Take home Exam due Wednesday, Aug 26. In class Exam will be the that morning class multiple choice questions.

Transcription

1 Announcements Take home Exam due Wednesday, Aug 26. In class Exam will be the that morning class multiple choice questions. Updated projects Aug 28: answer what lab chemistry needs to get done to realize my product to market. 1) Identify chemistry that needs to be done but eliminate that which can not be done (explain why it can not be as it must fly with me first). 2) Make a plan that your team will implement in the lab. Describe what each member in your group will do in the lab period. 3) Provide a detailed costing analysis of your product.

2 Unanswered Questions Why does the reaction rate depend on temperature when there is no T in the equation? Why are [A] and [B] multiplied in the rate law? What is happening at the molecular level? Why does the reaction rate depend on concentration of reactants to varying degrees? Focus so far has been macroscopic--much like the gas law. We turn now to kinetic molecular theory for kinetics ---collision theory and transition state theory.

3 If we plot k vs T data we observe that k increases exponentially as T increases. Exponential increase of k with T Svante Arrhenius showed that the rate constant, k,varies with temperature exponentially: k = A e E a/rt = A exp( E a /RT)

4 Arrhenius showed that the rate constant, k,varies with temperature according to: k = A e E a/rt = A exp( E a /RT) k is the reaction rate constant Ea = Activation energy (specific for any given reaction) A = Frequency factor (related to number of collisions and orientation factor for any given reaction) T = temperature in Kelvin R = gas constant (R = J/mol. K) rate = k [A] m [B] n rate = A exp (-Ea/RT) [A] m [B] n We can see now why rate varies with T

5 We can transform the Arrhenius equation to a more useful graphical form by taking the natural logarithm of each side. k = A e E a/rt = A exp( E a /RT) ln k = ln (A exp( E a /RT)) Take ln of both sides ln k = ln A + ln(exp( E a /RT)) Expanding ln k = E a R 1 T + ln A y = -m x + b This equation has the form y = mx + b

6 If we collect data for k vs T and then plot ln k vs 1/T we can determine Ea. ln k = E a R 1 T + ln A y = -m x + b Arrehenius Plot Temp ( C) k (M -1 s -1 ) Slope = -Ea/R E E E-04 ln k y E E-02 x 1/T (K -1 )

7 We can also transform the Arrhenius equation into something more useful using ratios of two different rate constants at two different temps. k 2 k 1 = A e Ea/RT2 A e E a/rt 1 = e E a/rt 2 E a /RT 1 ln ( k2 = e E a/rt 2 + E a /RT 1 k 2 = exp E ( a 1 1 k 1 R T 1 T 2 ) = E ( a 1 1 ) R T 1 T 2 k 1 ) If we have data for k2 at T2 we can determine k1 and some other temperature T1

8 The decomposition reaction of hydrogen iodide, 2HI(g) H 2 (g) + I 2 (g) has a rate constant of 9.51 x 10-9 L/mol. s at 500. K and one of 1.10 x 10-5 L/mol. s at 600. K. Find the activation energy, E a.

9 The decomposition reaction of hydrogen iodide, 2HI(g) H 2 (g) + I 2 (g) has a rate constant of 9.51 x 10-9 L/mol. s at 500. K and one of 1.10 x 10-5 L/mol. s at 600. K. Find the activation energy, E a. ln ( k2 k 1 ) = E a R ( 1 T 2 1 T 1 ) E a = R ln k 2 k 1 ( 1 T 2 1 T 1 ) 1 E a = (8.314 J/mol K) ln E a = 1.76 x 10 5 J/mol = 176 kj/mol ( L/mol s L/mol s ) ( K 1 ) K

10 Macroscopic reaction rates that we measure in the lab can be explained microscopically by two models called collision theory and transition state theory. 1. Collision Theory--in order for a chemical reaction to occur reactant molecules must collide. Only those molecules with a certain minimum energy (Ea activation energy) and correct spatial orientation will transform into products. 2. Transition State Theory--model used to describe the energetics and what reactants and products look like in transforming reactants to products.

11 Collision Theory explains the microscopic basis of the rate law and correlates chemical reaction rates to: Frequency of Collisions Collisions must occur for a reaction to occur. Activation Energy, Ea Reactant molecules must have a minimum kinetic energy in order for collisions to lead to a chemical reaction. Temperature Impacts Kinetic Energy (collision energy) Increases in T increase the fraction of molecules that have sufficient Ea to react. Orientation of Reactants Molecules must be oriented in a certain way in 3-D space in order for a collision to lead to a chemical reaction.

12 The Arrehenius equation links the macroscopic rate constant to fraction of molecular collisions with proper spatial orientation and Ea of collision theory. k = A e E a/rt Activation Energy Gas Constant k = pze E a/rt Temperature (K) Fraction of collisions with proper orientations Constant related to collision frequency Fraction of collisions with sufficient energy for reaction

13 The reactants are multiplied in the rate law because a product reflects the reality that reactants must collide to transform to products. A + B C rate = k [A][B] Why Multiplication? A A A A A B B B 4 collisions B Add another molecule of A 2 x 2 = 4 6 collisions 3 x 2 = 6 A A A We multiply because the rate depends on the number of collisions (which is found by multiplication) B B B 9 collisions 3 x 3 = 9

14 The reaction rate constant, k, increases as the activation energy (Ea) decreases and as temperature increases. The Effect of Ea and T on the Fraction of Collisions With Sufficient Energy to Transform The rate constant is altered by changing the fraction of energetic molecules.

15 Results of Collision Theory increased T increased average speed of particles increased collision frequency increased reaction rate Significance of activation energy: only those collisions with energy equal to, or greater than, E a can yield products. Decreasing Ea and/or increasing T enhances the fraction of productive collisions, f increases k and therefore the reaction rate. Reactants must collide with a minimum Ea and with the proper orientation to react.

16 Activation energy, Ea, is the minimum kinetic energy needed for a reaction to occur. Higher T results in greater fraction of molecules with energy > Ea Increasing T serves to produce a larger fraction of molecules > Ea f = e -E a /RT called the Boltzman factor

17 Because all chemical reactions are reversible there is an activation energy, Ea for both the forward reaction and the reverse reaction. A + B Reactants kforward kreverse C + D Products

18 Activation Energy--Analogy Ea Minimum Energy Needed To Get Over Hill Ea is a property of the height of the hill (i.e. the chem reaction) Ea Eb-a ΔH The Hill = Chemical System

19 A reaction energy diagram can help us understand activation energy (Ea). A + B C + D

20 The Arrehenius equation links the macroscopic to the microscopic insight of collision theory. k = A e E a/rt Activation Energy Gas Constant k = pze E a/rt Temperature (K) Fraction of collisions with proper orientations Constant related to collision frequency Fraction of collisions with sufficient energy for reaction

21 Reactants must be collide with the proper spatial orientation in order to transform from reactant to product. Consider the reaction between an atom of chlorine and a molecule of nitrosyl chloride: Cl + NOCl NO + Cl 2 An Effective Collision...Cl collides withcl There are two possible ways that Cl atoms and NOCl molecules can collide; one is effective and one is not. Ineffective Collision...Cl collides with O

22 Reactants must be collide with the proper spatial orientation in order to transform from reactant to product. Collision 1 Collision 2 Collision 3 Collision 4

23 Energy Transition State Theory explains the energetics and substances look like as they are transforming reactants to products. --Postulates intermediate structures called transition states or activated complex and energy barriers (activation energy) as a reaction occurs. --A reaction energy diagram depicts the transition state, activation energy, and thermodynamics. A Reaction Energy Diagram A + BC (Reactants) A...B...C (Transition State) Ea H Reation Progress AB + C (Products)

24 Reaction energy diagrams are used to depict the energetics and events that occur as reactants are transformed to products.

25 The activation energy, Ea can be viewed as the energy required to stretch and deform bonds forming a an activated complex or transition state. Ea H

26 Reaction energy diagrams include the energetics for enthalpies and levels of activation energy. H > 0 H < 0 H < 0

27 RED for the reaction between CH3Br and OH-

28 The proposed transition state in the reaction between: CH 3 Br + OH - ===> CH3OH + Br - The TS is trigonal bipyramidal; note the elongated C-Br and C-O bonds

29 A key reaction in the upper atmosphere is O 3 (g) + O(g) 2O 2 (g) The E a(fwd) is 19 kj, and ΔH rxn for the reaction is -392 kj. Draw a reaction energy diagram for this reaction, a transition state, and calculate E a(rev).

30 O 3 (g) + O(g) 2O 2 (g) The E a(fwd) is 19 kj, and ΔH rxn for the reaction is -392 kj. Draw a reaction energy diagram for this reaction, a transition state, and calculate E a(rev). Not to scale! transition state E a(rev) = ( ) kj = 411 kj ΔH rxn = Hf - Hi = 392kJ

31 A catalyst is any substance that increases the rate of a chemical reaction without itself being consumed in the reaction. A catalyst increases reaction rate in both directions by increasing k via lowering the activation energy, Ea of the reaction. Uncatalyzed Pathway Catalyzed Pathway The reaction thermodynamics (enthalpy, entropy) are unaffected! No change in the yield of the reaction! The catalyzed reaction proceeds via a different reaction mechanism than the uncatalyzed reaction.

32 Catalysts lower the activation energy (increasing k) relative to an uncatalyzed reaction. Both reactions have the same enthalpy, H Catalyst does not alter the yield of the reaction relative to uncatalyzed pathway. The reaction pathway (mechanism) is different in a catalyzed reaction Potential Energy Ea No Catalyst Reactants Ea Catalyzed Ea No Catalyst Catalyzed Reaction Progress Uncatalyzed Ea Uncatalyzed Products

33 There are two general classes of catalysts heterogeneous catalysts and homogeneous catalysts. Heterogeneous catalysis, the reactants and the catalysts are in different phases (solid catalysts in liquid reactions or solid catalyst for gases. Haber synthesis of ammonia Ostwald process for the production of nitric acid Catalytic converters In homogeneous catalysis, the reactants and the catalysts are dispersed in a single phase, usually a liquid. Acid catalysis Base catalysis

34 Catalysts are used in catalytic converters to reduce nitrogen oxides and to clean up unburned HC s CO + Unburned Hydrocarbons + O 2 catalytic converter CO 2 + H 2 O catalytic 2NO + 2NO 2 converter 2N 2 + 3O 2