Introduction to kinetics and catalysis Ing. Marcela Králová Ph.D., CEITEC EEA Grants Norway Grants

Kinetics Reaction lows Reactions orders and its determination Theory of chemical reactions Homogeneous catalysis Heterogeneous catalysis Photocatalysis Content

Deal with the rates of chemical processes Chemical processes – sequence of one or more single step Elementary process – transition between two atomic/molecular state separated by a potential barrier Activation energy Low barrier = fast reaction High barrier = slow reaction Elementary reactions Single reactive collision (bimolecular step) Dissociations/isomerisation (unimol.step) Termolecular step Goals of the study: Reaction mechanisms Absolute reaction rate Kinetics

Gas syringe method: For gas reaction Gas is collected in the syringe Push out against the plunger The volume can be read on the syringe Volume can be converted to a change in concentration Measuring the reaction rate

Changes in mass: For gas reaction Calculation of mass loss Gas escapes from the reaction flask Mass of precipitation: For reaction where the precipitation is formed Using stopwatch Measuring the reaction rate

Reaction rate: Rate at which reactants are used up Products are formed Units: concentration per time (mol.dm -3.s -1 ) N 2 + 3H 2  2NH 3 Reaction rate

Differential rate low: Changes of reaction rate with the concentration Reaction rate is proportional to the rate of conc. changes Rate is proportional to derivative of concentration Integrated rate low: Relates the concentration to time Rate laws

Differential rate law (Guldberg-Waag low): v = k[A] a [B] b [C] c k …………………… rate constant powers………….. partial order of the reaction with respect to the reactant overall order….. sum of the powers N 2 + 3H 2  2NH 3 Rate laws

Rate is independent on the concentration of the reactant Examples: some photochemical enzymatic catalyzed reactions reverse Haber process: Zero-order reactions 2NH 3 (g)  3H 2 (g) + N 2 (g)

Zero-order reactions A  P [k] = moldm -3 s -1

Half-life  1/2 : Required time for half of the reactants to be depleted Zero-order reactions

Rate is dependent on the concentration of one reactant Other reactant can be present, but each will be zero-order Examples: A  P A+B  P; where one component is in excess First order reactions

A → P [k] = s -1

Whenever the concentration of a reactant falls off exponentially, the kinetics follow the first order First order reactions

Half-life  1/2 : First order reactions

A + B → P Excess of one reactant Concentration of the other reactants can be include in rate constant First order reactions

Rate is dependent on the concentration of one second-order reactant (2A  P) two first order reactants (A+B  P) Second order reactions

2A → P [k] = dm 3 mol -1 s -1

Whenever the reciprocal of the concentration versus time is linear, the kinetics follow the second order Second order reactions

Half-life  1/2 : Second order reactions

A + B → P [k] = dm 3 mol -1 s -1

Summary Reaction order Differntial rate law Integrated rate low Charact kinetic plot Scope of kinetic plot Units of rate constant Zero[A] vs t-k moldm -3  s -1 Firstln[A] vs t-ks -1 Second1/[A] vs tk dm 3 mol -1  s -1

Overall reaction order: not deduced from chemical equation determined experimentally concentration measurement of one or more reactants Determination of rate low from experimental data

Integral method: Concentration as a function of time Comparison the time dependence LINEAR

Determination of rate low from experimental data Half lives: Only for reaction where is dependence:  1/2 = k/c A0 (N-1) two experiments with different c A0 receive two half time subtracted

Determination of rate low from experimental data Differential method: Initial concentration same for all reactants c A0 = c B0 = c C0 = c 0 v 0 = kc 0 N divide

Determination of rate low from experimental data Isolation method: Determination of partial reaction order Different initial concentration of same reactant c A0(1) = 2c A0(2) v 0 = kc 0 N divide

Explain: How chemical reaction occur Why reaction rate is different for different reaction Criteria: Sufficient kinetic energy (activation energy) Proper orientation Sufficient collision Collision theory

A and B are gasses Frequency of collision is proportional to the concentration of A and B Doubling of c A, the frequency of A-B collision double The rate at witch molecules collide affect the overall reaction rate Collision theory A + B → C

A and B are gasses Reactant sufficient kinetic energy to break the chemical bonds Reactants bonds are broken Products bonds are formed Reactants must be moving enough The minimum energy with which molecules must be moving is called activation energy Rate increases with the temperature Collision theory Activation energy

Sufficient activation energy not garantee succesfull collision Necessity of right orientation Molecules in liquid or gas – constant, random motion – probability of collision Effective collision - one in which molecules collide with sufficient energy and proper orientation, so that a reaction occur Collision theory Molecular orientation and Effective collision

Postulate the existence of hypothetical transition state It occurs between reactants and products state Formed species is called activated complex Based upon collision theory Theory of transition state

Activated complex: Reactant-product hybrid Exist at the peak of the reaction coordinate Transition state Theory of transition state

Factors determines if the reaction occur or not: Concentration of the activated complex The rate at which the activated complex breaks apart The mechanism by which the activated complex breaks apart Back to reactants Towards products Theory of transition state

Collision theory: Successful collision Enough energy Proper orientation Theory of transition state: Successful collision Enough energy Proper orientation Collision theory versus Theory of transition state PRODUCTS ACTIVATED COMPLEX

Catalysts: Reduced activation energy Increase the reaction rate Do not change during the reaction Affect the kinetics Do not affect the equilibrium state Homogeneous catalysis: Same phase as a reactants (g; l) Heterogeneous catalysis: Catalysts (s) and reactants (g; l) Catalysis

Examples: Acid catalysis Organometallic catalysis Enzymatic catalysis Advantage: Mix into the reaction mixture High degree of interaction: catalyst-reactant Disadvantage: Irrecoverable after the reaction Homogeneous catalysis

Mechanisms: Diffusion of reactants to surface of catalyst Adsorption of reactants onto the surface at active sites Interaction between the reactant and catalysts surface Chemical reaction Desorption of products Diffusion of products Heterogeneous catalysis

Advantage: Separation form the reaction mixture Disadvantage: Saturation of catalyst surface Heterogeneous catalysis

Absorption of UV light Creation of e - and h + Photocatalysis REDUCTION:͘ e - + O 2 → ● O 2 - ● O ● OOR → unstable products → CO 2 + H 2 O OXIDATION:͘ H 2 O + h + → ● OH + H + ● OH + org. molecule + O 2 → ● OOR → → → → → CO 2 + H 2 O

TiO 2 : High activity; chemical biological inertness Photostable; nontoxicity High recombination; absorption only in UV Application: Self-cleaning Depollution De-odorizing Photocatalysis

Thank you for your attention This project is funded by the Norwegian Financial Mechanism. Registration number: NF-CZ07-ICP Name of the project: „Formation of research surrounding for young researchers in the field of advanced materials for catalysis and bioapplications“