Chemical Kinetics Chapter 13. Chemical Kinetics Thermodynamics – does a reaction take place?...
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Transcript of Chemical Kinetics Chapter 13. Chemical Kinetics Thermodynamics – does a reaction take place?...
Chemical KineticsChapter 13
Chemical Kinetics
Thermodynamics – does a reaction take place?
Kinetics – how fast does a reaction proceed?
Reaction rate is the change in the concentration of a reactant or a product with time (mol/(Ls)).
A B
rate = -[A]t
rate = [B]t
[A] = change in concentration of A over time period t
[B] = change in concentration of B over time period t
Because [A] decreases with time, [A] is negative.
13.1
A B
13.1
rate = -[A]t
rate = [B]t
time
Br2 (aq) + HCOOH (aq) 2Br- (aq) + 2H+ (aq) + CO2 (g)
time
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Detector
[Br2] Absorption13.1
Br2 (aq) + HCOOH (aq) 2Br- (aq) + 2H+ (aq) + CO2 (g)
average rate = -[Br2]t
= -[Br2]final – [Br2]initial
tfinal - tinitial
slope oftangent
slope oftangent slope of
tangent
instantaneous rate = rate for specific instance in time13.1
Reaction Rates and Stoichiometry
13.1
2A B
Two moles of A disappear for each mole of B that is formed.
rate = [B]t
rate = -[A]t
12
aA + bB cC + dD
rate = -[A]t
1a
= -[B]t
1b
=[C]t
1c
=[D]t
1d
Write the rate for the following reaction:
CH4 (g) + 2O2 (g) CO2 (g) + 2H2O (g)
rate = -[CH4]
t= -
[O2]t
12
=[H2O]
t12
=[CO2]
t
13.1
Exothermic Reaction Endothermic Reaction
The activation energy (Ea ) is the minimum amount of energy required to initiate a chemical reaction.
13.4
A + B AB C + D++
Kinetic Molecular Theory
• Model of what happens to gas particles during experimentation– Large numbers of molecules in continuous
motion– Attractive and repulsive forces are negligible – Energy is transferred between molecules during
collisions, but average kinetic energy is unchanged (as long as temp is constant)
– Kinetic energy of molecules is proportional to the absolute temperature. (in K)
COLLISION THEORYCOLLISION 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
COLLISION THEORYCOLLISION 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 orhave more particles present
more successful collisions give particles more energy orlower the activation energy
CO(g) + NO2(g) CO2(g) + NO(g)
A Reaction Profile
Particles must collide with the proper geometry or orientation for atoms to come in direct contact and form the chemical bonds of the products. (steric factor)
• If both of these conditions are not met, particles will merely collide and bounce off one another without forming products.
http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/collis11.swf
• Although, the percentage of successful collisions is extremely small, chemical reactions still take place at a reasonable rate because there are so many collisions per second between reactant particles.
INCREASING THE RATEINCREASING THE RATE
• INCREASE THE SURFACE AREA OF SOLIDS
• INCREASE TEMPERATURE
• ADD A CATALYST
• INCREASE THE PRESSURE OF ANY GASES
• INCREASE THE CONCENTRATION OF REACTANTS
• INCREASE THE SURFACE AREA OF SOLIDS
• INCREASE TEMPERATURE
• ADD A CATALYST
• INCREASE THE PRESSURE OF ANY GASES
• INCREASE THE CONCENTRATION OF REACTANTS
The following methods may be used to increase the rate of a chemical reaction
INCREASING SURFACE AREAINCREASING 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.
CUT THE SHAPE INTO SMALLER
PIECES3
3
1
SURFACE AREA9+9+3+3+3+3 = 30 sq units
SURFACE AREA9 x (1+1+1+1+1+1) = 54 sq units
11
1
Effect increasing the temperature increases the rate of a reactionparticles get more energy so they can overcome the energy barrierparticle speeds also increase so collisions are more frequent
INCREASING TEMPERATUREINCREASING TEMPERATURE
ENERGY CHANGESDURING 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
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.
INCREASING TEMPERATUREINCREASING TEMPERATURE
MOLECULAR ENERGY
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MAXWELL-BOLTZMANN DISTRIBUTION OF
MOLECULAR ENERGY
MAXWELL-BOLTZMANN DISTRIBUTION OF
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
T1
T2
TEMPERATURE
T2 > T1
MAXWELL-BOLTZMANN DISTRIBUTION OF
MOLECULAR ENERGY
MAXWELL-BOLTZMANN DISTRIBUTION OF
MOLECULAR ENERGY
INCREASING TEMPERATUREINCREASING TEMPERATURE
MOLECULAR ENERGY
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Decreasing the temperature alters the distribution
• get a shift to lower energies/velocities• curve gets narrower and more pointed due to the smaller spread of values• area under the curve stays constant - it corresponds to the total number of particles
T1
T3
TEMPERATURE
T1 > T3
MAXWELL-BOLTZMANN DISTRIBUTION OF
MOLECULAR ENERGY
MAXWELL-BOLTZMANN DISTRIBUTION OF
MOLECULAR ENERGY
INCREASING TEMPERATUREINCREASING TEMPERATURE
MOLECULAR ENERGY
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REVIEWno particles have zero energy/velocitysome particles have very low and some have very high energies/velocitiesmost have intermediate velocitiesas the temperature increases the curves flatten, broaden and shift to higher energies
T1
T2
T3
TEMPERATURE
T2 > T1 > T3
MAXWELL-BOLTZMANN DISTRIBUTION OF
MOLECULAR ENERGY
MAXWELL-BOLTZMANN DISTRIBUTION OF
MOLECULAR ENERGY
INCREASING TEMPERATUREINCREASING TEMPERATURE
MOLECULAR ENERGY
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Ea
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.
MAXWELL-BOLTZMANN DISTRIBUTION OF
MOLECULAR ENERGY
MAXWELL-BOLTZMANN DISTRIBUTION OF
MOLECULAR ENERGY
NUMBER OF MOLECULES WITH SUFFICIENT ENERGY TO OVERCOME THE ENERGY BARRIER
INCREASING TEMPERATUREINCREASING TEMPERATURE
MOLECULAR ENERGY
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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
T1
T2
TEMPERATURE
T2 > T1
Ea
MAXWELL-BOLTZMANN DISTRIBUTION OF
MOLECULAR ENERGY
MAXWELL-BOLTZMANN DISTRIBUTION OF
MOLECULAR ENERGY
INCREASING TEMPERATUREINCREASING TEMPERATURE
MOLECULAR ENERGY
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EXTRA MOLECULES WITH SUFFICIENT ENERGY TO OVERCOME THE ENERGY BARRIER
• 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.
ADDING A CATALYSTADDING A CATALYST
WITHOUT A CATALYST WITH A CATALYST
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.
MOLECULAR ENERGY Ea
MAXWELL-BOLTZMANN DISTRIBUTION OF
MOLECULAR ENERGY
MAXWELL-BOLTZMANN DISTRIBUTION OF
MOLECULAR ENERGY
NUMBER OF MOLECULES WITH SUFFICIENT ENERGY TO OVERCOME THE ENERGY BARRIER
ADDING A CATALYSTADDING A CATALYST
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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
Ea
MAXWELL-BOLTZMANN DISTRIBUTION OF
MOLECULAR ENERGY
MAXWELL-BOLTZMANN DISTRIBUTION OF
MOLECULAR ENERGY
ADDING A CATALYSTADDING A CATALYST
EXTRA MOLECULES WITH SUFFICIENT ENERGY TO OVERCOME THE ENERGY BARRIER
MOLECULAR ENERGY
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• 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 reactantsdifferent phase to reactants e.g. CFC’s and ozone e.g. Fe in Haber process
Uses used in industry especially where an increase in temperature results ina lower yield due to a shift in equilibrium (Haber and Contact Processes)
CATALYSTS - CATALYSTS - A REVIEWA REVIEW
• increasing the pressure forces gas particles closer together
• this increases the frequency of collisions so the reaction rate increases
• many industrial processes occur at high pressure to increase the rate... but it can adversely affect the position of equilibrium and yield
The more particles there are in a given volume, the greater the pressure The greater the pressure, the more frequent the collisions The more frequent the collisions, the greater the chance of a reaction
INCREASING THE PRESSUREINCREASING THE PRESSURE
Increasing concentration = more frequent collisions = increased rate of reaction
INCREASING CONCENTRATIONINCREASING CONCENTRATION
However, increasing the concentration of some reactantscan have a greater effect than increasing others
Low concentration = fewer collisions Higher concentration = more collisions
Concentration
• A higher concentration of reactants leads to more effective collisions per unit time, which leads to an increasing reaction rate
• We are not increasing the amount being made for a given balanced equation with limiting reactants, we are only speeding up how quickly those products are made.
CuCO3(s) + 2HCl(aq) CuCl2(aq) + CO2(g) + H2O(l)
blueDescribe seven different ways to monitor the rate of the above reaction. State how each property would change as the reaction proceeds.
1. Mass of CuCO3(s) over time decreases
2. [HCl] over time decreases
3. [CuCl2] over time increases
4. Volume of CO2 over time increases
5. Mass of a open beaker over time decreases
6. Pressure of a closed beaker over time increases
7. Colour of the blue Cu2+ over time increases
CuCO3(s) + 2HCl(aq) CuCl2(aq) + CO2(g) + H2O(l)
blue
Describe five different ways to increase the rate of the above reaction.
1. Increase the temperature
2. Increase [HCl]
3. Add a catalyst
4. Increase the surface area of CuCO3(s)
5. Agitate
We can't change the nature of the reactant because then we wouldn't have the same reaction. Replacing HCl with H2SO4 would be faster but a different reaction.
