instructor: dr. marinella sandros

Instructor: Instructor:

Dr. Marinella SandrosDr. Marinella Sandros

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NanochemistrNanochemistryy

NAN 601NAN 601

Reaction mechanisms and Catalysis

“A sequential series of simple reactions which combine to form a larger, balanced chemical equation.”

For the overall reactionNO2 + CO CO2 + NO

The elementary reactions of the mechanism are

2NO2 NO + NO3 R1 = k1[NO2]2

molecularity = 2

NO3 + CO NO2 + CO2 R2 = k2[NO3][CO]molecularity = 2

For any elementary reaction◦ the order of the reaction wrt a reactant is its stoichiometric

coefficient in that step◦ the molecularity of the reaction is the sum of the stoichiometric

coefficients for that step. The molecularity corresponds to the number of molecules that actually collide in that step!

◦ NO3 is an intermediate.

Rule #1For any step, the order of that step w.r.t. a reactant is its stoichiometric coefficient in that step.For step 2 in the previous reaction

NO3 + CO NO2 + CO2 R2 = k2[NO3][CO]

NO2 + CO CO2 + NO

Two elementary reactions:2NO2 NO + NO3 (slow) R1 =

k1[NO2]2

NO3 + CO NO2 + CO2 (fast) R2 = k2[NO3][CO]

Overall reaction:NO2 + CO CO2 + NO Rate =

R1 =k1[NO2]2

The rate law for the overall reaction is the rate law for the

rate-limiting step!

Rule #2For a multistep reaction in which one step is much slower than the others,

overall rate = rate of slowest step

For any step, the order of that step w.r.t. a reactant is its stoichiometric coefficient in that step.

For a multi-step reaction in which one step is much slower than the others, the overall rate is equal to the rate of the slowest step

Write the rate equation for the overall reaction in terms of the reactants in the net reaction.

Step 1: NO (g) + NO (g) N2O2 (g) (fast)

Step 2: N2O2 (g) + Br2(g) 2 NOBr (g) (slow)

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K1

K-1

k2

2 NO (g) + Br2(g) 2NOBr (g)

Experimentally determined rate law:rate= k [NO]2[Br]

Show the following mechanism also produces a rate law consistent with experimentally

observed one?

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Step 1: NO (g) + NO (g) N2O2 (g) (fast)

Step 2: N2O2 (g) + Br2(g) 2 NOBr (g) (slow)

K1

K-1

k2

Step 2 is rate limiting!!!

Rate= k2 [N2O2][Br2]

K1[NO]2 = k-1 [N2O2 ]

K1[NO]2 = [N2O2 ]k-1

Rate= k2 K1[NO]2 [Br2] = k [NO]2[Br2]

k-1

A catalyst (Greek: καταλύτης, catalytēs) is a substance that accelerates the rate of a chemical reaction without itself being transformed or consumed by the reaction. (wikepedia)

A + B

C

ΔG

Ea

uncatalyzed

A + B +catalyst

C + catalyst

ΔG

Ea′

catalyzed

k(T) = k0e-Ea/RT

Ea′ < Ea

k0′ > k0

k′ > k

ΔG = ΔG

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Catalyst ??Efficiency depends on activity, properties & life of

the catalyst

• Examples:

• Ammonia synthesis – Promoted iron

• SO2 oxidation – Venadium Pentaoxide

• Cracking – Sylica, alumina

• Dehydrogenation – Platinum, Molybdenum

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A reaction may have a large, negative ΔGrxn, but the rate may be so slow that there is no evidence of it occurring.

Kinetic Vs. Thermodynamic

Conversion of graphite to diamonds is a thermodynamic favor process (ΔG -ve ).

C (graphite) --> C (diamond)

Kinetics makes this reaction nearly impossible (Requires a very high pressure and temperature over long time)

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Activation Energy

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Catalyst lowers the activation energy for both forward and reverse reactions.

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This means , the catalyst changes the reaction path by lowering its activation energy and consequently the catalyst increases the rate of reaction.

A catalyst that is present in the same phase as the reacting molecules.

Example:

The Decomposition of aqueous hydrogen peroxide

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

Very very very slow!!!!!!

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Catalyzed by Bromine:

Br2 (aq) + H2O2 (aq) 2 Br- (aq) + 2 H+ (aq) + O2 (g)

2 Br- (aq) + H2O2 (aq) + 2 H+ (aq) Br2 (l) + 2H2O (l)

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The two reactions together serve as a catalytic pathway for hydrogen peroxide decompostion.

Both of them must have significantly lower activation energies than the uncatalyzed decomposition.

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The catalyst exists in a different phase from the reactant molecules, usually as a solid in contact with either gaseous reactants or with reactants in a liquid solution.

Many industrially important reactions are catalyzed by the surfaces of the solids.

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Sulfuric acid synthesis (Contact process)

SO2 + O2, SO3 vanadium oxideshydration of SO3 gives H2SO4

Ammonia synthesis (Haber-Bosch process)

N2 + H2, NH3 iron oxides on aluminaconsumes 1% of world's industrial energy budget

Nitric acid synthesis (Ostwald process)

NH3 + O2, HNO3unsupported Pt-Rh gauze

direct routes from N2 are uneconomical

Hydrogen production by Steam reforming

CH4 + H2O, H2 + CO2 Nickel or K2OGreener routes to H2 by water splitting actively sought

Ethylene oxide synthesis

C2H4 + O2, C2H4Osilver on alumina, with many promotors

poorly applicable to other alkenes

Hydrogen cyanide synthesis (Andrussov oxidation)

NH3 + O2 + CH4, HCN Pt-RhRelated ammoxidation process converts hydrocarbons to nitriles

Olefin polymerization Ziegler-Natta polymerization

propylene, polypropylene

TiCl3 on MgCl2

many variations exist, including some homogeneous examples

Desulfurization of petroleum (hydrodesulfurization)

H2 + R2S (idealized organosulfur impurity), RH + H2S

Mo-Co on aluminaproduces low-sulfur hydrocarbons, sulfur recovered via the Claus

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http://en.wikipedia.org/wiki/Heterogeneous_catalysis

Heterogeneous catalysts are often composed of metals or metal oxides.

Initial step is usually Adsorption of reactants.

Adsorption binding of molecules to a surface, whereas absorption refers to the uptake of molecules into the interior of another substance.

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H H H H H H H H H H H H H H H H H H

H

H

HH

H

HH

H

H

H

H

H

H

H

H

H

HH

H2 absorption onpalladium hydride

H2 adsorption onpalladium

Adsorption occurs because the atoms or ions at the surface of a solid are extremely

reactive.

Example of heterogeneous catalysis is the reaction of hydrogen with ethylene:

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C2H4 + H2 C2H6 VERY SLOW!

However in the presence of finely powdered metal such as nickel or

palladium at room temperature and under <200 atm of hydrogen pressure.

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http://chemwiki.ucdavis.edu/@api/deki/files/1559/=Catalytic_Hydrogenation_Mechanism.jpg

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In physisorption

1. The bond is a van der Waals interaction

2. Adsorption energy is typically 5-10 kJ/mol. ( much weaker than a typical chemical bond )

3. Many layers of adsorbed molecules may be formed.

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For Chemisorption

1. The adsorption energy is comparable to the energy of a chemical bond.

2. The molecule may chemisorp intact (left) or it may dissociate (right).

3. The chemisorption energy is 30-70 kJ/mol for molecules and 100-400 kJ/mol for atoms.

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Adsorbent: surface onto which adsorption can occur.example: catalyst surface, activated carbon, alumina

Adsorbate: molecules or atoms that adsorb onto the substrate.example: nitrogen, hydrogen, carbon monoxide, water

Adsorption: the process by which a molecule or atom adsorb onto a surface of substrate.Coverage: a measure of the extent of adsorption of a species onto a surface

H H H H H H H H Hadsorbate

adsorbent

H H H H H

coverage θfraction of surface sites occupied

http://www.biology.hawaii.edu/171L/SummerSession1/Lab%204%20Enzyme%20Catalysis.pdf34

2NO (g) + 2N2O(g) 2N2(g) + 2NO2(g)

2 NO2(g) 2NO(g) + O2 (g)

a)What is the chemical equation for the overall equation?

b)Why is NO considered a catalyst and not an intermediate?

c) Is this an example of homogeneous or heterogeneous catalysis?

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(a) 2 N2O (g) 2N2(g) + O2(g)

(b) An intermediate is produced and then consumed. A catalyst is consumed but then reproduced. NO2 is the intermediate.

(c) Since NO is in the same state as reactant, it is homogeneous.

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instructor: dr. marinella sandros

Documents

overall rate

multistep reaction

rate equation

g n2o2 g faststep

net reaction

chemical reaction

reaction mechanisms

k2no3cooverall reaction