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Chemical Kinetics
• Factors that Affect Reaction rates
• Reaction Rates
• Concentration and Rate
• The Change of Concentration with Time
• Temperature and Rate
• Reactions Mechanisms
• Catalysis
Chemical Kinetics
Is the study of the rate at which reactions occur and also gives us information on how the reaction occurs (the Reaction Mechanism)
Factors that Affect Reaction Rates
Reaction rates depend on several factors
Physical State of the Reactants and Mixing
The concentration of the Reactants
The Temperature and Pressure at which the
reaction occurs
Catalysts
On a molecular level Reaction rates depend on the frequency with which molecules collide
The greater the frequency of collisions (with enough energy to break bonds), the faster the rate of the reaction
A quantitative definition of the rate of a chemical reaction is defined in terms of product(s) forming and reactant(s) disappearing per unit time
N2(g) + 3H2(g) → 2NH3(g)
Reaction Rates
Rate of ammonia formation can be expressed as
[NH3]t2 - [NH3]t1 = Δ[NH3] = 0.50M - 0Mt2 - t1 Δt 25s - 0s
= 0.50M / 25s = 0.02 M/s
(Rates are expressed as positive quantities, units: M or mol/L per second)
This is the average rate, it doesn’t gives us an actual rate at a given moment in time
Gives information on the rate at a particular moment, for this we plot the concentration of product or reactant with time and determine the slope at our time of interest
Instantaneous Rate
C4H9Cl(aq) + H2O(l) → C4H9OH(aq) + HCl(aq)
Reaction Rates and Stoichiometry
2HI(g) → H2(g) + I2(g)
In this case the rate of appearance (formation) of H2 and I2 is equal, but 2 mols of HI are consumed for every 1 mol of H2 and I2 formed, so we can express this as:Rate = - 1 Δ[HI] = Δ[H2] = Δ[I2]
2 Δt Δt Δt
The rate of HI disappearance is twice the rate of H2 and I2 appearance
This leads to the generalization that in a given reaction
aA + bB → cC + dD
Rate = -1 Δ[A] = -1 Δ[B] = 1 Δ[C] = 1 Δ[D]a Δt b Δt c Δt d Δt
In the reaction 2O3(g) → 3O2(g). O2 is formed at 2.0 x 10-5 M/s at a given instant, at what rate is O3 disappearing at this instant
Concentration and Rate Rate = k[A]m[B]n
NH4+(aq) + NO2
−(aq) N2(g) + 2 H2O(l)
The rate is proportional to the concentration of both reactants, doubling either
the concentration of NH4+ or NO2
- doubles the rate of the reaction, so we say
that the reaction order for NH4+ and NO2
- is 1 : we can express the rate law as
Rate = k [NH4+][NO2
-] m and n are both 1
For the Reaction: A + B → C + D
The experimentally data was tabulated as shown, Write the rate law for the reaction
Expt [A] M [B] M Init.Rate M s-1
1 0.1 0.1 0.001
2 0.1 0.2 0.002
3 0.2 0.1 0.004
What are the reaction orders with respect to reactants A and B and what is the order of the reaction overall?
The reaction orders can be found with the generalized formula:
Reaction order = {Log (rate 2/rate 1)} / {Log (concn 2/concn1)}
e.g. For NO2- = Log (2) / Log(2) = 1
From a rate law we can calculate the rate of reaction using the rate constant and initial reactant concentrations. We now need an equation that allows us to determine the concentration of reactants and products at any particular time
Zeroth Order Reactions
12.4 The Change in concentration with Time
Consider the reaction: A → products
Rate = - Δ [A] = k [A]o = k Differential Rate LawΔt
Intergrating this differential rate law gives:
[A]t - [A]o = -kt Ao is the initial concentration of A (at t = 0)At is the concentration of A at any time t after
From a rate law we can calculate the rate of reaction using the rate constant and initial reactant concentrations. We now need an equation that allows us to determine the concentration of reactants and products at any particular time
First Order Reactions
12.4 The Change in concentration with Time
Consider the reaction: A → products
Rate = - Δ [A] = k [A] Differential Rate LawΔt
Intergrating this differential rate law gives:
ln [A]t = -kt Ao is the initial concentration of A (at t = 0)[A]o At is the concentration of A at any time t after
Since ln[A]t/[A]o = ln[A]t - ln[A]o
We can get this equation in the form y = mx +c by re-arranging:
ln[A]t = -kt + ln[A]o
Plotting ln[A]t against t should give a straight line with slope = -k,and intercept = ln[A]o
Second Order Reactions
Again Consider the reaction: A → products
Rate = - Δ [A] =k [A]2 Intergrate to give:Δt
1 = kt + 1[A]t [A]o
So plotting 1/[A]t vs. t will be linear for a second order reaction
Half-Life
[A]t = 0.5[A]o
ln[A] t = -kt½ln[A]o
ln 0.5 = -kt½
= 0.693 = t½
k
First Order Reactions
Temperature and Rate
Kinetics Review
aA + bB → cC + dD Rate (r) = -1 Δ[A] = -1 Δ[B] = 1 Δ[C] = 1 Δ[D]a Δt b Δt c Δt d Δt
r = k[A]m[B]n
Zeroth order First order
Second orderArrhenius equation
[A]t - [A]o = -kt ln[A]t = -kt + ln[A]or= k r= k[A]
r= k[A]21 = kt + 1[A]t [A]o ln k = -Ea/RT + ln A
ln 0.5 = -kt½
t½ = 1k[A]o
k = Ae-Ea/RT R = 8.314 J/ mol K (T absolute)
Arrhenius Equation
ln k = -Ea/RT + ln A
Reaction Mechanisms
It provides an alternative pathway (mechanism), it alters either A or Ea or both
Catalysis
A catalyst increases the rate of a reaction (by lowering the Activation Energy), without being consumed in the reaction
Homogeneous Catalyst
Present in the samephase as the reactants
Heterogeneous Catalyst
Present in a different phase as the reactants (often metals)
Heterogeneous Catalysts
Exist in a different Phase to the reactants
For example catalytic hydrogentation of alkenes with a Ni catalyst and Catalytic Converters.
Reactants are adsorbed onto the surface of the metalReactants are free to move around the surfaceReactions happen (new sigma bonds form) on the surface
The products desorb
Enzyme Catalysis
Substrate (yellow) binds to enzyme (purple) in an active site
Each enzyme catalyses a specific reaction in the same way a key fits a given lock
The product is produced and the enzyme is unchanged.
An enzyme can increase a reaction rate up to 1018 times!
The product(s) leaving the active site is the rate determining step, once the products have left the site can be filled by another substrate molecule
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