Rates ofreactionandequilibrium slides

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Reaction Rates &
Equilibrium
HS Ch i t R id Learning Series
Chemistry Rapid L
i
S i
Wayne Huang, PhD
Kelly Deters, PhD
Russell Dahl, PhD
Elizabeth James, PhD

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1


Learning Objectives
By completing this tutorial you will learn…
What Kinetics studies.
Collision Th
C lli i Theory.
How factors affect rate.
Dynamic equilibrium.
Equilibrium constants.
Reaction Quotients.
Le Ch t li ’ P i i l
L Chatelier’s Principle.

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Concept Map
Previous content
Chemistry

New content

Kinetics
studies

Reaction
Rate
uses

Collision
Theory

Studies
Can be studied with

Forward = reverse

Matter

Equilibrium

Undergo

Equation with
Ratio of
products :
reactants

Equilibrium
Constant
Expression

Chemical
Reactions
When it’s
disturbed,
disturbed
follow

Le Chatelier’s
Principle

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2


Kinetics and Reaction
Coordinate Diagrams

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Definition - Kinetics
Kinetics – Study of the rates of reactions.
Reactants

Products

Reaction Rate – Rate at which reactants
produce products, i.e. how fast a reaction
takes place.
Rate = Δ[Product]/Δt = -Δ[Reactant]/Δt

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3


Definition - Collision Theory
Collision Theory – Defines 3 circumstances
to be met for a collision to occur.

1

Reactants must collide.

2

Collision must be at the correct
orientation.

3

Collision must have minimum
energy for reaction to occur.

Only a small fraction of collisions meet the requirements
and results in a successful reaction.
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Collisions Must Occur
In order for two molecules to react, they must come
into contact with one another.

F
F

O
N
O

There’s no way they’ll ever react if they don’t run into
one another!

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4


Collision with Correct Orientation
For a collision to result in a chemical reaction, it
must occur with the correct orientation.

O
F

F

F

N

F

O
N

O

O

Side-to-Side: This is not the correct orientation.
The reaction will not happen.
Head-to-Head:This is the correct orientation.
The reaction will happen.
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Collision with Enough Energy
For a collision to result in a chemical reaction, it
must occur with the minimum energy for reaction.

F

F

O
F

F

N

O
O
N
N

O

O
O

Slow Speed: The collision does not have enough
energy to produce a reaction.
High Speed: This collision had more energy (faster
moving molecules). A reaction will occur.
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5


Definition - Activation Energy

Activation Energy –
A ti ti
E
Energy that must be
overcome for a
reaction to occur.

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Reaction Coordinate Diagram
Reaction coordinate diagrams show the energy
changes throughout the reaction.
Activated complex
(Also called the transition state)

Energy

Activation E
A ti ti Energy

Products

Reactants
Reaction proceeds

Energy change for
reaction

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6


Activated Complex
What is an “activated complex”?
Energ
gy

Reactants

F2 + NO2

Activated Complex F2…NO2
Products F + FNO2

Reaction proceeds

O
F
F

F
F

F

N

F

O
N

O

O

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Definition - Endo & Exothermic
Endothermic Reaction –
The reaction takes in
energy…the
energy the products
have more energy than
the reactants.

heat

in

Exothermic Reaction –
g
The reaction gives off
energy…the products
have less energy than the
reactants.

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heat

off

Endothermic vs Exothermic Mnemonic: Endo- is to take in heat and Exo- is
to give off heat = “Endo is to Enter (in) and Exo is to Exit (off)!”


7


Endo- and Exothermic Diagrams

Energy

Reaction coordinate diagrams show whether a
reaction is endothermic or exothermic.

Endothermic
Products are higher
energy than reactants.

Energy

Reaction proceeds

Reaction proceeds
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Exothermic
Products are lower
energy than reactants.

Factors Affecting
Reaction Rates

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8


Surface Area of Reactants
How does the surface area of the reactants affect the
reaction rate?

Reactants must
collide in order
to react.

Larger surface
area means more
particles can
come in contact
with each other
at the same time.

More reactants
can collide at the
same time and a
fraction of those
will result in
reaction.

As surface area increases, reaction
rate increases.
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Concentrations of Reactants
How does the concentration of reactants affect the
reaction rate?

Only a small
fraction of the
collisions meet
the requirements
and result in a
reaction.

More reactants
mean more
collisions will
occur.

If more collisions
occur, more will
meet the
requirements
and result in a
reaction.

As the reactant concentration
increases, the reaction rate increases.
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9


Temperature
How does temperature affect the reaction rate?

Reactants must
collide with at
least energy
equal to the
activation
energy.

If molecules are
at a higher
temperature,
they have a
higher average
kinetic energy.

With higher
g
energy
molecules,
collisions will
have higher
energy and more
often result in
reaction.

For most reactions, as temperature
increases, reaction rate increases.

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Definition - Catalysts
Catalysts – Substance that increases the rate
of reaction without being used up
up.
A+B+C

D+C

“C” is the catalyst…it is present in the beginning and in the end.

Energy

Without Catalyst

With Catalyst (lowering activation energy)
Reaction proceeds
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Ex: Enzymes are catalysts in our body.


10


Catalysts
How do catalysts help speed up the reaction without
being used?
They increase the chances that a collision
will successfully produce a reaction.

For example, catalysts hold one or more of
the reactants in place to allow collisions to
occur with the correct orientation.
Once the reaction has occurred, the
catalyst releases the molecule(s) and finds
another one to help.

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Equilibrium

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11


Definition - Equilibrium
Reversible Reaction – A chemical reaction
that can proceed in both directions
(represented by a “ ”)
b
”).

Equilibrium – When the rate of the forward
reaction equals the rate of the reverse
reaction, i.e. Δ[Product]/Δt = -Δ[Reactant]/Δt
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Establishing Equilibrium
It takes time to establish equilibrium.
At first, there are only reactants
present. Only the forward reaction is
p
possible.

Reactants

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Products

But once there are products as well, they can begin to reform
p
,
y
g
reactants.
The reverse reaction becomes possible. Forward rate slows and
reverse rate increases until they are the same.
Once the rate of the forward and reverse process are equal, it is
at equilibrium.
When equilibrium is established, the numbers of products and
reactants doesn’t change…but the reaction keeps going.


12


Definition - Dynamic Equilibrium
Dynamic Equilibrium – The
reaction continues to
proceed in both directions,
but at the same rate.

The
Th number of products and reactants no longer
b
f
d t
d
t t
l
change, it may look as though the reaction has
stopped…
But the reaction continues!
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Equilibrium
Constants

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13


Definition - Equilibrium Constant
Equilibrium Constant Expression –
Equation showing the ratio of the
concentrations of prod cts to
products
reactants at equilibrium.
Concentration is symbolized with
brackets “[A]”.

Equilibrium Constant (K) – The
number calculated from the
equilibrium constant expression.
“K” is different for every reaction at every
temperature!
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Writing Equilibrium Constant Expressions
To write an equilibrium constant expression:

1

Write the concentration of products on the top take each one to a power of the coefficient in the
balanced equation.
b l
d
ti

2

Write the concentration of reactants on the bottom also take each to the power of the balanced
equation coefficient.

Example: Write the equilibrium constant expression for the following:
2 H2 (g) + O2 (g) 2 H2O (g)

K=

[H2O]
[H2] [O2]

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14


Definition: Homo- and Heterogeneous
Equilibrium
Homogeneous Equilibrium – All of
the species are the same state of
matter.
matter
2 H2 (g) + O2 (g)

2 H2O (g)

Heterogeneous Equilibrium –
There are at least 2 states of
matter.
2 H2 (g) + O2 (g)

2 H2O (l)

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Concentrations of Solids and Liquids
Pure solids and pure liquids have constant
“concentrations”.
If concentration (Molarity) = mole
liters
lit
And Density = grams
liters
And Molar Mass = grams
mole
Then f
Th for a pure solid or liquid, Molarity = grams / liters
lid
li id M l it
grams / mole
Or, Molarity =

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Density .
Molar Mass

Both Density and Molar Mass are constants—they don’t change.
Therefore, “concentration” of a pure solid or liquid is a constant.


15


“K” Expressions with Solids or Liquids
How does this affect the writing of Equilibrium Constant
Expressions?
If the “concentration” of a pure solid or liquid is constant,
then
th it will not change d i equilibrium and it is not written
ill t h
during
ilib i
d i
t itt
in the “K” expression.
2 H2 (g) + O2 (g)

2 H2O (g)

K=

[ H 2 O ]2
[ H 2]2 [O2 ]

2 H2 (g) + O2 (g)

2 H2O (l)

K=

1
[ H 2]2 [O2 ]

H2O is not included in this “K” expression because it’s a liquid.
Only gases and solutions are included in “K” expressions!
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Calculating “K” Example
Example: Solve for equilibrium constant for
Fe2O3 (s) + 3 H2 (g) 2 Fe (s) + 3 H2O (g) if the following
are concentrations at equilibrium: [H2] = 0.45 M and
[H2O] = 0.18 M
[H2]eq = 0.45 M

K=

[ H 2O ]3
[ H 2 ]3

K=

[H2O]eq = 0.18 M

[0.18]3
[0.45]3

K=?

Note that Fe2O3 and Fe were
not included in the K
expression as they are solids!

K = 0.064

Most instructors and textbooks do not require units for “K” as each
one would be different.
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16


Meaning of Equilibrium Constant
What general meaning can you get from the magnitude
of the equilibrium constant?

[Products]

If K is very large…

[Reactants]

There is a much larger ratio of products to reactants at
equilibrium.
The reaction is said to “lie to the right” (products are on
the right).
If K is very small
small…

[Products]

[Reactants]
There is a much smaller ratio of products to reactants at
equilibrium.
The reaction is said to “lie to the left”.
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Using “K” to Find Equilibrium Concentration
Example: Find the equilibrium concentration for NO if the
equilibrium constant for N2 (g) + O2 (g) 2 NO (g) is
1.24×10-4, and the other equilibrium concentrations are
[N2] = 0.166 M and [O2] = 0.145 M
[N2]eq = 0.166 M

K=

[O2]eq = 0.145 M
K = 1.24×10-4
[NO]eq = ? M

1.24 ×10 −4 =

[ NO ]2
[ N 2 ][O2 ]

[ NO ]2
(0.166 M )(0.145M )

(1.24 ×10 )(0.166M )(0.145M ) = [ NO]
−4

[NO]eq = 0.00173 M

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17


Reaction Quotient

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What is the Reaction Quotient?
Reaction Quotient is “Q”.
K

Q

Equilibrium Constant

Reaction Quotient

Expression is ratio of
products to reactants with
balanced equation
coefficients as powers.

Expression is ratio of
products to reactants with
balanced equation
coefficients as powers.

Only includes gases and
solutions.

Only includes gases and
solutions.

To solve for K, plug in
concentrations at
equilibrium.

To solve for Q, plug in
concentrations at any time.

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18


The Difference between K and Q
What exactly is the difference?
2 H2 (g) + O2 (g)

2 H2O (g)

K=

[ H 2 O ]2
[ H 2]2 [O2 ]

2 H2 (g) + O2 (g)

2 H2O (g)

Q=

[ H 2 O ]2
[ H 2]2 [O2 ]

The expressions for K and Q are the same.
To solve for “K”, plug in concentrations at equilibrium only.
To solve for “Q”, plug in concentrations at any time.

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Using Reaction Quotient
Reaction Quotient is used to determine if a system
is at equilibrium…and if it’s not, which way does it
need to go to get there.
[products now]

= Q

K =

[reactants now]
Q=K

Q>K

Q<K
38/54

[products at equilibrium]
[reactants at equilibrium]

[now] = [equilibrium]
[Products now] too
large.
[Reactants now] too
small.
[Products now] too
small.
[Reactants now] too
large.


System is at equilibrium.
System will make more
S t
ill
k
reactants to reach
equilibrium.

System will make more
products to reach
equilibrium.

19


Reaction Quotient Example
Example: For N2 (g) + O2 (g) 2 NO (g), if [N2] = 0.81 M, [O2] = 0.75 M
and [NO] = 0.030 M, is the reaction at equilibrium if
K = 0.0025? If not, which way will it go to reach equilibrium?

[N2] = 0.81 M

Q=

[O2] = 0.75 M

[ NO ]2
[ N 2 ][O2 ]

Q=

(0.030M ) 2
(0.81M )(0.75M )

[NO] = 0.030 M
K = 0.0025
At equilibrium = ?

Q = 0.0015

Q<K
Reaction is not at equilibrium.
More products will need to be made (and also thereby
reducing reactants) to have Q = K.
Reaction will go to the right (products) to reach equilibrium.
39/54

Le Chatelier’s
Principle

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20


Definition - Le Chatelier’s Principle
Le Chatelier’s Principle – If a system at
equilibrium is disturbed, it will shift to reestablish the equilibrium
equilibrium.
A system will try to undo
whatever you’ve done
(de-stress or count-change).

41/54

Le Chatelier’s Principle Mnemonic: Le Chatelier is to lesser the change
(stress) added = “Le Chatelier is Lessen (the) Change.”

Increasing Concentrations
How does adding a reactant or product affect a
system at equilibrium? Reactants Products

Adding a
reactant.

Adding a
product.

Q becomes too
small.

Q becomes too
large.

Reaction shifts
R
ti
hift
to right.
(To get rid of
extra reactants
and make more
products.)

Reaction shifts
to left.
(To get rid of
extra products
and make more
reactants.)

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21


Decreasing Concentrations
How does removing a reactant or product affect a
system at equilibrium? Reactants Products

Removing a
reactant.

Q becomes too
large.

Removing a
product.

Q becomes too
small.

Reaction shifts
to left.
(make more
reactants.)

Reaction shifts
to right.
(make more
products.)

43/54

Changes in Pressure
How does changing the pressure affect a system at
equilibrium? i.e. 2 Reactants 3 Products

Decrease
volume.
(i.e. Backward
reaction ).

Increase
volume.
(i.e. Forward
reaction ).

Pressure
increases.
(i.e. Favors
backward
reaction).

Pressure
decreases.
(i.e. Favors
forward
reaction).

Reaction shifts
to the side with
less moles (or
volume, i.e. left)
of gas to
decrease
pressure.

Reactions shifts
to the side with
the more moles
(or volume, i.e.
right ) of gas to
increase
pressure.

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22


Definition - Endo & Exothermic Reactions
Endothermic Reaction – The reaction takes in
heat energy…the products have more energy
than the reactants.
Energy is a reactant in the reaction.
Reactants + heat Products

Exothermic Reaction – The reaction gives off
heat energy…the products have less energy
than the reactants.

45/54

Energy is a product in the reaction.
Reactants Products + heat

Temperature and Endothermic
For endothermic, think of temperature (or energy)
as a “reactant”: Reactants + heat Products
Reaction shifts
R
ti
hift
to right.

Increase
temperature of
endothermic
reaction.

Increasing a
“reactant”.

Decrease
temperature of
endothermic
reaction.

Remove a
“reactant”.

(get rid of extra
reactants and
make more
products.)

Reaction shifts
to left.
(make more
reactants.)

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23


Temperature and Exothermic
For exothermic, think of temperature (or energy) as
a “product”: Reactants Products + heat
Reaction shifts
to left.

Increase
temperature of
exothermic
reaction.

Increasing a
“product”.

Decrease
temperature of
exothermic
reaction.

Remove a
“product”.

(get rid of extra
products and
make more
reactants.)

Reaction shifts
to right.
(make more
products.)

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Changes that Have No Effect
Some changes have no effect because they do not
affect the value of “Q”.
Adding a pure solid or liquid reactant or
g p
q
product.
Increasing pressure by adding an inert
gas.
Changing the volume of a reaction with
an equal number of moles of gas on
each side of the reaction
reaction.
Adding a catalyst.
A catalyst will speed up how fast
equilibrium is established—but not the
number of reactants and products once
it’s at equilibrium.
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24


Le Chatelier’s Examples
Example: Which way will the reaction shift for each of the
following changes:
NH4Cl (s)

NH3 (g) + HCl (g)

Removing NH4Cl

No change (it s a solid)
(it’s

Adding HCl

(Adding a product)

Adding Ne (g)

No change (it’s an inert gas)
(Goes to side with least gas moles)

Decreasing volume

Example: Which way will the reaction shift for each of the
following changes:
2 SO2 (g) + O2 (g)

2 SO3 (g) + Heat … an exothermic reaction

Increasing volume

(Goes to side with most gas moles)

Raising temperature

(Energy is a product)

Adding O2

(Adding a reactant)

Removing SO2

(Removing a reactant)

49/54

Le Chatelier in Industry
When companies need to make large amounts of
product, a reaction with a very small K is a problem.

Small K

Small ratio of
products to
reactants.

Lots of reactants
left over (wasting
money) and few
products made
(not making
money).

They can push the reaction towards the products.
e.g. Remove the products as they’re made, adjust pressure or
temperature as needed to push it to the right.
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25


Haber Process
The reaction to produce ammonia, NH3, is very
important to manufacturing.
N2 + 3 H2
2 NH3
(an exothermic reaction: ΔH = -92 kJ/mol)
92
Rate = Δ[NH3]/Δt = -(1/2)Δ[N2]/Δt = -(3/2)Δ[H2]/Δt
In order for the reaction to occur at a reasonable rate, the
temperature must be very high.
But when the temperature is high, the equilibrium
constant is very low.
A compromise is made and a catalyst is added to
increase the rate at the lower temperature.

The reaction yields 20%...the leftover reactants are
recycled and put back into the reaction again.
51/54

Learning Summary

Le Chatelier’s Principle
governs h
how a reaction at
ti
t
equilibrium will change
when disturbed.

Kinetics is the study of
the rates of reaction,
which are affected by
surface area,
concentration,
temperature and
catalysts.

Dynamic equilibrium is
established when the rates
of the forward and reverse
reactions are equal.

The
Th equilibrium constant
ilib i
t t
give the ratio of product :
reactants with the
stoichiometric ratios as
the powers.

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26


Congratulations
You have successfully completed
the core tutorial

Reaction Rates &
Equilibrium

Chemistry :: Biology :: Physics :: Math

What’s N t
Wh t’ Next …

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