Chemical Kinetics

Unit 11

Chemical Kinetics Chemical equations do not give us

information on how fast a reaction goes from reactants to products.

KINETICS: the study of reaction rates and their relation to the way the reaction proceeds, i.e. its mechanism

We can use thermodynamics to tell if a reaction is product – or reactant – favored

Only kinetics will tell us how fast the reaction happens!

Rate of Reaction A rate is any change per interval of time.

Example: speed (distance/time) is a rate!

Reaction rate = change in concentration of a reactant or product with time

Expressing a Rate

For the reaction A P

t

ARate

t

PRate

=

Appearance of product

Disappearance of reactant

Reaction Conditions & Rates

Collision Theory of Reactants Reactions occur when molecules collide

to exchange or rearrange atoms Effective collisions occur when

molecules have correct energy and orientation

Factors Affecting Rates

1. Concentrations (and physical state of reactants and products)

2. Temperature

3. Catalysts

Catalysts are substances that speed up a reaction but are unchanged by the reaction

Effect of Concentration on Reaction Rate

To propose a reaction mechanism,

we study the reaction rate and its

concentration dependence.

Rate Laws or Rate Expressions

The rate law for a chemical reaction relates the rate of reaction to the concentration of reactants.

For aA + bB cC + dD

The rate law is: Rate = k[A]m[B]n

The exponents in a rate law must be determined by experiment.They are NOT derived from the stoichiometry coefficients in an overall chemical equation.

Rate Laws & Orders of Reactions

Rate Law for a reaction:

Rate = k[A]m[B]n[C]p

The exponents m, n, and pAre the reaction orderCan be 0, 1, 2, or fractions (may be other

whole numbers in fictional examples)Must be determined by experiment

Overall Order = sum of m, n, and p

Interpreting Rate Laws If m = 1 (1st order)

Rate = k [A]1

If [A] doubles, then the rate doubles (goes up by a factor of 2)

If m = 2 (2nd order) Rate = k [A]2

If [A] doubles, then rate quadruples (increases rate by a factor of 4)

If m = 0 (zero order) Rate = k [A]0 If [A] doubles, rate does not change!

Rate = k[A]m[B]n[C]p

Rate Constant, kRelates rate and concentration at a given temperature.

General formula for units of k: M(1- overall order) time-1

Overall Order Units of k

0 M time-1

1 Time-1

2 M-1 Time-1

3 M-2 Time-1

Rate Law Problem:The initial rate of decomposition of acetaldehyde, CH3CHO, was measured at a series of different concentrations and at a constant temperature.

Using the data below, determine the order of the reaction – that is, determine the value of m in the equation

CH3CHO(g) CH4(g) + CO(g)

Rate = k[CH3CHO]m

CH3CHO (mol/L)

0.162 0.195 0.273 0.410 0.518

Rate (mol/L*min)

3.15 4.56 8.94 20.2 35.2

Strategy

Use the equation:

Pick any two points from the given data!

m

1

2m

1

m2

AA

AA

1 Rate2 Rate

Deriving Rate Laws

Rate of rxn = k[CH3CHO]2

Here the rate goes up by FOUR when the initial concentration doubles.

Therefore, we say this reaction is SECOND order overall.

Example:

Using the same set of data from the previous example, and knowing the order of the reaction, determine:

b) the value of the rate constant, k (w/ units!)

c) the rate of the reaction when

[CH3CHO] = 0.452 mol/L

Strategy: Use any set of data to find k. Solve for rate using k, rate order equation, and given

concentration.

The data below is for the reaction of nitrogen (II) oxide with hydrogen at 800oC.

2NO(g) + 2H2(g) N2(g) + 2H2O(g)

Determine the order of the reaction with respect to both reactants, calculate the value of the rate constant, and determine the rate of formation of product when [NO]=0.0024 M and [H2]=0.0042 M.

Strategy:Choose two experiments where concentration of

one reactant is constant and other is changed; solve for m and n separately!

Example:

The initial rate of a reaction A + B C was measured with the results below. State the rate law, the value of the rate constant, and the rate of reaction when [A] = 0.050 M and [B] = 0.100 M.

Experiment [A] (M) [B] (M) Initial Rate (M/s)

1 0.1 0.1 4.0x10-5

2 0.1 0.2 4.0x10-5

3 0.2 0.1 16.0x10-5

Potential Energy DiagramsMolecules need a minimum amount of energy for a reaction to take place.

Activation energy (Ea) – the minimum amount of energy that the reacting species must possess to undergo a specific reaction

Activated complex - a short-lived molecule formed when reactants collide; it can return to reactants or form products.

Formation depends on the activation energy & the correct geometry (orientation)

Potential Energy Diagram

Potential Energy Diagrams

Potential Energy Diagrams

Catalyzed Pathway

Catalysts lower activation energy!!!

Reaction MechanismsMechanism – how reactants are converted to products at the molecular level

Most reactions DO NOT occur in a single step! They occur as a series of

elementary steps

(a single step in a reaction).

Rate Determining StepRate determining step –

the slowest step in a reaction

COCl2 (g) COCl (g) + Cl (g) fast

Cl (g) + COCl2 (g) COCl (g) + Cl2 (g) slow

2 COCl (g) 2 CO (g) + 2 Cl (g) fast

2 Cl (g) Cl2 (g) fast

Getting the Overall Reaction

COCl2 (g) COCl (g) + Cl (g) fast

Cl (g) + COCl2 (g) COCl (g) + Cl2 (g) slow

2 COCl (g) 2 CO (g) + 2 Cl (g) fast

2 Cl (g) Cl2 (g) fast

2 COCl2 (g) 2 Cl2 (g) + 2 CO (g)

Adding elementary steps gives the

net (or overall) reaction!

Intermediates Intermediates are produced in one

elementary step but reacted in another

NO (g) + O3 (g) NO2 (g) + O2 (g)

NO2 (g) + O (g) NO (g) + O2 (g)

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

Catalysts Catalyst – a reactant in an elementary step

but unchanged at the end of the reactionA substance that speeds up the reaction but is

not permanently changed by the reactionBoth an original reactant and a final product

NO (g) + O3 (g) NO2 (g) + O2 (g)

NO2 (g) + O (g) NO (g) + O2 (g)

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

Example

Cl2 (g) 2 Cl (g) Fast

Cl (g) + CHCl3 (g) CCl3 (g) + HCl (g) Slow

CCl3 (g) + Cl (g) CCl4 (g) Fast

Identify: The rate determining step The overall (net) reaction The identity of any intermediates The identity of any catalysts

Example

H2O2(aq) + I1-(aq) H2O(l) + IO1-(aq) Slow

H2O2(aq) + IO1-(aq) H2O(l) + O2(g) + I1- (aq) Fast

Identify: The rate determining step The overall (net) reaction The identity of any intermediates The identity of any catalysts

Example

O3 (g) + Cl (g) O2 (g) + ClO (g) Slow

ClO (g) + O (g) Cl (g) + O2 (g) Fast

Identify: The rate determining step The overall (net) reaction The identity of any intermediates The identity of any catalysts