Rates of Reaction

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Rates of Reaction

• Chapter16

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RATE CHANGE DURING A REACTION
Reactions are fastest at the start and get
slower as the reactants concentration drops. In
a reaction such as A 2B gt C the
concentrations might change as shown

• Reactants (A and B)
• Concentration decreases with time
• Product (C)
• Concentration increases with time
• the steeper the curve the faster the
• rate of the reaction
• reactions start off quickly because
• of the greater likelihood of collisions
• reactions slow down with time as
• there are fewer reactants to collide

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MEASURING THE RATE
RATE How much concentration changes with time. It
is the equivalent of velocity.
THE SLOPE OF THE GRADIENT OF THE CURVE GETS LESS
AS THE REACTION SLOWS DOWN WITH TIME
CONCENTRATION
y
x
TIME

• the rate of change of concentration is found
  from the slope (gradient) of the curve
• the slope at the start of the reaction will
  give the INITIAL RATE
• the slope gets less (showing the rate is
  slowing down) as the reaction proceeds

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Rate of Reaction

• The rate of a reaction is defined as the change
  in concentration per unit time in any one
  reactant or product.

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Rate of Reaction

• The rate of a reaction is defined as the change
  in concentration per unit time in any one
  reactant or product.
• The rate of reaction depends on 5 factors.
• Nature of reactants
• Particle size
• Concentration
• Temperature
• Catalysts.

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Nature of reactants

• The reaction between acidified sodium dichromate
  and Ammonium Iron (11) sulphate is instantaneous.
  (Ionic)
• The reaction between acidified sodium dichromate
  and ethanal occurs much more slowly. (Covalent)

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Particle Size.

• Large Marble chips
• Small marble chips CaCO3 2HCl ? CaCl2 CO2
  H2O
• Powdered Marble
• The Rate of the reaction increases as the
  particle size decreases.
• The smaller the particle size the greater the
  surface area.
• Therefore the greater the number of collisions,
  the greater
• the number of successful collisions.

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INCREASING SURFACE AREA

• Increasing surface area increases chances of a
  collision - more particles are exposed
• Powdered solids react quicker than larger lumps
• Catalysts (e.g. in catalytic converters) are in
  a finely divided form for this reason
• In many organic reactions there are two liquid
  layers, one aqueous, the other non-aqueous.
  Shaking the mixture improves the reaction rate as
  an emulsion is often formed and the area of the
  boundary layers is increased giving more
  collisions.

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CUT THE SHAPE INTO SMALLER PIECES
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SURFACE AREA 993333 30 sq units
SURFACE AREA 9 x (111111) 54 sq units
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Coal Dust Explosion
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Concentration

• The greater the concentration the greater the
  rate of reaction
• This reaction is studied using the reaction
  between sodium
• thiosulfate and Hydrochloric acid.
• Na2S2O3 2HCl -gt S 2NaCl SO2
  H2O
• If the concentration of the reactants is
  increased, the number
• of collisions will also be increased.
• If the number of collisions is increased then the
  number of
• effective collisions will be increased.

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INCREASING CONCENTRATION
Increasing concentration more frequent
collisions increased rate of reaction
Low concentration fewer collisions
Higher concentration more collisions
However, increasing the concentration of some
reactants can have a greater effect than
increasing others
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Temperature

• The greater the temperature the greater the rate
  of reaction
• This reaction is studied using the reaction
  between sodium
• thiosulfate and Hydrochloric acid.
• Na2S2O3 2HCl -gt S 2NaCl SO2
  H2O
• The rate of a reaction increases as the
  temperature increases because more of the
  colliding molecules have the minimum activation
  energy needed to react.

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INCREASING TEMPERATURE
Effect increasing the temperature increases the
rate of a reaction particles get more energy so
they can overcome the energy barrier particle
speeds also increase so collisions are more
frequent
ENERGY CHANGES DURING A REACTION As a reaction
takes place the enthalpy of the system rises to a
maximum, then falls A minimum amount of energy
is required to overcome the ACTIVATION ENERGY
(Ea). Only those reactants with energy equal to,
or greater than, this value will react. If more
energy is given to the reactants then they are
more likely to react.
Typical energy profile diagram for an exothermic
reaction
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INCREASING TEMPERATURE
MAXWELL-BOLTZMANN DISTRIBUTION OF MOLECULAR ENERGY
NUMBER OF MOLECUES WITH A PARTICULAR ENERGY
MOLECULAR ENERGY
Because of the many collisions taking place
between molecules, there is a spread of molecular
energies and velocities. This has been
demonstrated by experiment. It indicated that
... no particles have zero energy/velocity
some have very low and some have very high
energies/velocities most have intermediate
velocities.
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INCREASING TEMPERATURE
MAXWELL-BOLTZMANN DISTRIBUTION OF MOLECULAR ENERGY
T1
NUMBER OF MOLECUES WITH A PARTICULAR ENERGY
T2
TEMPERATURE T2 gt T1
MOLECULAR ENERGY

• Increasing the temperature alters the
  distribution
• get a shift to higher energies/velocities
• curve gets broader and flatter due to the
  greater spread of values
• area under the curve stays constant - it
  corresponds to the total number of particles

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INCREASING TEMPERATURE
MAXWELL-BOLTZMANN DISTRIBUTION OF MOLECULAR ENERGY
NUMBER OF MOLECUES WITH A PARTICULAR ENERGY
NUMBER OF MOLECULES WITH SUFFICIENT ENERGY TO
OVERCOME THE ENERGY BARRIER
Ea
MOLECULAR ENERGY
ACTIVATION ENERGY - Ea The Activation Energy is
the minimum energy required for a reaction to
take place The area under the curve beyond Ea
corresponds to the number of molecules with
sufficient energy to overcome the energy barrier
and react.
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INCREASING TEMPERATURE
TEMPERATURE T2 gt T1
MAXWELL-BOLTZMANN DISTRIBUTION OF MOLECULAR ENERGY
T1
T2
NUMBER OF MOLECUES WITH A PARTICULAR ENERGY
EXTRA MOLECULES WITH SUFFICIENT ENERGY TO
OVERCOME THE ENERGY BARRIER
Ea
MOLECULAR ENERGY
Explanation increasing the temperature gives more
particles an energy greater than Ea more
reactants are able to overcome the energy barrier
and form products a small rise in temperature can
lead to a large increase in rate
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Catalysts
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Types of Catalysis

• Homogeneous catalysis. This is catalysis in which
  both the reactants and the catalyst are in the
  same phase. (Iodine snake experiment)
• Heterogeneous catalysis. This is catalysis in
  which the reactants and catalyst are in different
  phases. (Hydrogen peroxide (liquid) and Manganese
  dioxide (solid))
• Auto catalysis.One of the products in the
  reaction catalyses the reaction. (Permanganate
  ions and Fe2 ions.)

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Mechanisms of catalysis

• Intermediate compound theory.
• A B ? AB SLOW
• A C ? AC FAST
• AC B? AB C FAST
• The decomposition of hydrogen peroxide catalysed
  by the presence of I ions (iodine snake reaction)
  illustrates the formation of an intermediate.
• Overall Reaction
• 2H2O2 2H2O O2
• Step 1
• H2O2 I- H2O IO-
• Step 2
• H2O2 IO- H2O O2 I-
• Also
• Sodium Hydrogen tartrate Hydrogen peroxide
  Co2 ions (pink)
• Pink? Blue/Green ?Pink
• Cobalt Intermediate
  Cobalt

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Surface Adsorption Theory

• Methanol Methanal
• CH3OH
  HCHO
• Platinum
• 2H2 O2 ? 2H2O
• A good example is the reaction of Hydrogen and
  Oxygen to form
• water using finely divided Platinum as the
  catalyst.
• The Hydrogen and oxygen molecules settle on the
  surface of the
• catalyst. The adsorbed atoms form weak bonds
  with the metal
• atoms. Transition metals can act as catalysts
  because they have
• vacant d orbitals.
• The hydrogen and oxygen molecules then react to
  form water.
• The products leave the surface of the catalyst.
  (desorption)

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Catalytic converters.

• Exhaust fumes contain carbon monoxide (CO),
  nitrogen
• monoxide (NO), nitrogen dioxide (NO2) and unburnt
• hydrocarbons. a catalytic converter converts
  these gases to
• environmentally friendly gases.
• The catalytic converter consists of a thin
  coating of platinum,
• palladium, and rhodium on a ceramic or metal
  honeycomb inside
• a stainless steel case.
• 
  Pt/Pd/Rh (Temp 300 C)
• 2CO 2NO ? 2CO2 N2
• The un-burnt hydrocarbons react with oxides of
  nitrogen to form carbon dioxide nitrogen and
  water.

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Catalytic converter
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Collision Theory

• For a reaction to occur, the reacting particles
  must collide with each other.
• For the formation of product a certain minimum
  energy is required in the collision.
• Such a collision is called an effective
• collision.

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COLLISION THEORY

• Collision theory states that...
• particles must COLLIDE before a reaction can
  take place
• not all collisions lead to a reaction
• reactants must possess at least a minimum amount
  of energy - ACTIVATION ENERGY
• plus
• particles must approach each other in a certain
  relative way - the STERIC EFFECT
• According to collision theory, to increase the
  rate of reaction you therefore need...
• more frequent collisions increase particle
  speed or
• have more particles present
• more successful collisions give particles more
  energy or
• lower the activation energy

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The Activation energy

• The Activation energy is the minimum energy which
  colliding particles must have for a reaction to
  occur.

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Energy profile diagram

• A catalyst works by reducing the activation
  energy.

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ADDING A CATALYST

• Catalysts provide an alternative reaction
  pathway with a lower Activation Energy (Ea)
• Decreasing the Activation Energy means that
  more particles will have sufficient
• energy to overcome the energy barrier and
  react
• Catalysts remain chemically unchanged at the
  end of the reaction.

WITHOUT A CATALYST
WITH A CATALYST
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ADDING A CATALYST

• Catalysts provide an alternative reaction
  pathway with a lower Activation Energy (Ea)
• Decreasing the Activation Energy means that
  more particles will have sufficient
• energy to overcome the energy barrier and
  react
• Catalysts remain chemically unchanged at the
  end of the reaction.

WITHOUT A CATALYST
WITH A CATALYST
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ADDING A CATALYST
MAXWELL-BOLTZMANN DISTRIBUTION OF MOLECULAR ENERGY
NUMBER OF MOLECUES WITH A PARTICULAR ENERGY
NUMBER OF MOLECULES WITH SUFFICIENT ENERGY TO
OVERCOME THE ENERGY BARRIER
Ea
MOLECULAR ENERGY
The area under the curve beyond Ea corresponds to
the number of molecules with sufficient energy to
overcome the energy barrier and react. If a
catalyst is added, the Activation Energy is
lowered - Ea will move to the left.
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ADDING A CATALYST
MAXWELL-BOLTZMANN DISTRIBUTION OF MOLECULAR ENERGY
NUMBER OF MOLECUES WITH A PARTICULAR ENERGY
EXTRA MOLECULES WITH SUFFICIENT ENERGY TO
OVERCOME THE ENERGY BARRIER
Ea
MOLECULAR ENERGY
The area under the curve beyond Ea corresponds to
the number of molecules with sufficient energy to
overcome the energy barrier and react. Lowering
the Activation Energy, Ea, results in a greater
area under the curve after Ea showing that more
molecules have energies in excess of the
Activation Energy
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CATALYSTS - A REVIEW

• work by providing an alternative reaction
  pathway with a lower Activation Energy
• using catalysts avoids the need to supply extra
  heat - safer and cheaper
• catalysts remain chemically unchanged at the
  end of the reaction.
• Types Homogeneous Catalysts
  Heterogeneous Catalysts
• same phase as reactants different phase to
  reactants
• e.g. CFCs and ozone e.g. Fe in
  Haber process
• Uses used in industry especially where an
  increase in temperature results in
• a lower yield due to a shift in equilibrium
  (Haber and Contact Processes)

• CATALYSTS DO NOT AFFECT THE POSITION OF ANY
  EQUILIBRIUM
• but they do affect the rate at which
  equilibrium is attained
• a lot is spent on research into more effective
  catalysts - the savings can be dramatic
• catalysts need to be changed regularly as they
  get poisoned by other chemicals
• catalysts are used in a finely divided state to
  increase the surface area

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Monitoring the rate of production of oxygen from
hydrogen peroxide using manganese dioxide as a
catalyst

• Hydrogen peroxide decomposes into water and
  oxygen as follows
• H2O2(l) ? H2O(l) 1/2 O2(g)
• This occurs much too slowly to be
• monitored. However, manganese
• dioxide acts as a suitable catalyst,
• and the reaction occurs at a
• measurable rate.

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Method 2
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Questions

• 1. Why is the slope of the graph steepest in the
  early stages of the reaction?
• Since rate is proportional to concentration,
  the greatest rate, indicated by the steepest
  slope, is evident in the early stages when the
  concentration of hydrogen peroxide is at a
  maximum.
• 2. At what stage is the reaction complete?
• When the graph becomes horizontal.
  
• 3. What would be the effect on the graph of
  doubling the amount of
• manganese(IV) oxide?
• The increased surface area of catalyst would
  speed up the reaction, giving a steeper slope and
  an earlier completion. The volume of oxygen
  produced would be unchanged.

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• 4. Would doubling the manganese(IV) oxide create
  a practical difficulty? Explain your answer.
• Yes. The production of oxygen could become
  too quick for accurate monitoring.
• 5. What would be the effect on the graph of
  doubling the concentration of
• hydrogen peroxide?
• Increasing the concentration of a reactant
  would speed up the rate, as indicated by a
  steeper slope. Doubling the concentration would
  produce double the final volume of oxygen.
• 6. Would doubling the concentration of hydrogen
  peroxide create a practical difficulty? Explain
  your answer.
• Yes. The capacity of the collection vessel
  could be exceeded.

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Studying the effects on reaction rate of (i)
concentration and (ii) temperature

• The reaction used is that between a sodium
  thiosulfate solution and hydrochloric acid
• 2HCl(aq) Na2S2O3(aq) 2NaCl(aq)
  SO2(aq) S(s)? H2O(l)
• The precipitate of sulfur formed gradually
  obscures a cross marked on paper and placed
  beneath the reaction flask. The rate of reaction,
  and consequently the time taken to obscure the
  cross, depends on a number of variables such as
  temperature, concentration and volume. By varying
  one of these and keeping the others constant, the
  effect on rate can be studied.
• The inverse of the time taken to obscure the
  cross is the measure of reaction rate used in
  this experiment.

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Effect of concentration

• 1. Place 100 cm3 of the sodium thiosulfate
  solution into a conical flask.
• 2. Add 10 cm3 of 3 M hydrochloric acid to the
  flask, while starting the stop clock at the same
  time.
• 3. Swirl the flask and place it on a piece of
  white paper marked with a cross.
• 4. Record the time taken for the cross to
  disappear.
• 5. 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.
• 6. If the initial sodium thiosulfate
  concentration is 0.1 M, subsequent concentrations
  will be 0.08 M, 0.06 M, 0.04 M and 0.02 M
  respectively.

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• 7.Record the results in a table similar to the
  following
• Concentration of thiosulfate 0.l M 0.08 M 0.06
  M 0.04 M 0.02 M
• Reaction time (s)
• 1/time
• 8. Draw a graph of 1/time against concentration.
  This is effectively a graph of
• reaction rate against concentration.

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Effect of temperature

• Procedure NB Wear your safety glasses.
• 1 . Place 100 cm3 of 0.05 M sodium thiosulfate
  solution into a conical flask.
• 2. Warm the flask gently until the temperature is
  about 20 0C.
• 3. Add 5 cm3 of 3 M HCl, starting a stop clock at
  the same time, before proceeding.
• 4. Without delay, swirl the flask, place it on a
  piece of white paper marked with a cross, and
  record the exact temperature of the contents of
  the flask.
• 5. Record the time taken for the cross to
  disappear
• 6. Repeat the experiment, heating the thiosulfate
  to temperatures of approximately
• 30 0C, 40 0C, 50 0C and 60 0C respectively
  (before adding the HCl).
• 7. Record the results in a table similar to the
  following

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Questions

• Suggested Answers to Student Questions
• 1. What is the effect of increasing the
  concentration on the reaction time?
• The reaction time is decreased.
• 2. What is the effect of increasing the
  concentration on the reaction rate?
• The rate is increased.
• 3. What is meant by saying that two quantities
  are directly proportional?
• If one of the quantities is
  increased/decreased by a certain factor, the
  other changes in exactly the same way.
• 4. What is the effect of raising the temperature
  on the reaction time?
• The reaction time is decreased.
• 5. What is the effect of raising the temperature
  on the reaction rate?
• Suggest two factors responsible for the result
  observed.
• The rate is increased. The higher temperature
  results in greater kinetic energy of the
  particles present. This causes
• (i) more collisions per unit time, and
• (ii) a greater proportion of the collisions to
  have the activation energy needed for products
  to form.
• Both (i) and (ii) result in a rate increase.

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• 6. Suggest a reason why it is not recommended to
  carry out the experiment at temperatures higher
  than about 60 0C.
• The reaction occurs so quickly that it is not
  possible to measure the time
• accurately.
• 7. Which is the limiting reactant in the
  temperature experiment?
• 100 cm3 of 0.05 M Na2S2O3 contains
• 100/1000 x 0.05 0.005 moles
  Na2S2O3
• 5 cm3 of 3 M HCl
  contains
• 5/1000 x 3 0.015 moles HCl
• According to the balanced equation, the reacting
  ratio is Na2S2O3 HCl 12
• The amounts used are in the ratio
• Na2S2O3 HCl 0.005 0.015 1 3
• Clearly Na2S2O3 is the limiting
  reactant.