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Equilibrium
Section 3
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Mass of solute dissolved
When solute is first added to water
it begins to rapidly dissolve
time
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Mass of solute dissolved
As solute dissolves, some begins
to undissolve…
time
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Mass of solute dissolved
As long as more solute is dissolving
than undissolving, the concentration of
the solution continues to rise.
time
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Mass of solute dissolved
The “undissolving” starts out slow,
and speeds up as the solution
becomes more concentrated
time
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Mass of solute dissolved
At some point the rate of
un-dissolving and rate of
dissolving become the same.
time
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Mass of solute dissolved
At this point the solution has
become saturated.
Maximum = Saturated
time
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Notice that the mass
of dissolved solute
has reached a
maximum
Mass of solute dissolved
The solution’s concentration remains
constant over time since dissolving is
balanced by the undissolving
time
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Saturated solution
CuSO4 (s)  Cu2+
(aq)
+
SO42-
(aq)
Notice the double
arrow in the equation?
This means that the
dissolving is in
equilibrium with the
undissolving.
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Equilibrium
Balance between opposing changes
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Another example:
In a closed soda
bottle the dissolved
CO2 gas is in
equilibrium with the
gas above the liquid.
Soda
Dissolved CO2
CO2 (g)  CO2 (aq)
Though there are many examples of physical
equilibrium, we are also concerned with many
chemical reaction equilibria.
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What changes are in equilibrium here?
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Chemical equilibrium
In Reversible reactions the products of
the reaction will also react to reform the
starting reactants.
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Look carefully: In this reaction H2
molecules react with I2 molecules
to form HI molecules.
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HI molecules can then break
apart and reform H2‘s and I2‘s
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The reaction could be
written: H2 + I2  2HI
OR 2HI  H2 + I2 ?
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Notice that the products once
formed can go back to reactants.
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Graphing equilibrium
H2 + I2  HI
HI increases…to a point
H2 and I2 decrease…to a point
Starting with H2 and I2, once equilibrium is reached…
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Graphing equilibrium
H2 + I2  HI
HI increases…to a point
H2 and I2 decrease…to a point
…concentrations of H2 I2 and HI remain constant over time
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Also, starting with
the product; HI
the same
equilibrium is
reached - H2 I2
and HI remain
constant over time
We could of course write the equation
as the reverse: HI  H2 + I2
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H2 + I2  HI
HI  H2 + I2
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The Haber process
A chemistry classic
N2 + 3H2
NH3
As nitrogen and hydrogen react ammonia is formed
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Ammonia is produced during
the Haber Process
N2 + 3H2
NH3
This reaction starts off fast …but slows as reactants are used up
If you recall,
reaction
speed
depends on
effective
collisions
As
concentration
decreases
the reaction
slows down.
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Ammonia is produced during
the Haber Process
N2 + 3H2
NH3
As ammonia molecules collect …they can decompose to form the reactants
The reverse
reaction
doesn’t start
until some
NH3 is
formed
The
frequency of
collisions
between NH3
is fairly low in
the beginning
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Ammonia is produced during
the Haber Process
N2 + 3H2
NH3
The reverse reaction starts slow …and speeds up as ammonia collects
The
frequency of
collisions
between NH3
increases as
more NH3
molecules
are formed
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Ammonia is produced during
the Haber Process
N2 + 3H2
NH3
Because the reaction is reversible eventually an equilibrium is reached
The forward
reaction and
reverse
reaction are
happening at
the same
speed
The
concentration
of reactants
N2 and H2;
and the
concentration
of product
NH3 will
remain
constant over
time
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Equilibrium:
When rate of forward reaction equals RATE of reverse reaction
concentrations of both reactant and product will remain constant
..when concentration
remains constant
over time
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Equilibrium:
Notice: concentration of reactant
IS NOT equal to concentration of product
The process is affected by temperature and pressure
..when concentration
remains constant
over time
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Reversible reactions can reach equilibrium
Imagine fast
moving
molecules
collide
breaking
bonds and
rearranging
to form
products?
N2 + 3H2
2NH3
Molecules of
product can
also collide,
break apart
and
rearrange to
form the
reactant.
N2 + 3H2  2NH3
Which of the two reactions do you expect to predominate?
The forward reaction (At equilibrium there will usually be a lot more NH3)
For A  B
At what time does the system reach equilibrium?
How do you know?
Once concentrations stop changing (are constant) we know that the
forward and reverse reactions are in equilibrium with each other.
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LeChatier’s
Principle
Section B
Upsetting an Equilibrium
Dissolved CO2
CO2 (g)  CO2 (aq)
Visualize in your mind a bottle of soda which is closed
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Upsetting an Equilibrium
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Dissolved CO2
CO2 (g)  CO2 (aq)
The dissolving and un-dissolving of CO2 are in equilibrium.
Upsetting an Equilibrium
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What happens to the balance when the bottle is opened?
Upsetting an Equilibrium
Most of the CO2 gas escapes
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Upsetting an Equilibrium
CO2 (g)
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CO2 (aq)
Now the un-dissolving is the only change that can occur
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Upsetting an Equilibrium
CO2 (g)
CO2 (aq)
If the bottle is closed again, the CO2 gas can collect.
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Upsetting an Equilibrium
CO2 (g)
And the equilibrium can be re-established
CO2 (aq)
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Upsetting an Equilibrium
CO2 (g)
So a system thrown out of balance,
will always try to return to equilibrium
CO2 (aq)
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Upsetting an Equilibrium
CO2 (g)
This is called LeChatlier’s principle.
CO2 (aq)
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Upsetting an Equilibrium
CO2 (g)
Lechatlier (Le – shat – lee - aye)
CO2 (aq)
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Upsetting an Equilibrium
CO2 (g)
CO2 (aq)
Lechatlier says that a system at equilibrium will
“shift” to counteract the effect of a “stress”
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Upsetting an Equilibrium
CO2 (g)
CO2 (aq)
The stress: allowing CO2 gas to escape stops
the reverse reaction.
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Upsetting an Equilibrium
CO2 (g)
CO2 (aq)
The shift: dissolved CO2 comes out of solution
to replace the lost gas.
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Upsetting an Equilibrium
CO2 (g)
CO2 (aq)
And now the system can return to equilibrium
Upsetting an Equilibrium
Equilbria are stressed by changes in…
Concentration
Temperature
and Pressure
(…if gases are involved of course…)
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Factors Affecting Equilibrium
Changes (stresses) to a system cause equilibrium to
"Shift" left or right
Lets look again at the Haber process:
N2
+
3 H2
 2 NH3
1. Concentration: an increase in concentration
speeds up a reaction
ex: Adding N2 [increasing its concentration]:
the Forward reaction speeds up - (more collisions)
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Concentration
N2
+
3 H2

2 NH3
1. Concentration: increase in concentration speeds
up a reaction
The forward reaction is now faster than the reverse
and the Result:
H2 decreases, while NH3 increases
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Concentration
N2 +
3 H2

2 NH3
1. Concentration: increase in concentration
speeds up a reaction
A trick:
Draw a “pile” over the N2 and
shift right to get rid of the excess.
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N2 +
3 H2

2 NH3
Another example:
Removing NH3
the reverse reaction slows down - (fewer collisions)
(the forward reaction is now the faster one)
Reaction SHIFTS RIGHT (to replace the lost NH3 )
Result: N2 and H2 decrease, NH3 increases
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N2 +
3 H2

2 NH3
Another example:
Removing NH3
A trick:
Draw a hole under the NH3 and
shift right to fill it.
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N2 +
3 H2

2 NH3
LeCHATELIER'S PRINCIPLE: "When a
system at equilibrium is subjected to a stress,
the equilibrium will shift in the direction which
tends to counteract the effect of the stress.“
In other words: When you add something, the
reaction shifts to get rid of it, when you take
something away, it shifts to replace it.]
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2. Temperature
N2
+ 3 H2

2 NH3 + 92 KJ
Increasing temperature speeds up BOTH
forward and reverse reactions.
But… favors reactions which require energy
…endothermic
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2. Temperature
N2
+ 3 H2  2 NH3 + 92 KJ
Increasing temperature - reaction SHIFTS to
the LEFT - to use up extra heat
Result: N2 and H2 increase, NH3 decreases
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N2 + 3 H2  2 NH3
+
92 KJ
Decrease temperature?
reaction SHIFTS to the RIGHT
- to replace lost heat
Reverse reaction slows down more
(since there’s less heat available)
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3. Pressure
N2(g) + 3 H2(g)  2 NH3(g)
3. Pressure: - only affects gases
increasing pressure speeds up both forward and
reverse reactions by increasing concentration
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1 N2(g) + 3 H2(g)  2 NH3(g)
Pressure: depends on moles (coefficients) of reactants and
products
4 moles of reactants vs. 2 moles of products
Increase pressure: reaction SHIFTS to the RIGHT
to reduce # of moles - reduces pressure
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1 N2(g) + 3 H2(g)  2 NH3(g)
Pressure: depends on moles (coefficients) of reactants and
products
4 moles of reactants vs. 2 moles of products
Increase pressure: reaction SHIFTS to the RIGHT
to reduce # of moles - reduces pressure
reducing the number of moles reduces pressure!
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N2(g) + 3 H2(g)  2 NH3(g)
Pressure: depends on moles (coefficients) of reactants and
products
4 moles of reactants vs. 2 moles of products
Increase pressure: reaction SHIFTS to the RIGHT
to reduce # of moles - reduces pressure
Result: N2 and H2 decrease, NH3 increases
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What effect would a Catalyst have?
It Speeds up both reactions equally, favors both
forward and reverse reaction
Doesn’t affect the equilibrium / No shift occurs
For each reaction, identify the shift and the change in
concentrations of reactants and products that will occur:
PCl5(g) + heat  PCl3(g) + Cl2 (g)
Increase [PCl5]
PCl5(g) + heat  PCl3(g) + Cl2 (g)
increase in temp
PCl5(g) + heat  PCl3(g) + Cl2 (g)
Increase press
PCl5(g) + heat  PCl3(g) + Cl2 (g)
1 mole
2 moles
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C(s) + H2O(g) + energy  CO(g) + H2(g)
decrease [CO]
C(s) + H2O(g) + energy  CO(g) + H2(g)
decrease temp
C(s) + H2O(g) + energy  CO(g) + H2(g)
decrease press
C(s) + H2O(g) + energy  CO(g) + H2(g)
grind up the C(s) into powder
C(s) + H2O(g) + energy  CO(g) + H2(g)
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1. Define these in your own words:
Reversible reaction –
Dynamic equilibrium –
2. What is the significance of double arrows in an equation?
3. How is the term equilibrium used to describe an aqueous solution which is
saturated?
4. How do the rates of forward and reverse reactions compare at a state of dynamic
equilibrium?
5. How do the concentrations (or amounts) of reactants and products change once
equilibrium is reached?
6. State LeChatlier’s principle in simple English.
7. How is equilibrium position of this reaction affected by the following changes?
Explain each (shift left or right: hint – draw piles and move them, or holes and fill
them)
C(s) + H2O (g) + energy  CO (g) + H2 (g)
a. increasing temp
b. increasing gas pressure
c. adding H2
d. removing H2
8. What effect does each change have on the concentration of NH3 ?
Explain in terms of shift
N2 (g) + 3H2 (g)  2NH3 (g)
a. adding heat
b. Increasing pressure
c. adding a catalyst
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9. Can a change in pressure affect any reaction at equilibrium?
Explain.
10. Use LeChatlier’s principle to explain why pressure in a bottle of
soda increases again after it is capped. The equation for the dissolving
of Carbon dioxide gas in soda is shown below:
CO2 (g)  CO2 (aq)
11. Given the reaction at equilibrium:
N2 (g) + 3H2 (g)  2NH3 (g) + energy
What effect does raising the temperature have on the rate of the
forward reaction? Explain