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Week 7 of Dr. Ma's notes

by: Alexi Martin

Week 7 of Dr. Ma's notes CHEM 1200

Alexi Martin
GPA 3.58

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This covers Kinetics and Acid Base equilibirum
Chemistry II
Dr. Alexander Ma
Class Notes
Chemistry II
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This 5 page Class Notes was uploaded by Alexi Martin on Thursday March 10, 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 29 views. For similar materials see Chemistry II in Chemistry at Rensselaer Polytechnic Institute.

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Date Created: 03/10/16
Ch 15 Chemical Kinetics  example 1:  ​lnk2/2k1= ­Ea/R(1/300)­(1/31)   ln2 = Ea/8.31491.075x10^­4   ​ Ea= 5.35x10^4 J/mol  Collision Theory of Kinetics  ­ in order for most reactions to take place, molecules must collide at 10^9 collisions/sec  ­ molecules collide if:  ­ 1 enough E to break bonds  ­ 2 colliding in the proper orientation  ­ to be effective they must be fast enough to overcome bonding forming, some effective, some  are not  Effective Collisions  ­conditions net= effect  ­greater the frequency the greater the reaction rate  ­effective collision formation of the activated complex  ­ A is the frequency factor   ­collision frequency z # collisions/sec more=effective  k=A(e^ ­Ea/RT = pz(e^­ Ea/RT)  ­ orientation factor p<1 most, more complex<p, atoms where p equals p spherical, p>1 e­  transfer   ­ old bonds break and new bonds can form  Molecular interpretation­ reactant nature  ­ reactions occur faster in solutions­ mixing more particle contact, segregation more  collisions, break bonds that need to be broken  ­ different rates and increase in potential energy closer to the activated complex and a  reaction can lower the activation energy  Temperature  ­ increase temperature increases kinetic energy  ­ min amount kinetic energy causes potential energy to form activated complex  ­ an increase in the number of molecules that can overcome activation complex  Concentration/partial pressure  ­except 0 order reactions  ­more molecules=more collisions for effective collisions yield a biggers curve  multi­molecular collisions  ­ made up of several small reaction collisions 3 or fewer molecules  Reaction Mechanism  ­ nist reactions occur in small reactions with 1,2 or 3 molecules  ­ series of reactions: reaction mechanism  ­ knowing rate law helps us understand the sequence in the mechanism  example 2: H​ 2+2ICl­>2HCl+I2  1 H2+ICl­>HCl+HI  2HI+ICl­>HCl+I2    produced and consumed=intermediates(HI) the overall rate law cannot have  the concentrations of intermediates  ­reactions are elementary they cannot be broken down into simpler steps, molecules interact  directly without any other steps  Molecularity  ­ # reactant particles in an elementary step  ­ unimolecular 1 bi 2 ter 3( rare)  Rate laws for elementary reactions  ­ each step is its own reaction with its own rate law and Ea  ­ rate law overall is found experimentally  ­ can be deduced using equation of the slow or the rate law step  rate determining steps  ­ one step occurs slower than others  ­ result formation of products cannot occur faster than the slowest­ which determines the  overall reaction rate  ­ slowest step­ rate determining step  increases Ea  example 3:  NO2+CO­> NO+CO2 rate= k[NO2]2  NO2+NO­>NO3+NO​  slow rate determining step  NO3+CO­>NO2+CO2  Validate  1. elementary steps yields an overall reaction  2. rate law predicted must be consistent with experimentally determined rate law  Mechanisms with a fast initial step  ­ may contain intermediates  ­ substitute in rate law: left hand side must be just reactants  example 3: ​NO­>N2O2  H2+N2O2­>H2O+N2O​  slow rate=k[H2][N2O2]  H2+N2O­>H2)+HN  example4: ​ O3­> O2+O  O3+O­>2O2  rate f=rate r  K[O3]=k[O2][O] [O]=[O3][O2]^­1 rate=[O3]^2[O2]^­1  example5: ​A+B2­> AB+B slow rate determining step ​ate=k[A][B2]  A+B­>AB  2A+B2­>2AB  *Look at Clicker Problems*  Catalysts  ­affects rate by decreasing or increasing activation energy  ­consumed in early mechanism and formed in a later step   ­step up rate  can be hetero different or homo same (phases)  example 6:  HQ2R2+ Ri­> Q2R2­ +HR  Q2R2­. Q2R+R­ rate determining slow step    rate2=k2=[Q2R2­]= k[HQ2R2][R­]/[HR] r ate= k[HQ2R2][R­]/[HR]  Enzymes   ­ heterogenous biological reaction  ­ absorb substrate onto an active site­ products  ­ lock and key mechanism    Chapter 17 Acids and Bases (II)  ­ arrhenius acid­ produce H+ ions  ­ arrhenius base­ produces OH­ ions   ­ when they react they create water and a salt  HCl+NaOH­>NaCl+H2O  ­bronsted lowry acid­ H+ is transferred (proton) H donor   ­bronsted Lowry base­ H acceptor an atom with an unshared pair of electrons  H­A+B:­> :A­+H­B+  ­ Lewis acid e­ pair acceptor  ­ Lewis base e­ pair donor  Amphoteric substances :  ​can be an acid or base such as water  Strong or weak  ­ strong acid= strong electrolyte, all molecules ionize such as HCl  ­ strong base= strong electrolyte, all molecules ionize such as Ca(OH)2 strong  base(#nOH)= OH­ ion ion concentration  ­ weak acid/base weak electrolyte only some of the molecule will ionize  Strength of acid/base   ­ determined by finding equilibrium constant Ka or Kb  ­ farther towards products: stronger  ­ father towards reactants: weaker  Ka  ­ size of equilibrium  Ka=[acid][H3O+]/[HA]  ­ larger the Ka the stronger acid  Ion product of H2O  ­H3O+ and OH­ the same [H3O+][OH­]=1.0x10^­14 @ 25 degrees C  ­called Kw, water dissociation constant  ­one is given than the other can be calculated  ­ if [H3O+] is greater [OH­] must be lower  pH  ­pH of water is 7(neutral) greater than 7 is basic, less than 7 is acidic  ­pH= ­log[H3O]+   ­H3O+=10^­pH  ­pH+pOH=14  pK (unique to acid/ base strength)  pKa= ­log(Ka) or 10^­Ka=pKa  pKb= ­log(Kb) or 10^­Kb=pKb  ­ stronger acid smaller the pKa Kb=[OH­][HB]/[B]  ­ stronger the base the smaller the pKb  ­ larger the Kb stronger base  pH of a weak acid  ­solve equilibrium H3O+ using I.C.E. chart  Polyprotic Acids  ­more than 1 H+ donor  ­can have 1,2 or 3  ­ionizes in steps, each is removed sequentially  ­ each removal of H is harder  Acid­Base of salts  ­ salt water soluble ionic compounds  ­ cation strong base, weak acid anion are basic Na HCO3  ­ cation weak base anion strong acid NH4Cl  ­ stronger the acid, weaker the conjugate base  ­ strong the base weaker the acid  ­ weaker the acid, stronger the conjuagate base  ­ weaker the base   ­ stronger the acid  example 1: ​NO3­ netural HCO3 basic  Ka(Kb)=Kw  Metal cations as weak acids  ­ cations increase charged metals Al3+ is acidic  example 2: ​C5H5NH2+ acidic Ca2+ neutral Cr3+ acidic  ● think about where they come from  Acidity and Bascitiy of ions  ­anion of a strong acid neutral HCl HBr= Cl­ Br­ So4 2­  ­cations of strong bases are neutral Ca(OH)2 NaOH= Na+ K+ Mg+ Ca2+  ­anions of weak acids are basic HF= F­ CH3COO­  ­cations of weak bases are acidic NH3­> NH4+  ­small increase charged metal cations are weak acids Al3+ Cr3+  Classify Salt  ­salt cation is a counter ion of a strong base and anion of conjugate base of a strong acid=  neutral NaCl Ca(NO3)2 KBr  ­salt cation counter ion of a strong base and anion of a conjugate base of a weak acid = basic  Ca(C2H3O2)2 KNO2  ­salt cation is a conjugate acid of a weak base and anion is a conjugate base of a strong acid=  acidic NH4Cl  ­salt cation increased charged metal ion, anion conjugate base of a strong acid acidic Al(NO3)3  ­salt cation of conjugate acid of weak base when conjugate base of weak acid, pH depends on  strength of the acid or the base NH4F­ = acidic  Ka>Kb acidic Kb>Ka basic  example 3: S​ rCl2 neutral AlBr3 acidic CH3NH3NO3 acidic   NaCHO2 basic NH4F acidic KNO3 neutral COCl3 acidic Ba(HCO3)2 basic CH3NH3NO3 acidic  ionization of polyprotic acid  Ka1>Ka2>Ka3 except H2SO4  example 4:  12 M H2CO2  Ka= 4.3x10^­7 Ka2= 5.6x10^­11 *Ka2 does not affect pH*  H2CO3+H2O­>HCO3­+H3O+  HCO3­+H2O­>CO32­+H3O+  I 0.12 +x +x x^2/0.12=4.3x10^­7 x= ­log(2.27x10^­4)=3.64  C­x  H2SO4  ­strong acid completely dissociates use Ka2  example 5: HSO4­ +H2O­>SO42­+H3O+  I 0.01 0.01 0.01 x(0.01+x)/0.01­x use quadratic formula  C ­x +x +x x= ­log(0.045)=​.839  Binary acid  ­ strength increases across a period, H­C<H­N<H­O<H­F   ­ increase in bond energy the weaker the acid  ­ strength increases down a column H­F<H­Cl<H­Br<H­I  Oxyacids  ­more electronegative Y atom stronger HClO>HIO  ­more oxygens=stronger the acid  ­acidity increases down the group  ­acidity of oxyacids increases across a period HNO3>H2CO3>H3PO4>H3BO3  ­increase of oxidation # on central atom, stronger the oxy acid  example 6: ​cidity (least to most) H3AsO3<H3PO3<H3PO4<HNO3  acidity (least to most) HS­<H2S<HCl<HBr  basicity (least to most) NO3­<HCO3­<CO32­<BO33­   


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