­3]= ​ 2.06x10^­11 M/s  Find rate law: initial rate method  ­ experimental, reaction depends on molarity of reactants  ­ changing initial reactant changes initial rate  Rate=k[A]^n  ­ n=0 0 order rate=k  ­ n=1 first order rate=k[A]  ­ n=2 2nd order rate=k[A]^2  example 5:  rate=k[NO2]^n[CO]^n NO2 0.2/0.1 0.1=n=1 ​ rate=k[NO2]^2 k= 0.21 1/Ms  example 6:​ k[NH4+]^n[NO2­]^m= rate   NH4+=0.09600/0.0200=32.3/10/8=3=3^x=1  NO2=0.0404/0.0202=21.6/10.8=2=2^x=1  r ate=[NH4+][NO2­]k   k=  ~clicker problem~​  0.060/0.020=0.248/0.0276=3=3^x [ClO2]^3[OH­]=1  D. 230 1/M^2s  graphical method  ­must be determined experimentally  ­graph k vs. time or concentration vs time  Integrated Rate Law  ­depends on concentration  Half Life  ­length of time for ½ the reaction/reactants are depleted or is over  ­radioactive decay  ­1st order concentration of reactant is constant  Kinetics vs. Thermo  ­kinetics oscillates while thermo goes straight to the product  Zero order reactions  rate=k[A]^0=k  ­constant rate reaction  [A]= ­kt+[A]initial  ­[A] vs time is a straight line  t1/2=[A]initial/2k  ­rate is in M/s k=M/s  First order reaction  ­ rate= k[A]=[A]k  ­ ln[A]=kt+ln[A]initial  ­ t1/2=ln2/k=0.693/k  ­ half life is constant   ­ rate=M/s k=1/s  Second order reaction  ­rate=k[A]^2  ­ 1/[A]=kt+1/[A]initial  ­t1/2= 1/(k[A]0)  ­rate=M/s k=1/Ms  ­ half  life is inversely proportional to [A]0  example 7: ​2NO2­>2NO+O2  ​ linear 1/[A]0 second order  example 8: ​SO2Cl2­>SO2+Cl2 1st order   ln[A]= kt+ln[A]i, ln[A]=(2.9x10^­4)865+ln0.0225 ​[A]=0.0175 M  example 9: !​­> 2R 2nd order   1/[A]=kt+1/[A0], 1/0.0010=k500+1/0.010 =1 ​ .81/Ms=k  Graphical   ­ linear [A] vs time 0order  ­ linear ln[A] vs time 1st order  ­ linear 1/A vs time 2nd order  example 10: ​ t1/2=ln2/k 1st order t1/2=ln2/0.271 ​ .56 s  example 11: ​1/kA=t1/2 1/1.8(0.010)=​  55.56 s  C­14 dating  ­ t1/2= 5730 1st order decay process k=1.21x10^­4 1/yr  ­ C14:C12 r0=`.1x`0^­12  ­ at death C14 will decay and the ratio will be reduced  ­ lnr0/r1=1.21x10^­4t  The effect of temperature on rate  ­ changing t changing rate constant k increase temperature increases the rate  ­ A frequency factor and Ea activation energy  ­ K=A(e^ ­Ea/RT)  Activation E and activated complex  ­ E barrier to all reaction  ­ Ae amount of E needed to convert reactants to an activated complex (transition state)  ­ activated complex an increase in energy partially broken and partially formed bond  complexes  Arrhenius equation  ­ exponential factor between 0 and 1  ­ fraction of molecules can move it over  ­ greater energy barrier fewer molecules can go over it   ­ increase in temperature increase in kinetic energy increase in reaction rate  Arrhenius plot  ­ln k vs 1/t is a straight line  lnk= ­Ea/R(1/T)+lnA  Arrhenius equation two point form  ln(k2/k1)=­Ea/R(1/T2­1/T1)  example 12: NO2+CO­> CO2+NO  ln(567/2.57)= ­Ea/8.314(1/895­1/701​Ea=145 kJ/mol 


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