Transcript Chemical Equilibrium

Chemical Equilibrium
Chemistry .2
Midland High School
Mrs. Daniels
April 2007
Chemical Equilibrium
 Equilibrium is a BALANCE between two
opposing forces or processes
 What does it take for two people to be in
equilibrium on a teater totter?
 When playin tug o’ war?
 So, when does chemical equilibrium
happen?
Chemical Equilibrium
 When a reaction occurs until the
concentrations of products and reactants
no longer changes, the reaction is said to
have reached equilibrium.
 Does this mean that the reaction has
stopped completely?
 NO…it simply means that the RATE at
which product is being produced and the
reactants are being reformed is equal.
 Does this mean that the NUMBER of
products and reactants is equal? NO!
A Practical Example
 Imagine 5 students inside the classroom
and 25 students in the hallway.
 If 2 students leave the classroom through
one door which 2 students from the
hallway enter the classroom, does the
number of individuals in or out of the
room change?
 NO…The rate of entering is the same as
the rate of leaving.
 This is equilibrium…WITHOUT having
equal numbers of students in and out
Equal movement in both directions
means the rate is equal; therefore,
equilibrium is met
A chemical example
 Let’s look at a situation where chemical
equilibrium is met…
 When preparing a saturated solution of
sodium chloride, equilibrium is met and can
be written as follows:
 NaCl (s) <--> Na+ (aq) + Cl- (aq)
 Notice that the arrow is bidirectional
 This MUST be a closed system…no more
reactants or products are added from an
outside source
Relative Concentration
 There is NO WAY you can tell what the
concentrations of any of the products or
reactants are by just looking at the
equation; however, there is a ratio that
scientists were able to come up with after
much study:
 The law of chemical equilibrium
 Keq = [products] / [reactants]
 Remember that the [brackets] mean
“concentration of” whatever is inside
Keq
 The calculation of Keq isn’t quite that
simple, but conceptually it is.
 For the reversible equation
 aA + bB --> xX + yY
 With the lower case letters being the
number of moles
 Keq is calculated:
 [X]x[Y]y / [A]a[B]b
Keq
 Try writing out the formula for Keq for the
following reaction:
 4 NO(g) + 6 H2O(g) <--> 4 NH3 (g) + 5 O2 (g)
 [NH3]4[O2]5 / [NO]4[H2O]6 = Keq
 Practice:
 Try a few problems from your teacher
 Set up the equilibrium constants (Keq’s) for
these equations
Practice Problems:
 1. 2 NO2 (g) <--> N2O4 (g)
 2. N2 (g) + 3 H2 (g) <--> 2 NH3 (g)
 3. 2 NO
(g)
+2 H2 (g) <--> N2 (g) + 2 H2O
(g)
In the Lab
 Calculating the Keq in the lab involves
writing the correct formula for Keq,
substituting in the measured
concentrations
 H2 + I2 --> 2 HI
 [H2]= .0056 M
 [I2]= .00059 M
 [HI]= .0127 M
 Keq = [HI]2 / [H2][I2]
 Keq = (.0127)2 / (.0056)(.00059)
 Keq = 48.8
 If you measured different concentrations in
subsequent lab trials, calculate Keq and
then average all of the Keq values.
The Winds of CHANGE
 We’ve discussed that equilibrium can
exist only under conditions of constant
temperature, pressure, volume, and
concentration.
 Henri LeChatelier examined what occurs
when these factors do not remain
constant
 If one of these factors change, it is said
to put “stress” on the reaction
LeChatelier’s Principle
 When a stress is placed on a system in
equilibrium, the system will adjust to
remove the stress and to restore
equilibrium in the system
 1. Changes in Concentration
 This principle allows us to predict the
direction in which the equilibrium will
shift when one or more of the
concentrations of the products or
reactants is altered.
Predicting Direction
 You can use LeChatelier’s Principle to
predict the direction of the “teeter” or
“totter”
 Which direction will the reaction move
(toward the reactants or toward the
products) in order to reestablish
equilibrium?
 Let’s Practice!
Changes in Concentration
 CO(g) + 2H2 (g) <--> CH3OH (g)





What direction does the rxn shift if…
More CO is added?
Methanol is increased?
Methanol is removed?
Hydrogen gas source is reduced?
LeChatelier’s and Pressure
 2. The pressure of a system is directly
proportional to the number of gas
molecules present…
 So the only way to reduce the pressure is
to reduce the total number of molecules
in the system
 Increasing pressure on a gaseous system
causes the equilibrium to shift to the side
with the fewest number of molecules
 So, if the opposite is true and pressure is
decreased, then the eq shifts to the side
with the greatest number of molecules
Try this
 The following rxn has come to equilibrium
in a container:
 N2 (g) + 3 H2 (g) <--> 2 NH3 (g)
 In which direction will the rxn shift if the
pressure on the system above decreases?
 Left
 Why?
 Should the pressure on the above system
be increased or decreased to produce more
ammonia? Why?
Volume and LeChatelier’s
 3. When the volume of a rxn is reduced,
the molecules are crowded together.
 Decreasing the number of molecules can
decrease the stress.
 When the volume of a rxn involving
gases decreases, the eq shifts to the side
with the fewest number of molecules.
(when the gas rxn volume increases, the
eq shifts to the side with the greatest
number of molecules.)
Try these:
 Will the reaction shift toward the reactants
(left) or the products (right) side if the
VOLUME IS DECREASED:
 PCl5 (g) <--> PCl3 (g) + Cl2 (g)
 Left
 N2 (g) + 3 H2 (g) <--> 2NH3 (g)
 Right
 2CO (g) + O2 (g) <--> 2CO2 (g)
 Right
Temperature and LeChatelier’s
 4. Keq is temperature, so heating or
cooling the reaction will result in shifting
the eq to the left or right depending on
whether the rxn is endothermic or
exothermic.
 Exothermic:
 Increasing the temperature of an
exothermic rxn
 Reactants <--> products + heat energy
is like increasing a product, so the shift
will be to the left
Decreasing will have the opposite effect
 In an endothermic reaction however,
increasing the temperature would be like
increasing a reactant and would force the
shift to the right
 Reactants + heat energy <--> products
Ksp
 When an ionic solid is placed in water, an
equilibrium is established between the
ions in the saturated solution and the
excess solid phase.
 Ex.
 AgCl(s) <--> Ag+(aq) + Cl-(aq)
 or Ag3PO4 (s) <--> 3Ag+(aq) + PO4-3(aq)
 For each of these, we could write the Keq
expressions. This is called the Ksp or
solubility product
 Do NOT include the solid (in Ksp or in Keq)
 Write the Ksp for the above equations:
 [Cl-][Ag+] = Ksp
 And for the second reaction,
[Ag+]3[PO4-3] = Ksp
What does Ksp tell us?
The larger the Ksp, the more soluble a salt
is in water.
When doesn’t equilibrium occur?
 All of the reactions we’ve discussed have
been in EQUILIBRIUM, so do all reactions
reach equilibrium?
 What types of rxns don’t?
 Strong acid ionization, precipitations,
formation of a gas from an aqueous
solution, and the formation of water as a
product of the reaction
 Why don’t these types of rxns reach eq?