Rates

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Murray Low Udhir Chathuri
Rates Extents of Reactions
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Overview
Quiz A
What prevents reaction from going?
Increasing the Rate of Reaction
Temperature
Quiz B
Catalysis
Collisions
Rates Extents of Reactions
Measuring Change
Industry
Temperature
Quiz C
Quiz G
Pressure
Concn.
Mass of System
Establishing Equilibrium
The Equilibrium Constant
Le Chateliers Principle
Dynamic Eqm.
Quiz D
Quiz E
Quiz F
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Quiz A- Common misconceptions

• When solid copper carbonate reacts with excess
  acid, carbon dioxide is produced.
• The curves shown were obtained under two
  different sets of conditions.

X
Y
The change from X to Y could be brought about by
A. increasing the concentration of the acid.
B. decreasing the particle size of copper
carbonate.
C. adding a catalyst.
D. increasing the temperature.
E. decreasing the mass of copper carbonate.
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Quiz A- Common misconceptions
2. Chemical reactions are in a state of dynamic
equilibrium only when
A. the rate of the reverse reaction is equal to
the rate of the forward reaction.
B. the reaction involves a zero enthalpy change.
C. the activation energy of the forward reaction
equals the activation energy of the reverse
reaction.
D. the reaction goes to completion.
E. the ratio of the products to reactants is
equal to exactly 1.
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Quiz A- Common misconceptions

• X(g) is placed in a flask and the following
  reaction proceeds to equilibrium.
• X(g) ? Y(g) ?H ve
• Which of the following statements is correct?

A. The forward reaction rate increases as the
reaction gets going.
B. The forward reaction rate always equals the
reverse reaction rate.
C. The activation energy of the forward reaction
of the system considered will always be higher
than the activation energy of the reverse
reaction no matter the reaction conditions.
D. When equilibrium is re-established after a
disturbance (e.g. adding more Y) the rate of
the forward and reverse reactions will be equal
to those at the initial equilibrium.
E. The forward reaction is completed before the
reverse reaction begins.
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Quiz A- Common misconceptions

• The graph below shows the variation of
  concentration of a reactant , X, with time
• as the reaction proceeds.

What is the average reaction rate (mol dm-3 s-1)
during the first 30 seconds?
A. 0.0005
B. 0.005
C. 0.015
D. 0.020
E. The slope of the tangent to the curve at t
30 s
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Quiz A- Common misconceptions

• Excess marble chips (CaCO3) were added to 50 cm3
  of 1 M HCl. The experiment
• was repeated using the same mass of marbles
  chips and 50 cm3 of 1M CH3COOH.
• Which of the following would have been the same
  for both experiments?

A. The average rate of reaction.
B. The rate at which the first 2 cm3 of gas was
evolved.
C. The time taken for the reaction to be complete.
D. The mass of the marble chips remaining after
the reaction had stopped.
E. None of the above.
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What prevents reactions from going?
Ea
Energy Profile
2NO ? N2 O2
2NO
?H
N2 O2
Collisions
Reaction
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Increasing the Rate of Reaction
Orders of reactants
Rate constant
Non-examinable ! But useful for understanding.
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Increasing the Rate of Reaction
A. Temperature
Everyone knows that as we increase the
temperature, the reaction rate goes up
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T1
What makes reactions go faster?
T2
Maxwell Distribution of velocities
KEmin ½mv2min
KEmin gt Ea
KE ? T
T2 gt T1
vmin
T2
T1
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Increasing the Rate of Reaction
If we increase the temperature we will increase
the fraction of the molecules with the energy to
overcome the activation energy.
Where f fraction of molecular collisions with
energy greater or equal to Ea
If T ? then f ?
If f ? then k ?
If k ? then Rate?
Note If Ea ?due to a catalyst then f ? and the
Rate ?
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Increasing the Rate of Reaction
B. Catalysis
? Catalysts are employed to speedup the
attainment of equilibrium.
? Catalysts do not change the position of
equilibrium .
? Catalysts offer a different mechanism for the
reaction to occur .

• For the catalyst to be successful it must offer a
  mechanism with a
• lower activation energy, Ea.

• Catalysts can either be in the same phase as the
  reactant
• homogeneous catalysts -or a different phase
  heterogeneous
• catalysts.

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Increasing the Rate of Reaction
Ea
Energy Profile
N2(g) O2(g) ? 2NO(g)
Ea
2NO
?H
N2 O2
Catalyzed
Non catalyzed
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Increasing the Rate of Reaction
C. Increasing collisions
The more collisions there are the more likely
reactants are to react.
We can influence the amount of collisions by
1. Increasing the concentration of reactants
(pressure in the case of gases).
2. Increasing the surface area of solid reactants
(or solid catalysts).
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Quiz B- Increasing reaction rates
1. A catalyst lowers the activation energy of
a reaction from 100 kJ/mol to 50 kJ/mol.
What increase in f will result at 25 oC ?

• The activation energy for a reaction is known to
  be 50 kJ/mol of reactant. If the
• temperature is increased from 300 K to
  310 K what increase in the rate of reaction
• would you expect?

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Measuring Change
Consider the following reaction
Q
How could we measure the rate of this reaction?
A
Well lets consider what might change during this
reaction.
Mass of system
Temperature
Pressure
Concentration
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Measuring Change
A. Temperature
If the reaction is exothermic or endothermic then
the temperature will change.
But if the temperature changes then the rate of
reaction will change! Why?
Therefore if we want to measure the rate of
reaction we normally keep the temperature of the
reactor constant. This is called thermostatting
the reactor.
See Maxwell Distribution
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B. Pressure
Measuring Change
We know from Grade 11 that
Therefore if we keep the volume of the
thermostatted reactor constant then
const n
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Measuring Change
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Measuring Change
C. Mass of the system
This one is easy if the reactor is a closed
system then the mass cannot change. Mass must be
conserved.
If the reaction is taking place in an open system
then gases might leave the system and the mass
of the system will decrease.
In this particular case we have a closed system.
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Measuring Change
D. Concentration
This is by far the most common means of measuring
rates, because concentrations ALWAYS change
during a reaction. In this example
We need to be able to measure concentrations and
the most used technique is SPECTROSCOPY where a
reactant or product selectively absorbs (or
emits) electromagnetic radiation, which might
leave it coloured.
The intensity or amount of the radiation
absorbed (A) is proportional to the
concentration.
Normally chemists measure concentrations to
determine reaction rates.
A substance does not need to be coloured to
absorb electromagnetic radiation. CO2 for
instance is colourless as it does not absorb
visible radiation but it does absorb infra-red
radiation (the cause of global warming!).
Therefore we could use infra-red spectroscopy to
monitor CO2.
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Quiz C- Measuring change
1. Which of the following reactions could we
use PTOTAL for the measurement of the
reaction rate?
a) 4NH3(g) 5O2(g) ? 4NO(g) 6H2O(g)
b) O3(g) NO(g) ? NO2(g) O2(g)
c) CaO(s) 3C(s) ? CaC2(l) CO(g)
2. Predict the necessary experimental
conditions to be able to measure the rate of
the following reaction by monitoring the mass
of the system.
CaO(s) 3C(s) ? CaC2(l) CO(g)

• Sulfur dioxide gas can be oxidized to sulfur
  trioxide gas by oxygen gas. In a
• particular experiment a stoichiometric amount of
  sulfur dioxide and oxygen
• are allowed to react in a glass vessel. If
  the initial pressure in the flask is 3 atm ,
• what would be the predicted final pressure if
  the reaction goes to completion
• and the temperature remains constant.

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Quiz C- Measuring change
4. Calculate the rate of CO2 production (moles of
CO 2 / dm3 sec) from an industrial process
given the data in the two graphs shown below.
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Quiz C- Measuring change
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Establishing Equilibrium
Closed System
Open System
(v)
?
(l )
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Establishing Equilibrium
Solution A at t 0 sec
3
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Dynamic Equilibrium
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Quiz D- Establishing Equilibrium

• Which of the following systems are at equilibrium
  ?(Assuming they have been left long
• enough to establish equilibrium.)

a) A boiling kettle.
b) A sealed thermostatted test-tube containing a
drop of water.
c) A completely flat battery.
d) A water tank under the following conditions
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The Equilibrium Constant
The Law of Mass action, which is independent of
kinetic theory, states that for a reaction-
aA bB ? cC dD
the ratio
will be a constant when the system is at
equilibrium. This constant is known as the
equilibrium constant, Kc.
One must remember that the value of Kc is
specific to a particular reaction equation(i.e.
one in which the stoichiometric factors are
fixed) and that it is specific to a given
temperature.
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The Equilibrium Constant
For a gaseous reaction the equilibrium constant
can be expressed as a ratio of the partial
pressures of the products and reactants. For
example the equilibrium constant for the reaction

aA(g) bB(g) ? cC(g) dD(g)
could be expressed as
Note that if one uses this expression the
pressures must be quoted in atmospheres. Also
note that the numerical value of KP might
not be the same as the numerical value of KC.
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Quiz E- Equilibrium Constants

• The reaction between nitrogen gas and oxygen gas
  to form nitrogen dioxide gas
• is shown below.

N2(g) 2O2(g) ? 2NO2(g)
a) Write down an expression for the equilibrium
constant, Kc, for this reaction.
b) Kc is equal to about 2.6 10-15 for this
reaction at 25 oC. In a 1 L flask at 25 oC
there are 1.0 10 13 molecules of N2, 3.0
1015 molecules of O2 and 1.0 1012 molecules
of NO2.
i) Is this system at equilibrium?
ii) If this system is not at equilibrium, in
what direction will the reaction proceed?

• What is the numerical value of the equilibrium
  constant for each of the following
• reactions?

i) ½N2(g) O2(g) ? NO2(g)
ii) 2NO2(g) ? N2(g) 2O2(g)
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Quiz E- Equilibrium Constants

• At a given constant temperature, a 1 L flask
  initially containing 0.100 mol of SO2 and
• 0.150 mol of O2, is allowed to come to
  equilibrium. 80 of the SO2 is found to have
• reacted to form SO3. Calculate the equilibrium
  constant for the reaction-

O2(g) 2SO2(g) ? 2SO3(g)
3. For the endothermic reaction
2SO3(g) ? O2(g) 2SO2(g)
state the effect on the equilibrium constant that
the following disturbances will have.
a) Increasing the concentration of SO3(g).
b) Decreasing the concentration of SO2(g).
c) Doubling the size of the reaction flask.
d) Decreasing the temperature.
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Le Chateliers Principle
Note By changing the position of equilibrium we
are not changing the equilibrium constant if the
temperature remains constant. Equilibrium
constants can only change if one changes the
temperature of the reaction vessel.
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Quiz F- Influencing the Equilibrium

• Consider one of the gaseous equilibria involved
  in the industrial preparation
• of nitric acid by the Ostwald process.

2NO(g) O2(g) ? 2NO2(g) ?H ve
What qualitative effect would the following
disturbances have on the position of the
equilibrium?
a) An increase in PNO.
b) An increase in temperature.
c) A decrease in reactor volume.
d) An increase in pressure via the addition of
an inert gas e.g. Ar.
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Quiz F- Influencing the Equilibrium
2. Consider the following equilibria below
BaCO3(s) (aq) ? Ba2(aq) CO32- (aq)
CO32- (aq) H2O(l) ? CO2(g) 2OH-(aq)
What qualitative effect would the following
disturbances have on the position of the
equilibrium?
a) Making the particle size of the BaCO3 smaller.
b) Decreasing the pH of the aqueous solution.
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Industry
They want to do things as quickly and as
efficiently as possible. However, sometimes,
doing something quickly might not mean doing
the same thing efficiently.
Thermodynamic (i.e. enthalpy ) considerations
might sometimes clash with kinetic (i.e. rate)
considerations.
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Quiz G- Equilibrium rates in industry

• Predict the conditions of temperature and
  pressure required to increase
• productivity in the following industrial
  processes.

a) N2(g) 3H2(g) ? 2NH3(g)
?H -ve
b) 4NH3(g) 5O2(g) ? 4NO (g) 6H2O(g) ?H -ve
c) CaO(s) 3C(s) ? CaC2(l) CO(g) ?H
ve
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Quiz G- Equilibrium Rates in Industry
2. An industrial process to convert X into Y has
the following stoichiometry.
2X(g) ? Y(g) ?H -ve
The reaction is catalysed by a solid heterogenous
catalyst.
The catalyst comes in two different physical
forms
Which of the following set of experimental
conditions would an industrial chemist choose to
optimize the reaction?