Transcript Rates of Reaction

Rates of Reaction
Objectives
• To understand that a chemical
reaction involves collisions between
particles
• To be able to describe the factors
which will affect the rate of a
chemical reaction.
Some Everyday Chemical
Reactions
• Burning wood
• Fruit ripening
• Getting a tan
• Cooking food
How Reactions Happen
• A chemical reaction
involves a collision
between particles.
• The particles collide and
make new substances
• The particles which react
are called the reactants
• The substances which
are made are called the
products
Reactions Happen at
Different Speeds
• There are chemical reactions that occur
very slowly and others that occur very
quickly
RUST FORMATION
FIREWORKS
Rates of Reaction
• The rate of a reaction is how quickly a
reaction proceeds
• It may be defined as the change in
concentration in unit time of any one
reactant or product
Rate = Change in concentration (molL-1)
Time taken (s)
Calculate the rate of the
following
H2O2 → H2O + ½ O2
The initial concentration of the reaction
is 5 molL-1, ten seconds later this has
decreased 3.5 molL-1
What is the rate?
Calculate the rate of the
following
2SO2 + O2 → 2SO3
The initial concentration of the reaction
is 15 molL-1, 20 seconds later this has
decreased 12.5 molL-1
What is the rate?
Average and Instantaneous
Rate
• The average rate of reaction is average
rate over the course of the reaction
• The instantaneous rate of reaction is
the rate at a particular point in time
during the reaction
Controlling the Rates of
Reaction
• Being able to control
the speed (rate) of
chemical reactions is
important both in
everyday life
(cooking) and when
making new materials
on an industrial
scale.
Factors Affecting Rate of
Reaction
1. Concentration of reactants
2. Temperature of reaction
3. Particle size of solid reaction
4. Nature of reactants
5. Presence of catalyst
1. Concentration of
Reactants
• The
higher
the
concentration
of
reactants the higher
the probability of
collisions
between
reactant molecules
There are less red
particles in the
same volume so
there is less chance
of a collision
There are more red
particles in the same
volume so there is
more chance of a
collision so the
reaction goes faster
2. Temperature of
Reaction
• An increase in
temperature brings
about an increase in
reaction rate.
• You give more
energy to the
system in the form
of heat
3. Particle Size
• Particle size (finely
divided particles
react faster)
• Molecules can only
collide at the
surface.
• Smaller particles
bigger surface
area
Dust Explosions
• A dust explosion is
the explosive
combustion of a dust
suspended in air in
an enclosed location
• Any solid material
that can burn in air
will do so at a rate
that increases with
increased surface
area
Grain Dust Peril
• In 1998, at a series of
explosions occurred at a grain
elevator facility in Haysville,
Kansas
• There were seven fatalities as a
result of the explosions.
• It is not the actual grain that is
ignited, but the fine, thick dust
which is released during the
loading process when grain
particles rub against each other
4. Nature of Reactants
• Ionic compounds react faster than covalent
• In reactions bonds are broken and form
• When an ionic compound is placed in water it
dissociates
• It takes more energy to break covalent bonds
Activation Energy (Eact)
• For a reaction to happen, energy is required.
• Activation energy is the minimum energy
with which particles need to collide to cause
a reaction
• This is different
according to the
type of bonds
of the reactants.
• The activation energy may be shown on a
reaction profile diagram (right)
• These diagrams show energy as a barrier that
needs to be overcome by the reactants before
they become products
• The difference between the energy of
the reactants and the energy of the
products is the heat of reaction (ΔH)
Endothermic reaction
• A reaction in which
heat is taken in.
• In an endothermic
reaction heat is
taken in from the
surroundings and the
products formed
have more energy
than the reactants.
• It is written as +
ΔH
Exothermic Reaction
• A reaction in which
heat is liberated.
• In an exothermic
reaction heat is lost
to the surroundings
and the products
formed have less
energy than the
reactants.
• It is written as - ΔH
5. Catalysts
• A catalyst is a substance that alters the
rate of a chemical reaction but is not
consumed (used up) in the reaction.
• In most cases it makes the reaction go
faster. Some catalysts make a reaction
go slower and are called negative
catalysts or inhibitors. Eg Calcium
propionate added to bread to make it
stay fresh longer (ie. It slows down
staleness)
• The catalyst does not get used up in the
reaction.
• It gives the reaction the energy to get
started
General Properties of
catalysts
1. Catalysts are recovered chemically
unchanged at the end of a reaction
(Eg. Manganese dioxide used to speed up
decomposition of Hydrogen peroxide
has exactly the same chemical
properties before and after the
reaction)
2. Catalysts tend to be specific, even
though a catalyst may catalyse one
reaction it may not have any effect on
a similar reaction
Enzymes in the body are examples of
catalysts that are very specific
Know two examples
• Protease breaks down proteins such as
blood stains on clothes and are used in
washing powders
• Catalase breaks down hydrogen
peroxide in the body
3. Catalysts need only be present in very
small amounts
• Increasing the amount of catalyst does
not greatly affect the rate of a
reaction and in cases where it does it is
usually something to do with the
reaction itself
4. Catalysts help reactions reach
equilibrium quicker but do not change
what the equilibrium of a reaction is
• 5. Action of catalysts may be destroyed
by catalytic poisons for example lead in
petrol can destroy the catalytic
converters in cars
Arsenic is a poison that inhibits the
action of certain enzymes in the body
Types of Catalysis
• Chemists have discovered 3 types of
catalysis
1. Homogenous Catalysis
2. Heterogeneous Catalysis
3. Autocatalysis
Homogenous Catalysis
• This describes when reactants and
catalysts are in the same phase and
there is no boundary between them
• Eg Iodine Snake where Potassium
Iodide catalyses the decomposition of
H2O2 to release oxygen both catalyst
+ Reactant are liquids
Heterogeneous Catalysis
• Catalysis where the catalyst and
reactants are in different phases Eg. a
liquid and a solid
• There is a boundary between the
catalyst and the reactants
• Eg. MnO2 catalyses the decomposition
of H2O2 to release oxygen both
catalyst + Reactant are liquids
Autocatalysis
• When one of the products of the
reaction catalyses the reaction
• In this type of reaction it occurs slowly
at first but as the reaction proceeds it
gets quicker this is because the
products drive the reaction forward.
Theories of Catalysts
• Speed up a
reaction by giving
the reaction a new
path.
• The new path has
a lower activation
energy and more
molecules have
this energy.
• The reaction goes
faster.
Think of a catalyst as a
tunnel through a mountain
CATALYST
By lowering the activation energy
a catalyst makes it possible to
carry out a reaction at lower
temperatures (lower energy)
Mechanisms of Catalysis
• The mechanism of catalysis tells us how
the catalyst works
• There are two main mechanism of
catalysis for you to study
1. Intermediate Formation theory
2. Surface Adsorption theory
Intermediate Formation Theory
of Catalysts
• Homogeneous catalysts sometimes work
by reacting with reactants to form
unstable intermediate products
• The intermediate exists for a very
short time and reacts with the other
reactant to give the final product and
regenerate the catalyst
X+A + B → C
AX + B
See it with you own eyes
• Oxidation of Potassium Sodium Tartrate
using hydrogen peroxide
• Catalyst in this reaction is Cobalt ions
which give a pink colour
• The intermediate is a green colour
which appears as carbon dioxide + steam
are given off
• The pink colour is restored at the end
indicating the Cobalt ions have not been
used up
Surface Absorption Theory
of Catalysts
• Heterogenous catalysis of gas reactions
by metals
• The reaction happens on the surface of
the metal
• The reaction occurs at the active site
of the catalyst
• A catalyst can have multiple active sites
Stages of reactions of ethene
H
H H
H C
C H
H
H
H
CATALYST
Reactants get absorbed onto
catalyst surface. Bonds are
weakened
H
H H
H
H H
C C
H
CATALYST
H
Bonds Break
H
H
H
H
C
C
H
H
H
H
CATALYST
New bonds formed
H
H
H
H
C
C
H
H
H
CATALYST
H
Second bond forms and
product diffuses away from
catalyst surface, leaving it
absorb fresh reactants
H
H
H C
H
H
C H
H
H
CATALYST
Catalytic Poisons
•
•
•
Catalysts can be poisoned they can become less efficient and
sometimes they no longer work at all
In heterogeneous catalysis particles that poison the catalyst
(lead / arsenic) are absorbed more strongly onto the catalyst
surface than the reactant particles
Catalytic poisons block the active sites of enzymes
How it works (Catalytic
Converter)
• The catalyst in the
converter speeds up
reactions that
reduce atmospheric
pollution
• The catalyst remains
unchanged at the
end of the reaction
• The catalyst is a mix
of transition metals
(platinum, rhodium,
palladium)
Reactions Catalysed
• Carbon monoxide is converted to carbon
dioxide by reaction with oxygen
CO + ½ O2 → CO2
• Carbon monoxide can react with
nitrogen monoxide to give carbon
dioxide
2CO + 2NO → 2CO2 + N2
• Unburnt hydrocarbons are oxidised to
carbon dioxide and water
C8H18 + 12½ O2 → 8CO2 + 9H2O
Environmental Benefits
• Reduction in emissions of toxic gases including
unburnt hydrocarbons
• Reductions in photochemical smog
Smog in Beijing
Mandatory Experiment 14.1
Monitoring the rate of production of
oxygen from hydrogen peroxide, using
manganese dioxide as a catalyst
Hydrogen peroxide→Oxygen + Water
How to layout your results
T
0 15
Vol 0 60
30
45
60
90
110
134 150 170 184 190 197 197
O2
Cm3
60
15
Time secs
120 150 180 210
oxygen production by decomposition of hydrogen peroxide
70
60
Oxygen/cm3
50
When the reaction is
complete the amount of
oxygen stops increasing
40
30
As the reaction proceeds
more oxygen is produced
20
10
0
0
2
4
6
Time/mins
8
10
12
Average and Instantaneous Rates
• The instantaneous rate of reaction is the rate
at a particular point in time during the
reaction
• You use your graph to find the instantaneous
rate
oxygen production by decomposition of hydrogen peroxide
70
60
Oxygen/cm3
50
Draw a tangent to the curve
40
30
20
10
0
0
2
4
6
Time/mins
8
10
12
oxygen production by decomposition of hydrogen peroxide
70
60
V2
50
Oxygen/cm3
V1
Connect the tangent to the X and
Y axis using straight lines
40
30
20
10
0
0
t1
2
t2
4
6
Time/mins
8
10
12
Calculations
Instantaneous rate= tan θ= Δv = v2 – v1
(Slope)
Δt t2 – t1
= 52-45 (cm3)
3 - 2 (min)
= 7 (cm3)
1 (min)
instantaneous rate = 7cm3 min-1
Mandatory Experiment 14.2
Studying the effects on the reaction rate
of (i) concentration and (ii)
temperature, using sodium thiosulfate
solution and hydrochloric acid
What’s happening?
2HCl(aq) + Na2S2O3(aq)
→
2NaCl(aq) + SO2(aq) + S(s)↓ + H2O(l)
Place 100 cm3 of the sodium thiosulfate
solution into a conical flask.
Add to the flask, while starting the stop
clock at the same time.
• Repeat the experiment using 80, 60, 40 and 20
cm3 of. sodium thiosulfate solution
respectively. In each case, add water to make
the volume up to 100 cm3 and mix before
adding HCl.
Concentration of
Thiosulphate
0.1 M
0.08 M
0.06 M
0.04 M
0.02 M
Reaction time
(s)
Rate of Reaction
(1/time)
(s-1 )
Draw Graph
1
time
Concentration of Thiosulphate