Temperature dependence of reaction rates Typically rates of
Temperature dependence of reaction rates Typically rates of reactions double for every 10oC rise in temperature, k =Ae(E a /RT ) Arrhenius equation Ea: activation energy A: frequency factor Ea lnk =lnA
RT An Arrhenius plot of ln k against 1/T is used to determine Ea and A The higher the Ea the stronger the temperature dependence of the rate constant
Collision Theory Collisions between two (or more) atoms/molecules required for a reaction. However, every time two reactants collide they may not react As temperature increases: atoms/molecules collide more frequently kinetic energy of atoms/molecules increases Collision theory: reaction occurs only if the reactants collide with a kinetic energy of at least the activation energy, and
they do so in the correct orientation. Kinetic energy is important Orientation is important Cl N O 2 AB -> A2 + B2
2 NOCl 2 NO + Cl2 Animation 1 Animation 2 Animation 3 k =Ae(E a /RT ) Ea
lnk =lnA RT The factor e-Ea/RT: fraction of molecules that have at least the minimum energy required for reaction. For an Ea = 40 kJ/mol Temperature (K) e-Ea/RT
298 9.7 x 10-8 400 5.9 x 10-6 600
3.3 x 10-4 A: reflects orientation effect or steric effect Measuring k as a function of T Ea to be determined k2 Ea 1 1 ln = ( ) k1
R T2 T1 Reaction coordinate diagram Activated complex or transition state - highest energy along reaction coordinate Reactants must collide with sufficient energy to reach this point and collide in a preferred
orientation to form the activated complex E = (Ea)forward - (Ea)reverse Higher temperatures favor products for an endothermic reaction and reactants for an exothermic reaction Endothermic reaction: Ea(forward) > Ea(reverse) Exothermic reaction: Ea(forward) < Ea(reverse)
CH3OH(aq) + H+(aq) CH3OH2+(aq) CH3OH2+(aq) + Br- (aq) CH3Br + H2O(aq) Catalysis Catalyst: a compound which speeds up the rate of a reaction, but does not itself undergo a chemical change. Simple mechanism A + catalyst intermediates
intermediates B + catalyst Overall: AB Concentration of catalyst is included in k; hence k varies with concentration of catalyst Presence of a catalyst provides an alternate path with a lower Ea
2H2O2(aq) 2H2O(aq) + O2(g) In the absence of a catalyst, Ea = 76 kJ/mol In the presence of a catalyst (I-); Ea = 57 kJ/mol; rate constant increases by a factor of 2000 H 3C
CH3 C H H (g) C
H cis-2-butene CH3 C H 3C C
(g) H trans-2-butene Catalyzed by I2 Pt C2H4(g) + H2(g) C2H6 (g)
Example of heterogenous catalysis A catalyst does not effect the thermodynamics of the reaction G is not affected by catalyst; neither is K Equilibrium concentrations are the same with and without catalyst; just the rate at which equilibrium is reached increases in the presence of a catalyst K = k1/k-1; catalyst speeds up both the forward and reverse
reaction Enzymes Practically all living reactions are catalyzed by enzymes; each enzyme specific for a reaction. Enzymes typically speed up rates by 107 - 1014 times rate of uncatalyzed reactions Ea for acid hydrolysis of sucrose: 107 kJ/mol Ea for catalyzed acid hydrolysis of sucrose: 36 kJ/mol Rate increase of 1012 at body temperature
E + S ES ES P + E Poisoning a catalyst Arsenic poisoning: Ingestion of As(V) as AsO43- results in reduction to As(III) which binds to enzymes, inhibiting their action Nerve gases - block enzyme-controlled reactions that allow nerve impulses to travel through the nerves.
Catalytic Converters Incomplete combustion of gasoline produces CO, hydrocarbon fragments (CmHn) High temperature in the engine causes oxidation of N 2 to NO and NO2 Conversion of these pollutants to less harmful compounds is speeded up in the presence of catalysts. 2 NO(g)
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Hess's Law Definition of Hess's Law If a reaction is carried out in a series of steps, ΔH is the sum of all the enthalpy changes CH4(g) + O2(g) → CO2(g) + 2H2O(g) ΔH=-802 2H2O(g) → 2H2O(l) ΔH=-88 ----- CH4(g)...