# Unit #3 Electron Configuration & Periodicity

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```					                 Unit #8 Chemical Equilibrium
Upon completion of this unit, the student should be able to:

1. assemble an equilibrium constant expression, K, given a balanced equation
involving gases.
2. estimate the value of the equilibrium partial pressures of the reactants and products
in a system given the original partial pressures of all products and reactants and the
partial pressure of one product or reactant.
3. determine the value of the equilibrium constant, K, given the original partial
pressures of all products and reactants and the partial pressure of one product or
reactant.
4. predict the effect of adding reactants, removing products, changing pressure or
temperature on the equilibrium position of a reaction.
5. Calculate the solubility of an ionic compound given Ksp.

Assignments                     Text Sections          Due Dates
 HW #1                                     12.1
page 343 #2, 4
 HW #2                                     12.2
page 344 #6, 8, 12, 14, 24 & WS           12.3

Mini-Lab

Quiz #1
 HW #3                                     12.4
Worksheet                                 12.5
 HW #4                                     16.1
Worksheet

Quiz #2

Formal Lab

Unit Exam
12.1 Equilibrium Systems (page 321)

( __________________ ) N2O4 (g)  2 NO2 ( ________-__________ )

A. At the beginning of the reaction only the __________ N2O4 exists. As soon as the _______-
____________ NO2 begins to form, the ___________ reaction begins to occur as well.

B. Eventually, the rate of the _________ reaction and the rate of the __________ reaction become
equal. This is called ___________ _____________.

C. At dynamic equilibrium, the concentrations of all products and reactants remain ___________.

D. The relationship between the amount of product and reactant at equilibrium is called the
___________ ____________ ( ___ ) for the system.

E. As the concentration (partial pressure for gases) of the reactants decreases, the partial pressure of
the products _______________. The ratio of the coefficients is proportional to the partial
pressures.

N2O4 (g)  2 NO2 (g)
Times                  0 sec                   1 sec                     2 sec         3 sec
P N2O4                3.0 atm                 2.0 atm
P NO2                  0 atm                                            3.0 atm       3.0 atm

CH4 (g) + 2 O2 (g)  CO2 (g) + 2 H2O (g)
Times                 0 sec                   1 sec                     2 sec         3 sec
P CH4                1.0 atm                 0.5 atm
P O2                2.0 atm                                           0.5 atm
P CO2                 0 atm                                                         0.75 atm
P H2O                 0 atm

12.2 Equilibrium Constant Expressions (page 324)

a A (g) + b B (g)  c C (g) + d D (g)

   The partial pressures of the products appear in the ______________
   The partial pressure of the reactants appear in the _______________
   The pressures are raised to the power of their _________________

K = ________________________________
Examples:
4 NH3 (g) + 6 NO (g)  5 N2 (g) + 6 H2O (g)                 2 NO2 (g) + 7 H2 (g)  2 NH3 (g) + 4 H2O

K = _____________________________                           K = ___________________________

Rules for using Equilibrium Constants

1. K is meaningless unless accompanied by a ____________ _____________

2. If the coefficient in a balanced equation is multiplied by a factor n, the equilibrium constant is
raised to the ______ _______________.

3. The equilibrium constants for forward and reverse reactions are the ____________ of each other.

4. If a reaction can be expressed as the sum of two or more reactions, K for the overall reaction is
the ______________ of the equilibrium constants of the _______________ reactions.

5. Liquids and solids _____ _________ part of an equilibrium expression.

6. Ions in solution ________ part of the equilibrium expression and are entered as their ________
concentration.

Example: If the reaction N2O4 (g)  2 NO2 (g) has a value for K of 11 at 100 C

a. What would be value of K for the decomposition of 3 moles of N2O4?

b. What would be the value of K for the synthesis of half a mole of N2O4?

12.3 Determination of K (page 329)

Equilibrium constants can be calculated using the partial pressure values of the __________ and
______________ in a balanced equation at equilibrium.
Example:
Consider the reaction equation: CO(g) + 2 H2 (g)  CH3OH (g) @ 30 C
The beginning pressure of the CO and the H2 were each 1.50 atm. The equilibrium pressure of the
CO is 1.12 atm. Calculate K at this temperature.

12.4 Applications of the Equilibrium Constant (page 330)

A. A very small value for K means that at equilibrium there are very few _____________. A very
large value for K means that at equilibrium there are very few ____________. A value near "1"
means there are ___________ amounts of both reactants and products.
B. An expression "Q" known as the _________ ____________, can be used to determine if
equilibrium has been reached and in which ________________ the reaction must proceed to reach
equilibrium.
C. "Q" is calculated the same way as "K" is calculated and then compared to "K".
1. If Q < K, the reaction proceeds from _______ to _______ to reach equilibrium.
2. If Q > K, the reaction proceeds from _______ to _______ to reach equilibrium.
3. If Q = K, the reaction is already at ___________.

Example:
Consider the reaction     CO2 (g) + C(s)  2 CO(g) @ 25C, K = 1.17
The partial pressure of CO2 is 0.010 atm and the partial pressure of CO is 0.00015 atm. Determine
the value of Q and determine the direction of the reaction if not at equilibrium.

12.5 Effect of Changes in Conditions Upon an Equilbrium System (page 334)

Le Chatelier's Principle- If a system at equilibrium is disturbed _____________________________
________________________________________________________________________________
________________________________________________________________________________.
A. Change in Concentration
   If a chemical system at equilibrium is disturbed by adding a ____________ product or reactant,
the reaction will proceed in such a direction as to ___________ part of the added species.
   If a gaseous species is removed, the system shifts to ___________ part of that species.
Example:         Consider N2O4 (g)  2 NO2 (g)

1. Adding N2O4 will cause a shift to the _____________
2. Adding NO2 will cause a shift to the ______________
3. Removing N2O4 will cause a shift to the ___________
4. Removing NO2 will cause a shift to the ____________

B. Change in Pressure
   When a system is compressed, increasing the pressure, the shift is in the direction that
___________ the total number of moles of gas.
   When a system is expanded, decreasing the pressure, the shift is in the direction that
___________ the total number of moles of gas.
A decrease in pressure caused an equilibrium shift in which direction?

1. 2 CO2 (g)  2 CO (g) + O2 (g) ____________
2. H2 (g) + I2 (g)  2HI (g)      ____________
3. H2 (g) + I2 (s)  2HI (g)      ____________

C. Change in Temperature
   An increase in temperature causes an ______________ reaction to occur.
   A decrease in temperature would cause the ______________ reaction to occur.
Example:
N2O4 (g)  2 NO2 (g)   H = +57.2

1. An increase in temperature would cause a _____________ shift.
2. A decrease in temperature would cause a ______________ shift.

   Unlike changes in concentration or pressure, changes in temperature cause a change in the value
of _________.
   If the forward reaction is ________________, K _________ as temperature increases.
   If the forward reaction is _______________, K __________ as temperature increases.
16.1 Solubility Product Constant (page 425)

Precipitation reactions, like all reactions, reach a position of ___________________.

A net ionic equation can be written for the equilibrium between a precipitate (solid) and its
corresponding ions in solution (aqueous).

SrCrO4 (s)  Sr2+(aq) + CrO42-(aq)

The corresponding equilibrium expression is as follows:

Ksp = [Sr2+] [CrO42-]

Where Ksp equals the _____________ ____________ _______________ at a particular temperature,
and the values in square brackets equal the ____________ ______________ of the ions.

SrCrO4 does not appear in the equilibrium expression because it is a ____________.

Example 16.1
Write Ksp expressions for Ag2CrO4 and Ca3(PO4)2.

Ksp and the Equilibrium Concentrations of Ions

If the value of Ksp is known, the concentration of one ion at equilibrium can be used to find the
concentration of the other ion.

Example 16.2
Ca3(PO4)2 has a Ksp value of 1 x 10-33. What would be the concentration of PO43- in equilibrium
with the solid if [Ca2+] is 1 x 10-9 M.

Ksp and Precipitation Formation

1. If Q > Ksp, the solution is ________________, and a _____________ forms.

2. If Q < Ksp, the solution is _________________, and ___ ______________ forms.

3. If Q = Ksp, the solution is _________________, and we are at the point of ______________.
Sodium chromate is added to a solution in which the original concentration of Sr2+ is 1.0 x 10-3 M.
Ksp for strontium chromate is 3.6 x 10-5.
a.) will a precipitate form when the concentration of CrO42- = 3.0 x 10-2 M?

b.) will a precipitate form when the concentration of CrO42- = 5.0 x 10-2 M?

Ksp and Water Solubility

Calculate the water solubility of
a.) CaCO3 in mol/L (Ksp = 4.9 x 10-9 )

b.) BaF2 in g/L (Ksp = 1.8 x 10-7)

Ksp and the Common Ion Effect

A common ion comes from at least ______ sources.

An ionic solid is ______ soluble in a solution containing a __________ _____ than it is in _______.
This is consistent with Le Chatelier's Principle.

Example 16.6 (without math)
How would the solubility of CaCO3 in water compare with its solubility in a solution of Na2CO3
solution?

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