Transcript Kinetics

International
Baccalaureate Chemistry
Topic 6 - Kinetics
Review: Reaction Types

Sythesis:
A + B → AB
2Na + Cl2 → 2NaCl

Decomposition:
AB → A + B
2H2O → 2H2 + O2

S. Replacement:
A+ + BC → B+ + AC
Cu+ + AgNO3 → Ag+ + Cu(NO3)2

D. Replacement:
AB + CD → AD + BC
KI + Pb(NO3)2 → KNO3 + PbI2

Combustion:
CxHy + O2 → CO2 + H2O
CH4 + 2O2 → CO2 + 2H2O

Neutralization:
HX + MOH → MX + H2O
HCl + NaOH → NaCl + H2O

Redox:
Still to come
Kinetics: The Study of
Reaction Rates






A Chemical reaction is the process in which chemical
species react to form new substances.
Not all chemical reactions take place at the same rates.
Some reactions are very fast (explosions) and some are
very slow (a car rusting).
Reaction Rate: A change in the concentration of reactants
or products per unit of time
R = Conc. of reactant used
(D conc)
Time interval
(D time)
= Conc. of Product Formed
Time interval
Units are usually Mol/s = mol/L•s = mol•dm-3•s-1
Other units are possible if properties other than
concentration are used.
Kinetics


Chemical reactions occur at different
rates.
Chemical reaction that occur in nature:
– Car rusting (Slow)
– Engine Combustion (Fast)
– Oxidation of copper on pennies (Slow)
– Food spoiling at high temperatures (Fast)
– Graphite → Diamond (Slow)
Factors Affecting
Reaction Rates
1. Nature of Reactants
2. Surface Area:
Dust vs. solid
3. Concentration:
0.01 M vs 3 M
Zn + HCl →
4. Temperature:
Ice melting in hot vs cold water
5. Catalyst
H2O2 decomposition normal vs with MnO2
Example

If 3.0 mol dm-3 of HCl are used up
after 15 s in a certain reaction with Zn
metal, find the average rate of
reaction.

3.0M = 0.2 M
15 s
1s
Reaction Rates
Rates of reactions can be determined by
monitoring the change in concentration of
either reactants or products as a function of
time. D[A] vs Dt
7
Dependents


The rate of a chemical reaction is a measure
of the rate at which reactants are consumed
which is equal to the rate at which products
are formed:
Rate = -D[R] = D[P]
Dt
Dt
Where:
– [R] and [P] are the concentration of the
reactants and products respectively
– Dt is a change in time
General Rate of Reaction
aA+bB→cC+dD
Rate of reaction = rate of disappearance of reactants
=-
1
Δ[A]
a
Δt
1 Δ[B]
=-
b
Δt
= rate of appearance of products
=
1
Δ[C]
c
Δt
=
1
Δ[D]
d
Δt
Example



MnO4-(aq) + 8H+(aq) + 5Fe2+(aq)  Mn2+(aq) + 4H2O(l) + 5Fe3+(aq)
Rate = -D[MnO4-] = 1 D[Fe3+]
Dt
5 Dt
The rate of MnO4- consumption is equal to 1/5 the
rate of Fe3+ production.
Any property that differs between reactants and
products can be used to measure reaction rate.
–
–
–
–
–
Absorption of color or light
Volume, mass pressure of a gas
pH
Electrical conductivity
Temperature
Example

Rate of change of concentration with time.
2 Fe3+(aq) + Sn2+ → 2 Fe2+(aq) + Sn4+(aq)
t = 38.5 s
Δt = 38.5 s
[Fe2+] = 0.0010 M
Δ[Fe2+] = (0.0010 – 0) M
2+
Rate of formation of Fe2+= Δ[Fe ]
Δt
= 0.0010 M
38.5 s
= 2.6x10-5 M s-1
Graphs


In all cases, a graphical plot of the
measured property vs. time will allow
determination of the slope of the graph,
which will be equal to the reaction rate.
Example:
Cu(s) + 4HNO3(aq)  Cu(NO3)2(aq) + 2H2O(l) + 2NO2(g) + heat

The rate of the above equation can be
found by measuring several different
properties:
Possible Properties to Measure
1.
Color change:
Rate = DColor Intensity
DTime
2.
Temperature Change:
Rate = DTemperature
DTime
3.
Pressure Change:
Rate = D Pressure
DTime
4.
Mass Change:
Rate = DMass
DTime
For the initial rate, draw a gradient (Tangent) at t = 0.
These rates will be greater than those taken later on in the reaction
(t=x)
Determining and Using an Initial Rate of Reaction.
H2O2(aq) → H2O(l) + ½
O2(g)
-(-2.32 M / 1360 s) = 1.7 x 10-3 M s-1
-(-1.7 M / 2600 s) =
6 x 10-4 M s-1
Rate =
-Δ[H2O2]
Δt
Reaction Rates
C4H9Cl(aq) + H2O(l)  C4H9OH(aq) + HCl(aq)
[C4H9Cl] M
In this reaction, the
concentration of
butyl chloride,
C4H9Cl, was
measured at various
times, t.
Reaction Rates
C4H9Cl(aq) + H2O(l)  C4H9OH(aq) + HCl(aq)
Average Rate, M/s
The average rate of
the reaction over
each interval is the
change in
concentration divided
by the change in
time:
Reaction Rates
C4H9Cl(aq) + H2O(l)  C4H9OH(aq) + HCl(aq)


Note that the average
rate decreases as the
reaction proceeds.
This is because as the
reaction goes forward,
there are fewer
collisions between
reactant molecules
(because there are
fewer molecules).
Collision Theory



Chemical reactions involve collisions of reactant
particles.
Not all collisions lead to a chemical reaction.
For molecules to react (effective collisions), they
must:
– Collide with each other
– Collide with correct orientation/geometry (Steric factor)
– Collide with sufficient Kinetic Energy (meet the minimum activation
energy (Ea) requirement)


Ineffective collisions involve particles that rebound
essentially unchanged.
The rate of reaction depends upon the frequency of
collisions and fractions of those collisions that are
effective.
Collision Theory

Kinetic-Molecular theory can be used to
calculate the collision frequency.
– In gases 1030 collisions per second.
– If each collision produced a reaction, the rate
would be about 106 M s-1.
– Actual rates are on the order of 104 M s-1.
 Still a very rapid rate.
– Only a fraction of collisions yield a reaction.
Collision Theory
Furthermore, molecules must collide with the
correct orientation and with enough energy
to cause bonds to break and new bonds to
form
Factors Affecting
Reaction Rate





Nature of Reactants
Surface Area
Affect Collision Rate
Concentration
Temperature
Affect proportion with required Ea
Catalyst
Nature of Reactants


Generally, chemical reactions which involve a lot of old
bond breaking and new bond forming are slow.
Example 1:
5Fe2+(aq) + MnO4-(aq) + 8H+(aq)  5Fe3+(aq) + Mn2+(aq) + 4H2O(l)

Example 2:
5C2O42+(aq) + 2MnO4-(aq) + 16H+(aq)  10CO2(g) + 2Mn2+(aq) + 8H2O(l)

Which reaction would occur faster?
– In both reactions, MnO4-(aq) produces Mn2+(aq)
– Change is accomplished by two different reactions.

Using Fe2+(aq), there are no bonds to break, thus the reaction would be faster.

In C2O42-(aq), there are lots of bonds to break, thus the reaction would be
slower.
Physical States


Another aspect to consider with respect to the
nature of the reactants is the physical state the
reactants are in (solid, liquid, gas, aqueous).
Example:
–
–
–
–
Solid + Gas → Slow
Liquid + gas → Slow
Gas + gas → Fast
Aqueous + Aqueous → Very fast
Surface Area


The more surface area exposed to or in
contact with the reactants, the faster the
reaction will take place.
Smaller particles have more available surface
area than larger particles.
Example:

Surface Area only applies to heterogeneous
reactions (ones involving different states)

Example 1:
Wood + O2(g) → Products + Heat

Example 2:
Mg(s) + 2HCl(aq) → H2(aq) + MgCl2(aq)

Example 3:
2NaCl(aq) + Cu(OH)2(aq) → 2NaOH(aq) + CuCl2(aq)

How would you increase the reaction rate of
the above three reactions?
Concentration


If the concentration of a reactant is increased, we can
expect the reaction rate to increase because more
effective collisions may occur due to the presence of
more reactant particles in a given volume.
NOTE: Reducing volume has the same effect as adding
more reactant.
2M H2
6M H2
Temperature

A temperature increase always increases the rate of
a reaction. (IB Question: State)

Reason: There will be more collisions between

Recall: For reactant particles to react and form

reacting particles at higher temperatures because
they are moving faster. (IB Question: Explain)
products, they must collide with sufficient kinetic
energy to break old bonds (i.e. collisions must be
successful).
The minimum amount of energy required for a
successful collision is called its activation energy
(Ea).
Maxwell–Boltzmann Distributions

Temperature is
defined as a
measure of the
average kinetic
energy of the
molecules in a
sample.
• At any temperature there is a wide
distribution of kinetic energies.
28
Maxwell–Boltzmann Distributions


As the temperature
increases, the curve
flattens and broadens.
Thus at higher
temperatures, a larger
population of
molecules has higher
energy.
If the dotted line represents the activation energy, as the
temperature increases, so does the fraction of molecules that can
overcome the activation energy barrier.
As a result, the reaction rate increases.
29
29.
Activation Energies
•How does the change in Temperature affect
the reaction rate? (possible exam question)
•A temperature increase increases the number
of particles with the necessary activation
energy and therefore increases the reaction
rate. (Full mark answer)
•Activation Energy (Ea) is the minimum energy
required for a successful collision, assuming
optimum collision geometry.
Catalysts


A catalyst is a substance that increases the reaction
rate without apparently being consumed in the
chemical reaction.
Example: The decomposition of hydrogen peroxide
MnO2


2H2O2  2H2O + O2
The rate of this reaction is increased by adding the
catalyst MnO2.
How do Catalysts work?
– They lower the Ea so that a greater percentage of
molecules will have sufficient energy to react (E  Ea).
Catalysts
The catalyst is H+
Potential Energy Change


To help explain Catalysts, we can also look at the
Potential Energy Change in a chemical reaction.
All reactions take place in Three Steps:
1. Reactant particles approach each other. As they approach
they begin to slow down because of electron repulsive
forces. This means K.E. is decreasing and Potential
Energy (P.E.) is increasing. If they have sufficient
energy (E  Ea) they will:
2. Collide. This group of atoms formed on collisions is called
the activated complex. If the molecules have enough
activation energy they will react on collision.
3. Once product molecules form, they will separate because
of repulsive forces. As they separate, KE will increase and
PE will decrease.
Reaction Mechanisms
•
The molecularity of a process tells how many
molecules are involved in the process.
•
The rate law for an elementary step is written
directly from that step.
Reaction Mechanisms



4HBr(g) + O2(g) 2H2O(g) + 2Br2 (g)
For this reaction to proceed as written, a
five-particle collision would have to occur.
Chances of this happening are low and
therefore we assume that this reaction, like
most, occurs through a series of steps called
a “Reaction Mechanism”.
A reaction mechanism is a sequence of
simple 2-particle reactions (usually) by
which a complex reaction takes place.
The Steps
Step 1
Step 2
Step 3
HBr + O2
 HOOBr
(Slow)
Intermediate
HOOBr + HBr  2 HOBr
(Fast)
Intermediate
HOBr + HBr
 H2O + Br2 (Fast)
Step 4 HOBr + HBr 
Net Reaction: 4 HBr(g) + O2(g)
H2O + Br2

(Fast)
2H2O(g) + 2Br2(g)
Characteristics of a Reaction Mechanism
1.
2.
3.
4.
5.
Each step in the mechanism is usually bimolecular (2 Particle)
and never more than trimolecular (3 Particle).
The sum of the steps in the reaction mechanism must give the
balanced equation for the chemical reaction. (Intermediates
must be consumed)
It is always the slowest step in a series of steps which
determine the rate of the overall reaction. (called the Rate
Determining Step or R.D.S.)
The only way to increase the rate of the reaction is to speed
up the slowest step. In the above reaction, this can be done by
increasing the [HBr] and /or [O2] because these are the two
reactants involved in the first step of the reaction sequence.
If a reactant is not in the Rate Determining Step, then
changing its concentration will have no effect on the reaction
rate.