EEA Grants Norway Grants Introduction to kinetics and

EEA Grants Norway Grants Introduction to kinetics and catalysis Ing. Marcela Králová Ph. D. , CEITEC 20. 4. 2015

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

Kinetics • 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

Measuring the reaction rate • 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

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 + 3 H 2 2 NH 3

Rate laws • • 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 + 3 H 2 2 NH 3

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

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

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

First 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

First order reactions • 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

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

Second order reactions 2 A → P [k] = dm 3 mol-1 s-1

Second order reactions • 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 Charact kinetic plot Scope of kinetic plot Units of rate constant Zero [A] vs t -k mol dm-3 s First ln[A] vs t -k s-1 k dm 3 mol-1 s-1 Reaction order Second Differntial rate law Integrated rate low 1/[A] vs t -1

Determination of rate low from experimental data • 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. A 0(N-1) • two experiments with different c. A 0 • receive two half time subtracted

Determination of rate low from experimental data • Differential method: • Initial concentration same for all reactants • c. A 0 = c. B 0 = c. C 0 = 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. A 0(1) = 2 c. A 0(2) • v 0 = k c 0 N divide

Collision theory • 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 + B → C • 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 Activation energy • 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 Molecular orientation and Effective collision • 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

Theory of transition state • 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

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

Catalysis • 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)

Homogeneous 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

Heterogeneous 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

Photocatalysis • Absorption of UV light • Creation of e- and h+ REDUCTION: e- + O 2 → ●O 2●O - + ●OOR 2 → 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

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

Thank you for your attention • This project is funded by the Norwegian Financial Mechanism. Registration number: NF-CZ 07 -ICP-1 -040 -2014. Name of the project: „Formation of research surrounding for young researchers in the field of advanced materials for catalysis and bioapplications“