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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 - PowerPoint PPT Presentation
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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