Transcript advanced chem- kinetics and equilibrium

Ch. 12-13
Reaction Kinetics and Equilibrium
Reaction Kinetics
• Looks at the reaction process and
the factors that help us predict
reactions
Stability
• Thermodynamically Stable:
– Reaction does not spontaneously
occur
• Kinetically Stable:
– Spontaneous
– Reaction is occurring so slow it is
undetected ( but things are still
reacting)
– Ex. Decomposition of H2O2 (needs
brown bottle)
Reaction Mechanism
• Rxn occurs in a series of steps
• Reaction Mechanism:
– Series of reaction steps that must
occur for a reaction to go to completion
– Each step has 2 particles colliding
Ex. A + B -> C (step 1)
C + D -> E (step 2)
E + F -> G (step 3)
Total Rxn : A+ B + D + F  G
• There were intermediates in
between that you never saw (C, E)
• you only see the original reactants
and the final products
• Intermediates:
– Something that appears in the series
but not in the final product
Ex. N2O  N2 + O (step 1)
N2O + O  N2 + O2 (step 2)
Total rxn. : 2N2O  2N2 + O2
(O was an intermediate)
• Clock Reactions
• Reaction Mechanism- teaching it
Homogeneous Reaction
• All reactant(s) and product(s) are in
the same phase
Heterogeneous Reaction
• Reaction that takes place at the
interface between 2 phases
• Zn (s) + HCl (aq)  H2 (g) + ZnCl2 (aq)
( HCl bubbles on the surface of the Zn)
Collision Theory
In order for a rxn. to occur their
particles must collide & those
collisions must result in interactions
• Collisions must:
– Collide w/ enough energy
– Have particles positioned in a way that
enables them to react
– Rate of reaction song
Factors that affect
reaction rates:
1. Nature of reaction:
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•
Dependent on the type of bond
involved
Ionic reaction rates, faster than
covalent
2. Stirring:
•
Molecules in faster motion increase
probability that the particles will hit &
collide w/ enough energy
3. Crushing:
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Smaller pieces increase the surface
area so there are more possible sides
for collisions
Lycopodium Small scale- creamer Mythbusters-creamer
4. Concentration: (video)
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Quantity of matter that exists in a unit
of volume
Increasing concentration increase #
of collisions therefore increasing rate
Ex. Double the concentration  4x
the collisions
5. Pressure (works for gases)
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Increase pressure, decrease volume
So you have the same # of molecules
in a smaller space, more molecules
per unit of volume (i.e. higher
concentration)  more collisions that
could occur  increase rate
6. Temperature:
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Measure of average kinetic energy
(frequency of collisions)
Increase that frequency , the
collisions increase
Increase temperature does 2 things:
A. Heating up molecules, moves them
faster, more chances for collisions
B. More kinetic energy in molecules
increase the motion of particles, easier to
get over that activation energy, rate of
reaction will increase
Commercial Break
What is this?
A “Cattle List”
(get it, a catalyst)
Ha, ha, ha, ha, ha, ha,
ha,ha, ha, ha, ha (I crack
myself up)
7. Catalyst
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A chemical that increases the speed of
the reaction but remains chemically
unchanged
Doesn’t change the normal position of the
equilibrium
Same amounts of product will be formed
w/ or w/out the catalyst –just takes longer
Types: homogeneous & heterogeneous
•
Sugar/sulfuric
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•
Heterogeneous Catalyst
• Surface catalyst
• Ex. metal oxides, platinum
• Works by adsorption – the
adherence of one substance to the
surface of another
• Catalyst has specific lumps that hold
the chemicals in the right position to
react (increase the chance of them
coming together)
• Catalytic converter:
– Platinum honeycomb structure (more
surface area)
– Pollution  SO2, CO2, NO
• Converter lets H2O react w/ gases to
convert them to weak acids (more
complete combustion)
Homogeneous Catalyst
• In same phase as reactants
• Forms an intermediate or activated
complex
• Reactant reacts better w/ the
catalyst than the other reactant
• Ex. Sulfuric acid in ester reaction
enzymes
Colbalt chloride
8. Entropy
•
Chemical systems tend to achieve
the lowest possible energy state
(more stable)
Law of Disorder – states that things
move spontaneously in the direction
of maximum chaos
Entropy
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Can be thought of as– measure of the
disorder of the system or the
randomness (more stable)
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Entropy
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More exact definition- measure of the
number of possible ways that the energy
of a system can be distributed; related to
the freedom of the system’s particles to
move and the number of ways they can
be arranged (energy dispersal)
• Misconceptions about Entropy
• This view of the second law of thermodynamics is very
popular, and it has been misused. Some argue that the
second law of thermodynamics means that a system can
never become more orderly. Not true. It just means that in
order to become more orderly (for entropy to decrease),
you must transfer energy from somewhere outside the
system, such as when a pregnant woman draws energy
from food to cause the fertilized egg to become a complete
baby, completely in line with the second line's provisions.
• Entropy of gas is greater than liquid
or solid
• Entropy increase when a substance
is divided into parts
• Entropy increase w/ increase in
temperature
9. Inhibitors
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Prevents reaction from happening for a
certain length of time (delays reaction)
Not opposite of catalyst
Ex. Lemon juice on apples
A + B  AB
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w/ inhibitor: A-inh + B  no rxn.
Once inhibitor is used up then: A + B  AB
Energy Diagrams
• Activation Energy:
– Energy required to start a chemical
reaction
– High activation energy  few collisions
have enough energy for a reaction 
get slow undetected reaction
• Activated Complex:
– Product formed when reactants have
collided w/ sufficient energy to meet
activation energy requirement
Energy Diagram:
Exothermic Rxn
-releases energy, lower energy after
rxn.
Energy diagram:
Endothermic Rxn.
-absorbed energy, higher energy after
rxn.
• Endo thermic/exothemic song
Energy Diagram:
Catalyst
-w/ catalyst product formed faster
-lowers the activation energy requirement
Reaction Rate:
• Rate of disappearance of one of the
reactants or rate of appearance of
one of the products
• Unit: (mole/L)/s
– Change in molarity per second
• Reaction rate song (second time)
Rate Law:
• Rate is dependent on the
concentration of the reactants
• Expression relating the rates of
reaction to the concentration of
reactants
• [ ] = concentration
A + B  AB
• Rate = k [A] [B]
• k= specific rate constant
(proportionality constant relating to
concentration – value changes
depending on rxn.)
Ex. H2 + I2 => 2HI
rate= k [H2 ] [I2 ]
Exp. 1- [H2 ] = 1.0M
[I2 ] = 1.0M
rate= .20 M/s
k=?
.20= k [1.0M] [1.0M]
k = .20
Exp. 2 - [H2] = .5 M [I2] = .5 M
rate = ?
k= .20
rate= k [H2 ] [I2 ]
rate = .20 [.5] [.5]
rate= .05 M/s
• The rate law for elementary
reactions is just the product of the
reactants, reactions that have more
than one step you would need to
figure out the order of reaction.
Rate Determining Step
• The step or reaction in the series
that is slower than all the others 
the reaction rate is dependent on
this
Ex. Person going 45 in the left lane on
I-94
Reaction Order or Order
of Reaction
• Changing the concentration of
substances taking part in a reaction
usually changes the rate of reaction
• A rate equation shows this effect
mathematically
• Orders of reaction are a part of this
rate equation (helps us describe the
reaction )
• Orders of Reaction are always found
by doing experiments
Elementary Reactions
• A reaction with no intermediate steps (very
rare) – not a reliable way to determine
order
• One can determine the order with the
coefficients
• Rate is proportional to the concentration of
the reactants raised to the power of the
coefficients
Rate expressed as:
aA + bB  cC + dD
Rate = k [A]a [B]b
( a and b are the coefficients)
Reaction Order
• Can determine reaction order
experimentally or graphically
• Experimentally:
• Gather data and see what happen to
rates if you change the concentration (1st
order- double [ ] doubles rate, 2nd order –
double[ ] quadruples rate, zero order- rate
constant with any [ ] )
• Graphically:
• Plot concentration vs. time – identify
which graph gives you a linear graph
– Zero Order: Linear Graph [A] vs time
– 1st order: Linear graph ln[A] vs time
– 2nd order: Linear graph 1/[A] vs. time
• Sum of the power to which all the
reactant concentrations are raised
(always defined in terms of reactant
concentrations (no products))
• Overall order = a + b
(exponents added together)
Finding overall Order
Ex. Rate = k [A] [B]2
Rate is 1st order for reactant A
Rate is 2nd order for reactant B
Overall order =(a + b) = 3rd order
-if you double [A] = doubles rate
-if you double [B] = quad. rate
Practice Problems
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Rate Law:
Reaction:
– 2NO(g) + Cl2(g) 2 NOCl (g)
• Using the following data, calculate the
rate law and constant.
[NO]
[Cl2]
0.10
0.10
Rate
(Δ[ ]/Δt)
0.18
0.10
0.20
0.36
0.20
0.20
1.45
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What is the rate law?
Rate = k [NO]2 [Cl2]
What is the order of the reaction
with respect to NO?
• 2nd order
• What is the order of the reaction with
respect to Cl2
• 1st order
• Using the data and rate law,
calculate the rate constant.
• k = 180
• Assign: p 567 #21, #27 a,b #30 a,
39, 65 a,b
Equilibrium
-use for reversible reactions
Equilibrium
• Means a state of balance
• We will look at dynamic equilibrium –
where changes are taking place but
the overall balance is maintained
(happening at the molecular level)
• Not Static equilibrium- where nothing
is moving any more (see-saw)
• Use equilibrium w/ reversible
reactions, where reactants convert
to products and products convert to
reactants simultaneously
• Ex. Equilibrium reaction
2SO2(g) + O2(g) ↔ 2SO3 (g)
Can also use these symbols
↔,
2SO2(g) + O2(g) ↔ 2SO3 (g)
Steps:
1. In a reversible rxn., the rate of the
reverse process is zero at the
beginning. At that point no
products are going back to
reactants.
2. As the concentration of products
build up, some products start
converting to reactants (reverse
rxn. starts)
2SO2(g) + O2(g) ↔ 2SO3 (g)
3. As reactants are used up their
concentration decreases (forward
reaction slows down)
4. As products build up, their
concentration increases (reverse
reaction speeds up)
5. Eventually the products are going
to reactants at the same rate as
the reactants are going to
products; the rxn. has reached
equilibrium
Chemical Equilibrium
• Forward & reverse reactions are
taking place at same rate (no net
change in actual amts. of products
or reactants in the system)
2SO2(g) + O2(g) ↔ 2SO3 (g)
C
O
N
C
E
N
T
R
A
T
I
O
N
100
75
50
25
0
SO3
SO2
O2
TIME 
(Have twice as much SO2 as O2 initially, then a mixture of 3 gases is obtained
at equilibrium)
Equilibrium Position
• Given by relative concentration of
reactants & products at equilibrium
• Doesn’t mean exactly 50%/50%
concentration at equilibrium
• The position indicates what is
favored at equilibrium
A
B= product bond is weak and
you have mostly reactants at equil.
A
B = product bond is strong and
you have mostly products at equil.
(larger arrow indicates the favored
direction)
• Catalyst speed up the forward and
reverse reactions equally (the
activation energy is reduced by the
same amount)
• Catalysts don’t affect the amts. of
reactants or products present at
equil. ( just decrease the time to get
to equil.)
Equilibrium Constant (Keq)
• Use of constant is a concise way of
stating whether reactants or
products are favored in a rxn.
• The constant #’s relate the amt. of
reactants to products at equil.
• Keq show the ratio of products to
reactants
• If Keq > 1 products favored at equil.
(spontaneous rxn)
• If Keq < 1 reactants favored at equil.
(non-spontaneous rxn)
Ex. aA + bB  cC + dD
(a) moles of reactant A react w/ (b)
moles of reactant B and give (c)
moles of product C and (d) moles of
product D
Keq (equil. constant)=
ratio of product concentrations to
reactant concentrations w/ each
concentration raised to a power
given by the # of moles of that
substance in the balanced chem.
rxn.
-Keq is dependent on temp., as temp
changes Keq changes
aA + bB  cC + dD
Keq = [C]c x [D]d
[A]a x [B]b
H2 + I2  2HI
Write the Keq equation for this.
Keq = [HI]2
[H2] [I2]
N2O4 (g)  2 NO2 (g)
•
This is a homogeneous equil.- all
substances are in same phase
1. Write Keq equation:
Keq= [NO2]2
[N2O4]
N2O4 (g)  2 NO2 (g)
2. Calculate the Keq if:
[NO2] = .0045 mol/L
[N2O4] = .030 mol/L
Keq= [NO2]2
[N2O4]
Keq = [.0045]2
.030
Keq = 6.8 x 10 -4
Keq = 6.8 x 10
-4
3. What is favored at equilibrium?
Reactants (Keq<1)
N2 (g) + O2(g)  2NO(g)
1. Write the Keq
Keq = [NO]2
[N2] [O2]
N2 (g) + O2(g)  2NO(g)
2. If [N2] and [O2] = .72 M and
Keq = 4.6 x 10 –31. What is [NO]?
4.6 x 10-31 = [NO]2
[.72] [ .72]
2.38 x 10 –31 = [NO]2
4.9 x 10 –16 M = [NO]
Book problems
• P. 614 (new book)
• # 17, 21,22,23,37
• P 587 #24-26 (old book)
Le Chatelier’s Principle
demo
• Delicate balance exists between
reactants and products in a system
at equilibrium
• If equil. conditions are changed –
system shifts to restore equilibrium
• Any application of stress to the
system disrupts the system
• Henri LeChatelier studied the
changes in systems w/ stress
application
LeChatelier’s Principle =
-if a stress is applied to a system in
a dynamic equilibrium, the system
changes to relieve the stress
-stress types: concentration,
temperature, pressure
1. Changes in Concentration
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•
Change amounts of reactants or
products
System changes to minimize the
original change
Ex. CO2 + H20
H2CO3
• Adding a reactant always pushes a
reversible reaction in the direction of
the products
– Shifts reaction to right ()
– Forms more product, uses up excess
reactants
Ex. CO2 + H20
H2CO3
• Removing a reactant always pulls a
reversible rxn in the direction of the
reactants
– Shifts reaction to the left ()
– Forms more reactants (less products)
Ex. CO2 + H20
H2CO3
• If product is added at equil. the
reaction shifts to the formation of
more reactants
– Shifts to the left ()
– Forms less product, more reactant
• If product is removed at equil.
reaction shifts to direction of
products
– Shifts to the right ()
– Forms more product, less reactant
– Removal is a trick used by chemists to
increase the yield of a desired product
(ex. Hens & eggs)
2. Change in temperature
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Increasing the temperature causes
the equilibrium position of a reaction
to shift in a direction that absorbs the
heat energy
Keq changes if temp. changes
Reversible rxns are endothermic in
one direction, exothermic in the other
Effect of a change depends on which
is endo
In a reaction
-addition of heat  favors endo side
-removal of heat  favors exo side
2SO2 + O2<-> 2SO3 + heat (kcal)
• Think of heat as a product
•  exo direction(points towards heat)
•  endo direction (points away from
heat)
• If I add heat– shifts away from heat
to cool system (shifts left-toward
reactants)
• If I cool – shifts towards heat (shifts
right-toward products)
3. Change In Pressure
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Affects gas phase only
Affects the number of moles
Similar effect as increasing
concentration of any gas
Le Chat. States that if the pressure
on an equilibrium system is changed
the rxn. is driven in a direction that
relieves that stress
• An increase in pressure (decreases
volume) favors the side w/ the least
moles
• A decrease in pressure (increases
volume) favors the side w/ more
moles
• If moles are equal on both sides- no
pressure effect
Pressure Increase
PCl5 (g) <-> PCl3 (g) + Cl2(g)
1 mole
2 moles
- favors side w/ least moles
- Shifts toward left (reactants)
-  PCl5
 PCl3
 Cl2
Pressure decrease
2PbS(s) +3O2(g)<->2PbO(s) +2SO2(g)
-look only at gases
3 moles
2 moles
-favors side w/ most moles
-shift toward left (reactants)
O2
SO2
test
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18 multiple choice
4 short answer
2 calculations
3 essay