6.1 DETERMINING THE RATE OF A CHEMICAL RXN
What is the rate of a chemical rxn?
The speed at which reactants are used up and products are
formed
EXPERIMENTS TO MEASURE RATE OF RXN
Consider this rxn:
CaCO3(s) + 2HCl(aq)  CaCl2(aq) + CO2(g) + H2O
We can measure the rate of rxn in 2 ways
1. Measure rate that CO2 is produced
2. Measure the rate at which the mass decreases
MEASURING CO2 PRODUCTION
CO2(g)
Delivery
Tube
Measuring
Cylinder
HCl(aq)
CaCO3(s)
Water
MEASUREMENT OF THE RATE MASS DECREASES
Mass decreases as CO2
is given off
Measure mass
decrease every 10 sec
Plot the data
Cotton
Wool
HCl(aq)
CaCO3(s)
Balance
RATE OF RXN DEFINED
Change in concentration of reactants or products
per unit of time
(Time could be 1 sec, 1 min, etc)
Average rate= (change in concentration)/(time)
Unit are dm-3 s-1, dm-3 min-1, etc
If a given amount of reactant is used up, the same
amount will be produced
6.2 COLLISION THEORY
What is the product of particles colliding?
Reactions
For collisions to result, there are two conditions that
must be fulfilled
1. Collision must involve more than a certain minimum
amount of energy
2. Molecules must collide with the correct orientations
COLLISION MUST INVOLVE A CERTAIN MIN.
AMOUNT OF ENERGY
To react, particles must collide with sufficient energy
The min. energy must result in the activation energy (Ea)
Ea  is the minimum amount of energy that colliding
particles must posses for a collision to result in a rxn
If particles collide and do not overcome Ea, they will just
bounce off each other
Collisions resulting in a reaction are called successful or
effective collisions
Not all rxns that overcome Ea will result in rxn
MOLECULES COLLIDE WITH CORRECT ORIENTATION
If molecules do not collide with correct orientation they will not react
Main factors affecting rate of rxn
1. Concentration of reactants
2. Pressure for reactions involving gases
3. Surface area of solid reactants
4. Temperature
5. Catalysis
EFFECTS OF CONC.
Does higher reactant concentrations, how will it
effect the rate?
With more particles in a given volume, there will
be more collisions
Increases the chances of successful collisions
EFFECTS OF PRESSURE
How will increased pressure effect rate?
Similar to concentration increase
Increases collision frequency
Only reactions involving gases are significantly affected by changing
pressure
EFFECT OF SURFACE AREA OF SOLID REACTANTS
Reactions generally only occur on the surface of solids
How can surface area of a solid be increase?
Finely divide the solid to open up more areas for reactions
Allows for more potential collision opportunities
RELATIONSHIP BETWEEN TEMP AND THE
ENERGY OF PARTICLES IN A GAS
How does temp effect the movement of particles?
Ideal gas: kinetic energy of the particles in a gas is
proportional to its temp in K
If temp is doubled, average energy is generally doubled
Does not depend on identity of gas, however lighter
molecules will be traveling faster than heavier ones
EFFECTS OF TEMP ON RATE
How does temp effect rate?
The rate will generally double for every 10 K increase in temp
Why?
Collision frequency increases due to particles moving faster
The also collide harder
Increases the chance for collision
MAXWELL-BOLTZMANN
Distribution of molecular kinetic energies at a particular temp
Only a few particles with high energy and only a few low energy
Most particles have average energy
CATALYSIS
Catalyst A substance that increases the rate of a chemical rxn
without itself being used up in the rxn
Often written above the yield arrow in the equation
Also allow the rxn to proceed by an alternative pathway of lower
activation energy
May be homogeneous (same state as reactants) or heterogeneous
HL2
6.3 THE RATE EXPRESSION
Rate equation/ expression
Consider A B
Rate is directly proportional to [A]
Rate= k[A]
k is the rate constant
The rate expression is experimentally determined equation
relating to the rate of rxn to the concentration of
substances in the rxn mixture
RATE EXPRESSION
General equation
xA + yB  C + D
Rate= k[A]m[B]n
Rate constant a constant of proportionality relating the
concentrations in the experimentally determined rate
expression to the rate of the rxn
Only constant at a particular temp
ORDER OF RXN
In respect to a particular reactant is the power of the
reactant’s concentration in the experimentally
determined rate equation
m and n are the orders of the reactants
Overall order is m+ n
Rate expression can be determined experimentally
EXPERIMENTAL DETERMINATION OF RATE EXPRESSION
Consider, A + B C
The initial rate can be taken because we know the
initial concentrations of A and B
An experiment with a fixed amount of B and varied
concentrations of A is performed
Then do the same with fixed A and varied B
Use data to determine orders of A and B
DETERMINING ORDER OF REACTION AND RATE
EXPRESSION FROM EXPERIMENTAL DATA
Given the reaction 2A B
Experiment
[A]/ mol dm-3
Rate/ mol dm-3
s-1
1
2
3
1.0
2.0
5.0
0.60
1.2
3.0
CONT
We want to determine:
1. The order with respect to A
2. The rate expression
3. The value of the rate constant (w/ units)
4. The rate of reaction when [A]= 1.3 mol dm-3
ZERO-ORDER RXNS
Rate independent of concentration
Rate equation is: rate=k
Units of k are conc. time-1
Could be
mol dm-3 s-1
mol dm-3 h-1
etc.
FIRST –ORDER RXNS
Rate is directly proportional to the
concentration
Half- life is first-order
Half-life is related to rate constant:
rate constant= 0.693/ half-life
Rate equation: rate=k[A]
Units of k are time-1
SECOND-ORDER RNS
Rate of rxn is proportional to
concentration squared
Rate expression: rate=k[A]2
UNITS OF RATE CONSTANT
Overall order Units of k
Example of units
0
Concentration time-1 mol dm-3 s-1
1
time-1
s-1
2
Concentration-1
time-1
mol-1 dm3 s-1
3
Concentration-2
time-1
mol-2 dm6 s-1
6.4 THE ARRHENIUS EQUATION
Shows the variation of the rate constant with temperature
As temp. increases, the rate constant increases exponentially
Equation: k=Ae-Ea/RT
A pre-exponential factor, A-factor or frequency factor
Relates to frequency and orientation of collision
A constant that varies slightly with temperature
e-Ea/RT  the fraction of collision where E> Ea (Energy is greater than
activation energy)
Not all reactions where E> Ea will result in collision
THE OTHER ARRHENIOUS EQUATION
May also be written as lnk= (-Ea/R) x (1/T) + lnA
This form is used to solve for activation energy, if these
procedures are followed
1. Conduct a series of temperature-varied experiments
2. Calculate rate constant for each temp.
3. Plot a graph of lnk (y-axis) vs. 1/T (x-axis)
Temp in K
Slope = -Ea/R
(R= gas constant)
EFFECTS OF A CATALYST ON RATE CONSTANT
If the rate equation is: rate= k[A][B]
Catalyst increases rate constant
6.5 MECHANISMS OF REACTIONS
Consider the reaction,
2NO2(g) +F2(g)  2NO2F(g)
If this rxn were to occur in a single step, all 3 molecules would have to
collide in correct orientation at the same time
We could assume that if the concentration was increased the
chances of proper collision would increase
This makes the rate dependent upon reactant concentrations
Rate equation will be: rate= [NO2]2[F2]
The 2 superscript comes from the coefficient in the balanced
equation
CON’T
The rate derived from experimentation was found to be
rate= [NO2][F2]
This suggests that the rxn does not occur in a single step
Thus, this reaction (as many more) must occur in multiple
steps
The chances of molecules colliding in perfect orientation
simultaneously is quite low
SUGGESTEDMECHANISMS OF THIS RXN
NO2 + F2  NO2F +F
Step 1
NO2 + F  NO2F
Step 2
2NO2 + F2  2NO2F
Overall Equation
Step 1:
rate= k1[NO2][F2]
Step 2:
rate= k2[NO2][F]
CON’T
Step 1 is the same as the overall equation, thus must be the step
determining the rate for the overall rxn
Called the rate-determining step (the slow step)
Step 2 is the fast step and does not influence the overall rate of rxn
to a great extent
thus, the concentrations of the step species are not included in the
rate equation
GENERIC EXAMPLE
A + 2B  C
B+BQ
Step 1
rate- determining step
Q+AC
Step 2
fast
Mechanism
B+BQ
Q+AC
2B + A  C
Thus, rate= k[B]2
since it is the only species remaining from the first
step
RATE DETERMINING STEP AS SECOND STEP
B + B  Q
Q+AC
Step 1
Step 2
fast
rate-determining step
Process of determining overall rate is basically the same as if the first step
as the slow step
Rate= k[Q][A]
Because Q is produced by 2 B molecules in the first step, we can replace [Q]
with [B]2
Rate= k[B]2[A]
ANOTHER MECHANISM
A + B  S
step 1
fast
S+BC
step 2
rate-determining step
Reactants involved up to and including the rate determining step are
included in rate equation
Rate= k[B]2[A]
REACTION INVOLVING A CATALYST
CH3COCH3(aq) + I2(aq)  CH3COCH2I(aq) + HI(aq)
The rxn is acid (H+) catalysed
Experimental rate expression is
Rate= k[CH3COCH3][H+]
Does not include I2, so it is only involved after the rate-determining
step
CON’T
Proposed mechanism:
CH3COCH3 + H+  X
X + I2  CH3COCH2I + HI + H+
rate-determining step
fast
Catalyst is involved in the rate- determining step but is regenerated in
second step and does not appear in the overall chemical equation
X is an intermediate
H+ cancels out
SN1 VS SN2 MECHANISMS
(CH3)3CBr + OH-  (CH3)3COH + Br –
This is a nucleophilic substitution
Experimental rate expression:
rate= k[(CH3)3CBr]
Since OH- is not included, it is in the fast step
Suggested mechanism:
(CH3)3CBr  (CH3)3C+ + Br –
rate- determining step
(CH3)3C+ + OH-  (CH3)3COH
fast
CON’T
This is considered SN1
S= substitution
N= nucelophilic
1= molecularity of rate- determining step
Molecularity # of ‘molecules’ that react in a particular step (usually
rate-determining)
SUMMARY OF MECHANISM RULES
1. Mechanism must agree with overall stoichiometric equation
2. Maximum of 2 particles can react in any one step
3. All species in rate equ. Must appear in mechanism in or before
the rate-determining step
4. The power of a particular reactant’s concentration in the rate equ.
Indicates the # of times it appears in the mechanism up to and
including the rate determining step