1. Welcome Back to Chemistry Standards 101!
Mrs. Michele L. Cramer
Instructor of Chemistry
Albright College

2. Chemistry Standards for 2014-15
HS-PS1-1
Use the periodic table as a model to predict the relative properties of elements based on the patterns of electrons in the outermost energy level of atoms.
Clarification Statement: Examples of properties that could be predicted from patterns could include reactivity of metals, types of bonds formed, numbers of bonds formed, and reactions with oxygen.
HS-PS1-6
Refine the design of a chemical system by specifying a change in conditions that would produce increased amounts of products at equilibrium.*
Clarification Statement: Emphasis is on the application of Le Chatelier’s Principle and on refining designs of chemical reaction systems, including descriptions of the connection between changes made at the macroscopic level and what happens at the molecular level. Examples of designs could include different ways to increase product formation including adding reactants or removing products.

3. Introducing Chemical Equilibrium
Water demonstration
Difference between a reaction that runs to completion and one that is at equilibrium
Video demonstration
Combustion of octane- reaction that goes to completion
Nitrogen dioxide equilibrium- equilibrium reaction

4. The Concept of Equilibrium
Chemical equilibrium occurs when a reaction and its reverse reaction proceed at the same rate.
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5. The Concept of Equilibrium
As a system approaches equilibrium, both the forward and reverse reactions are occurring.
At equilibrium, the forward and reverse reactions are proceeding at the same rate.
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6. A System at Equilibrium
Once equilibrium is achieved, the amount of each reactant and product remains constant.
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7. Depicting Equilibrium
Since, in a system at equilibrium, both the forward and reverse reactions are being carried out, we write its equation with a double arrow.
N2O4 (g)
2 NO2 (g)
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8. The Equilibrium Constant
Forward reaction:
N2O4 (g)  2 NO2 (g) Rate = kf [N2O4]
Reverse reaction:
2 NO2 (g)  N2O4 (g) Rate = kr [NO2]2
Therefore, at equilibrium
Ratef = Rater
kf [N2O4] = kr [NO2]2
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9. The Equilibrium Constant
Rewriting this, it becomes
The ratio of the rate constants is a constant at that temperature, and the expression becomes kf kr
[NO2]2
[N2O4] =
Keq = kf kr
[NO2]2
[N2O4] =
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10. The Equilibrium Constant
Consider the generalized reaction
aA + bB
cC + dD
The equilibrium expression for this reaction would be
Kc =
[C]c[D]d
[A]a[B]b
[products]
[reactants] =
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11. Equilibrium Can Be Reached from Either Direction
As you can see, the ratio of [NO2]2 to [N2O4] remains constant at this temperature no matter what the initial concentrations of NO2 and N2O4 are.
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12. Equilibrium Can Be Reached from Either Direction
This is the data from the last two trials from the table on the previous slide.
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13. Equilibrium Can Be Reached from Either Direction
It doesn’t matter whether we start with N2 and H2 or whether we start with NH3: we will have the same proportions of all three substances at equilibrium.
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14. What Does the Value of K Mean?
If K>>1, the reaction is product-favored; product predominates at equilibrium.
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15. What Does the Value of K Mean?
If K>>1, the reaction is product-favored; product predominates at equilibrium.
If K<<1, the reaction is reactant-favored; reactant predominates at equilibrium.
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16. The Reaction Quotient (Q)
Q gives the same ratio the equilibrium expression gives, but for a system that is not at equilibrium.
To calculate Q, one substitutes the initial concentrations on reactants and products into the equilibrium expression.
[products]init
[reactants]init
Q =
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17. the system is at equilibrium.
If Q = K,
the system is at equilibrium.
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18. there is too much product, and the equilibrium shifts to the left.
If Q > K,
there is too much product, and the equilibrium shifts to the left.
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19. there is too much reactant, and the equilibrium shifts to the right.
If Q < K,
there is too much reactant, and the equilibrium shifts to the right.
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20. Le Châtelier’s Principle
“If a system at equilibrium is disturbed by a change in temperature, pressure, or the concentration of one of the components, the system will shift its equilibrium position so as to counteract the effect of the disturbance.”
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21. The Haber Process
The transformation of nitrogen and hydrogen into ammonia (NH3) is of tremendous significance in agriculture, where ammonia-based fertilizers are of utmost importance.
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22. The Haber Process
If H2 is added to the system, N2 will be consumed and the two reagents will form more NH3.
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23. The Haber Process
This apparatus helps push the equilibrium to the right by removing the ammonia (NH3) from the system as a liquid.
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24. Heat in Chemical Reactions
Endothermic- a process that absorbs heat from the surroundings, heat is put into the reaction
reactants + heat  products DH = +
Exothermic- a process that releases heat to its surroundings
reactants  products + heat DH = -

25. Heat in Equilibria Treat heat as you would another reactant or product
If endothermic, heat is a reactant
reactants + heat products DH = +
If exothermic, heat is a product
reactants products + heat DH = -

26. The Effect of Changes in Temperature
Co(H2O)62+(aq) + 4 Cl(aq)
CoCl4 (aq) + 6 H2O (l)
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27. LeChatelier’s Principle Examples
Cr2O7-2(aq) + H2O CrO4-2(aq) + 2H+(aq)
Blue bottle demo

28. Using Le Châtelier’s Principle to Predict Shifts in Equilibrium
Consider the equilibrium
In which direction will the equilibrium shift when (a) N2O4 is added, (b) NO2 is removed, (c) the temperature is decreased?

29. Using Le Châtelier’s Principle to Predict Shifts in Equilibrium
Consider the equilibrium
In which direction will the equilibrium shift when (a) Cl2(g) is removed, (b) the temperature is decreased?
What could you do to increase the production of PCl5?