AnnouncementsTake home Exam due Wednesday, Aug 26. In class Exam will be the that morning class. 15-20 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.

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?

Why does the reaction rate depend on concentration of reactants to varying degrees?

What is happening at the molecular level?

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.

If we plot k vs T data we observe that k increases exponentially as T increases.

Exponential increase of k with T

k = A e!Ea/RT = A exp(!Ea/RT)

Svante Arrhenius showed that the rate constant, k ,varies with temperature exponentially:

Arrhenius showed that the rate constant, k ,varies with temperature according to:

• 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 = 8.314 J/mol . K)

k = A e!Ea/RT = A exp(!Ea/RT)

rate = A exp (-Ea/RT) [A]m [B]n

rate = k [A]m [B]n We can see now why rate varies with T

We can transform the Arrhenius equation to a more useful graphical form by taking the natural logarithm of each side.

k = A e!Ea/RT = A exp(!Ea/RT)

ln k = ln (A exp(!Ea/RT))

ln k = ln A + ln(exp(!Ea/RT))

ln k =!Ea

R

1

T+ ln A

y = -m x + b

Take ln of both sides

Expanding

This equation has the formy = mx + b

Temp(°C)

k(M-1 s-1)

283 3.52E-07

356 3.02E+05

393 2.19E-04

427 1.16E-03

508 3.95E-02

ln k

1/T (K-1)

Slope = -Ea/R

∆x

∆y

ln k =!Ea

R

1

T+ ln A

y = -m x + b

If we collect data for k vs T and then plot ln k vs 1/T we can determine Ea.

Arrehenius Plot

We can also transform the Arrhenius equation into something more useful using ratios of two different rate constants at two different temps.

k2

k1=

A e!Ea/RT2

A e!Ea/RT1= e!Ea/RT2 ! !Ea/RT1

= e!Ea/RT2 + Ea/RT1

k2

k1= exp

Ea

R

!1

T1! 1

T2

"

ln

!k2

k1

"=

Ea

R

!1

T1! 1

T2

"

If we have data for k2 at T2 we can determine k1 and some other temperature T1

The decomposition reaction of hydrogen iodide,

2HI(g) H2(g) + I2(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, Ea.

The decomposition reaction of hydrogen iodide,

2HI(g) H2(g) + I2(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, Ea.

ln

!k2

k1

"= !Ea

R

!1

T2! 1

T1

"

Ea = !R lnk2

k1

!1

T2! 1

T1

"!1

Ea = 1.76 x 105 J/mol = 176 kJ/mol

Ea = !(8.314 J/mol K) ln

!1.10" 10!5 L/mol s

9.51" 10!9 L/mol s

" !1

600 K! 1

500 K

"!1

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.

Macroscopic reaction rates that we measure in the lab can be explained microscopically by two models called collision theory and transition state theory.

2. Transition State Theory--model used to describe the energetics and what reactants and products look like in transforming reactants to products.

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.

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!Ea/RT

Fraction of collisionswith sufficient energyfor reaction

Constant related to collision frequency

Fraction of collisionswith proper orientations

k = pZe!Ea/RT

Gas Constant

Temperature (K)

Activation Energy

The reactants are multiplied in the rate law because a product reflects the reality that reactants must collide to transform to products.

A

A

B

B

4 collisions2 x 2 = 4

A

AB

BA

6 collisions

Add another molecule of A

3 x 2 = 6A

A

B

B

A B

9 collisions3 x 3 = 9

A + B C rate = k [A][B] Why Multiplication?

We multiply because the rate depends on the number of collisions (which is found by multiplication)

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.

Results of Collision Theory

Significance of activation energy: only those collisions withenergy equal to, or greater than, Ea can yield products.

Decreasing Ea and/or increasing T enhances the fraction of productive collisions, f increases k and therefore the reaction rate.

increased T increased average speed of particles increased collision frequency increased reaction rate

Reactants must collide with a minimum Ea and with the proper orientation to react.

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-Ea/RT

called the Boltzmanfactor

Because all chemical reactions are reversible there is an activation energy, Ea for both the forward reaction and the reverse reaction.

ProductsReactants

A + B C + Dkforward

kreverse

Activation Energy--Analogy

Ea

Eb-a

ΔH

Ea is a property of the height of the hill (i.e. the chem reaction)

Ea Minimum Energy Needed To Get Over Hill

“The Hill” = Chemical System

A + B C + D

A reaction energy diagram can help us understand activation energy (Ea).

The Arrehenius equation links the macroscopic to the microscopic insight of collision theory.

k = A e!Ea/RT

Fraction of collisionswith sufficient energyfor reaction

Constant related to collision frequency

Fraction of collisionswith proper orientations

k = pZe!Ea/RT

Gas Constant

Temperature (K)

Activation Energy

• Consider the reaction between an atom of chlorine and a molecule of nitrosyl chloride:

Reactants must be collide with the proper spatial orientation in order to transform from reactant to product.

• There are two possible ways that Cl atoms and NOCl molecules can collide; one is effective and one is not.

Cl + NOCl → NO + Cl2

An Effective Collision...Cl collides withCl

Ineffective Collision...Cl collides with O

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

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

∆H

Ea

A....B....C(Transition State)

(Reactants)

A + BC

AB + C(Products)

Reation Progress

Ener

gy

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

∆H

Ea

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

Reaction energy diagrams include the energetics for enthalpies and levels of activation energy.

∆H > 0 ∆H < 0 ∆H < 0

RED for the reaction between CH3Br and OH-

The proposed transition state in the reaction between: CH3Br + OH- ===> CH3OH + Br-

The TS is trigonal bipyramidal; note the elongated C-Brand C-O bonds

A key reaction in the upper atmosphere is

O3(g) + O(g) 2O2(g)

The Ea(fwd) is 19 kJ, and ΔHrxn for the reaction is -392 kJ. Draw a reaction energy diagram for this reaction, a transition state, and calculate Ea(rev).

O3(g) + O(g) 2O2(g)The Ea(fwd) is 19 kJ, and ΔHrxn for the reaction is -392 kJ. Draw a reaction energy diagram for this reaction, a transition state, and calculate Ea(rev).

Ea(rev) = (392 + 19) kJ = 411 kJ

transition stateNot to scale!

ΔHrxn = Hf - Hi = 392kJ

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.

The reaction thermodynamics (enthalpy, entropy) are unaffected!

The catalyzed reaction proceeds via a different reaction mechanism than the uncatalyzed reaction.

Uncatalyzed Pathway Catalyzed

Pathway

No change in the yield of the reaction!

Catalysts lower the activation energy (increasing k) relative to an uncatalyzed reaction.

Both reactions

have the same enthalpy, ∆H

Reaction Progress

Pote

ntia

l Ene

rgy

Uncatalyzed

Catalyzed

Ea Uncatalyzed

Ea Catalyzed

Reactants

Products

Ea No

Catalyst

Ea No

Catalyst

Catalyst does not alter the yield of the

reaction relative to

uncatalyzed pathway.

The reaction pathway

(mechanism) is different in a catalyzed

reaction

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.

In homogeneous catalysis, the reactants and the catalysts are dispersed in a single phase, usually a liquid.

• Haber synthesis of ammonia

• Ostwald process for the production of nitric acid

• Catalytic converters

• Acid catalysis

• Base catalysis

CO + Unburned Hydrocarbons + O2 CO2 + H2Ocatalytic

converter

2NO + 2NO2 2N2 + 3O2catalytic

converter

Catalysts are used in catalytic converters to reduce nitrogen oxides and to clean up unburned HC’s