CHEM1612 - Pharmacy Week 12: Kinetics – Catalysis

CHEM1612 - PharmacyWeek 12: Kinetics – Catalysis

Dr. Siegbert SchmidSchool of Chemistry, Rm 223Phone: 9351 4196E-mail: [email protected]

Unless otherwise stated, all images in this file have been reproduced from:

Blackman, Bottle, Schmid, Mocerino and Wille,     Chemistry, John Wiley & Sons Australia, Ltd. 2008

     ISBN: 9 78047081 0866

Lecture 33 - 3

Energy Landscape in Chemical Reactions

A + B

C + D

Activatedstate

Ea (forw)

Ea

(rev)

Exothermic reaction Endothermic reaction

A + B

C + D

Ea (forw)

Ea

(rev)

A + B C + D

Forward reaction is faster than reverse Reverse reaction is faster than forward

Larger Ea smaller klower rate

Figure from S

ilberberg, “Chem

istry”,

McG

raw H

ill, 2006.

Lecture 33 - 4

By the way: why the product?

elementary reaction A + B→ C

rate = k [A][B]

Why does rate depend on the product of reactant concentrations?

Rate proportional to the number of collisions of A and B

No. collisions = product of the number or particles present

2×2 =4

3×2 =6

3×3=9

Lecture 33 - 5

Transition State If reactants come together with enough energy and the right

orientation, they combine to form a transition state (or activated complex).

This species is half-way betweenthe reactants and the productsbut is not neither. Transition states are very unstable (cannot be isolated).

Blackman Figure 14.10

Lecture 33 - 6

Nature of the transition state in the reaction between CH3Br and OH-.

CH3Br + OH- CH3OH + Br -

transition state or activated complex

Transition StateFigure from

Silberberg, “C

hemistry”,

McG

raw H

ill, 2006.

Lecture 33 - 7

Reaction Energy DiagramFigure from

Silberberg, “C

hemistry”,

McG

raw H

ill, 2006.

Lecture 33 - 8

Transition state in elementary steps

Ea1 > Ea2, therefore Ea1 is the slow step and Ea2 is the fast step.

Two transition states.

Blackman Figure 14.11

k = A e – Ea / R T

Lecture 33 - 9

Transition state in elementary steps

Step 1 NO2 + F2 →NO2F + F

Step 2 NO2 + F → NO2F

Overall 2 NO2 + F2 → 2NO2F

Figure from S

ilberberg, “Chem

istry”,

McG

raw H

ill, 2006.

Lecture 33 - 10

A catalyst increases the rate of a chemical reaction.

A catalyst is not consumed or changed in the overall process.

A catalyst provides an alternative reaction pathway of lower activation energy: more molecules have the minimum energy required for successful reaction and the reaction proceeds at a faster rate.

Catalysis

A + B C + Dk

Lecture 33 - 11

Catalysis A catalyst speeds both the forward and reverse reaction, so

does NOT affect the position of the equilibrium.

Does not change the equilibrium constant Keq = k1/k-1; even though k1 and k-1 may be much larger for the catalyzed reaction.

Does not change the G0 for the reaction. A catalyst can be homogeneous (one phase with reactants

and products) or heterogeneous (more than one). Many catalysed reactions are zero-order. A small quantity of catalyst affects the reaction rate for a large

amount of reactant.

Lecture 33 - 12

Demo: Catalytic Decomposition of Hydrogen PeroxideManganese dioxide MnO2 is used to catalyse the decomposition of H2O2.

Prepare a slurry of MnO2 and NaOH, and add H2O2.

The rapid decomposition of H2O2 occurs with production of a grey foam:

2 H2O2 → 2 H2O + O2

Lecture 33 - 13

Let’s look at the decomposition of H2O2 again.

2 H2O2(aq) 2 H2O(l) + O2(g)

Catalyst Ea (kJ mol-1) Rel. rate of reaction

None 75.3 1I– 56.5 2.0 x 103

Pt 49.0 4.1 x 104

Catalase 8.0 6.3 x 1011

Bombardier beetle (Brachinus fumans)

Catalysis

Lecture 33 - 14

Let’s closely examine the reaction of H2O2 with I–:

Rate law: rate = k[H2O2][I–].

Reaction occurs in two steps:

Step 1:H2O2 + I– H2O + IO–

Step 2: H2O2 + IO– H2O + O2(g) + I–

Note that I– regenerated during reaction and it does not appear in overall reaction.

I– acts as a homogeneous catalyst for H2O2 decomposition.

Homogeneous Catalysis

Lecture 33 - 15

A B

Heterogeneous catalysis: most important industrially.

Catalytic converters:

First converter (A): RhCatalyses: 2 NO(g) N2(g) + O2(g)

Second converter (B):Pt/Pd Catalyses: 2 CO(g) + O2(g) 2

CO2(g)

Heterogeneous Catalysis

Catalytic converter

Lecture 33 - 16

The metal-catalyzed formation of ammonia FeN2(g) + 3H2(g) → 2NH3(g)

Both substrates must bind to a free active site on the Fe surface before the reaction can proceed.

Increasing the concentration of either gas cannot increase the rate of reaction (i.e. rate independent of concentration).

Lecture 33 - 17

Enzymes Catalysts of

biological reactions Complex 3D

structure Huge molar mass Active site attracts

substrates through intermolecular forces

Haber process (500 atm and 450 °C; Nitrogenase (1 atm and 25°C)

Enzyme-substrate complex of elastase and small peptide

Lecture 33 - 18

All enzymes are proteins, but not all proteins are enzymes.

Enzymes must possess catalytic activity.

The part of the enzyme tertiary structure that is responsible for the catalytic activity is known as the “active site”.

Active site

Structure of the enzyme Hexokinase from X-ray data.

Each enzyme catalyses a single chemical reaction on a specific substrate molecule with high selectivity.

Enzymatic Catalysis

Lecture 33 - 19

Enzymatic Catalysis

Michaelis-Menten mechanism: enzyme-substrate complex ES.

Product + enzyme

Efficient: rate enhancements of 108 to 1020 possible Specific (one enzyme per reaction) Low tolerance to temperature and pH changes Lock-and-key model (E. Fischer, 1894) Induced fit model (D. Koshland, 1958)

Lecture 33 - 20

Hexokinase alters its conformation to fit around the substrate molecule (D-glucose).

Enzyme and substrate adapt to accommodate one another.

“Enzymes are molecules that are complementary in structure to the transition states of the reactions they catalyze”.

Linus Pauling (1948)

Enzymatic Catalysis

Lecture 33 - 21

Enzymes can distinguish between enantiomers:

Only one of the enantiomers can be used as a substrate for this enzyme.

Enzymatic Catalysis

Lecture 33 - 22

uncatalyzed reactioncatalyzed reaction

Ea(uncat) – E a(cat) = EaFree energy

Reaction co-ordinate

If Ea = 10 kJ mol-1 55-fold rate acceleration (at 25°C).

If Ea = 20 kJ mol-1 3000-fold rate acceleration (at 25°C).

If Ea = 40 kJ mol-1 107-fold rate acceleration (at 25°C).

…..

Enzymatic Catalysis

Lecture 33 - 23

The Arrhenius equation indicates that in order to increase the rate of a reaction:

The temperature must be increased,

Ea must be decreased, and/or

The reactants must be positioned so as to maximise the reaction efficiency.

Increasing the temperature is not an option for most biological reactions, so the remaining options are exploited by Nature.

Enzymatic Catalysisk = A e – Ea / R T

Lecture 33 - 24

SummaryCONCEPTS Elementary reactions, reaction mechanisms Dependence of reaction rate on temperature and orientation Arrhenius equation and its implications Activation energy and transition states Catalysis

CALCULATIONS Express reaction rate in terms of reactant/product

concentrations Derive rate law of a reaction from experimental data on

reactant consumption/product formation Derive rate law of a reaction, knowing its elementary steps Calculate k, A, T or Ea for a reaction using Arrhenius equation

Lecture 33 - 25

In the atmosphere:

O2 + hn (<200 nm) O + O k1 O + O2 O3 assume very fast O3 + hn (210-300 nm) O2 + O k2 O3 + O 2 O2 k3

Kinetics very complicated since UV intensity will vary so much in time and place.

At equilibrium, rate of ozone creation and destruction will be the same: steady state approximation.

The Ozone System

Lecture 33 - 26

Chlorofluorocarbons provide an additional pathway for ozone decomposition.

CF2Cl2 + hn CF2Cl• + Cl• Cl• + O3 ClO• + O2 removing ozone ClO• + O• Cl• + O2

Cl is regenerated during the reaction Cl will stick around until eventually reacts to HCl and is

precipitated out.

The Ozone Hole