1. Day 3 Chemical Equilibrium
Lecture Presentation
Unit 7
Day 3 Chemical Equilibrium
James F. Kirby
Quinnipiac University
Hamden, CT
Edited by M. Day

2. Warm Up R [HA] [A-] [H3O+] I 0.650 ? C E 0.250
Consider the following reaction:
HA + A → H3O+
Fill in this RICE table R
[HA]
[A-]
[H3O+] I
0.650 ? C E
0.250

3. Agenda Equilibrium Phet Lecture: Q vs. K LeChatelier’s Principle
Practice Problems

4. Is a Mixture in Equilibrium? Which Way Does the Reaction Go?
To answer these questions, we calculate the reaction quotient, Q.
Q looks like the equilibrium constant, K, but the values used to calculate it are the current conditions, not necessarily those for equilibrium.
To calculate Q, one substitutes the initial concentrations of reactants and products into the equilibrium expression.

5. Comparing Q and K Nature wants Q = K.
If Q < K, nature will make the reaction proceed to products.
If Q = K, the reaction is in equilibrium.
If Q > K, nature will make the reaction proceed to reactants.

6. Calculating Equilibrium Concentrations
If you know the equilibrium constant, you can find equilibrium concentrations from initial concentrations and changes (based on stoichiometry).
You will set up a table similar to the ones used to find the equilibrium concentration, but the “change in concentration” row will simple be a factor of “x” based on the stoichiometry.

7. An Example
A L flask is filled with mol of H2(g) and mol of I2(g) at 448 °C. Given a Kc of 50.5 at 448 °C, what are the equilibrium concentrations of H2, I2, and HI? H2(g) + I2(g) ⇌ 2 HI(g)
initial concentration (M)
1.000
2.000
change in concentration (M) –x
+2x
equilibrium concentration (M)
1.000 – x
2.000 – x 2x

8. LeChâtelier’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.”

9. How Conditions Change Equilibrium
We will use LeChâtelier’s Principle qualitatively to predict shifts in equilibrium based on changes in conditions.

10. How Conditions Change Equilibrium
We will use LeChâtelier’s Principle qualitatively to predict shifts in equilibrium based on changes in conditions.

11. Example (continued)
Set up the equilibrium constant expression, filling in equilibrium concentrations from the table.
Solving for x is done using the quadratic formula, resulting in x = or

12. Example (completed)
Since x must be subtracted from M, makes no physical sense. (It results in a negative concentration!) The value must be So
[H2]eq = – = M
[I2]eq = – = M
[HI]eq = 2(0.935) = 1.87 M

14. Warm Up
What three conditions did we manipulate in the lab that shifted the equilibrium of the five systems tested?

15. Lab Report: Procedure, Data, Discussion of Theory A
Lab Report: Procedure, Data, Discussion of Theory A. HIn (aq) ⇌ H+ (aq) + ln-(aq) H+(aq) + OH-(aq)  H2O(l)
Procedure
Observations
Explanation
Initial color of water and bromothymol blue
Solution is green
The green color shows that the pH of distilled water is between 6 and 7.6, and the indicator is a mix of Hln and I-
Add 0.1 M HCl
Solution turned yellow
pH < 6.0
Add 0.1 M NaOH
Solution turned green and finally ended on blue, pH > 7.6
Additional drops of HCl and NaOH
Solution changes from yellow (with HCl) to blue (with NaOH)

16. Lab Report: Procedure, Data, Discussion of Theory B
Lab Report: Procedure, Data, Discussion of Theory B. Cu2+(aq) + 4NH3 (aq) ⇌ [Cu(NH3)4]2+(aq) H+(aq) + NH3(aq)  NH4+(aq)
Procedure
Observations
Explanation
Initial color of copper solution
Solution is blue
Add concentrated NH3 drops
Solution turned light blue and solid formed; with more NH3, solid dissolved and turned deep blue
Add 1.0 M HCl
Deep blue faded to lighter blue and solid formed again
Additional drops of NH3
Solid dissolved again and solution turned deep blue

17. Lab Report: Procedure, Data, Discussion of Theory C
Lab Report: Procedure, Data, Discussion of Theory C. [Co(H2O)6]2+(aq) + 4HCl (aq) + Heat ⇌ [CoCl4]2-(aq) + 6H2O(l) Ag+(aq) + Cl-(aq)  AgCl(s)
Procedure
Observations
Explanation
Add 6.0 M HCl drops (tube A)
Solution turned blue
Add 0.1 M AgNO3 drops (tube B)
White solid precipitate and solution turned pink
Add distilled water drops (tube C)
Solution turned pink
Add 5-6 grains CaCl2 to tube C
Crystals dissolved and solution turned blue
Test tube C placed in ice water bath for 2-3 minutes
Test tube C placed in hot bath for 2-3 mins

18. Lab Report: Procedure, Data, Discussion of Theory D
Lab Report: Procedure, Data, Discussion of Theory D. 2CO2 (g) + H2O (l) ⇌ CO2 (aq) + H+(aq) + HCO3-(aq)
Procedure
Observations
Explanation
Initial color of solution
Green color = pH 4.4
Pull back on syringe to decrease pressure
Solution turned teal, pH = 4.8
Push syringe to increase pressure
Solution turned green, pH = 4.4

19. Lab Report: Procedure, Data, Discussion of Theory E
Lab Report: Procedure, Data, Discussion of Theory E. Mg(OH)2 (s) ⇌ Mg2+(aq) + 2OH-(aq) H+(aq) + OH-(aq)  H2O (l)
Procedure
Observations
Explanation
Initial color of milk of magnesia and universal indicator
Purple solution with white solid suspended in liquid
White solid is undissolved Mg(OH)2. Purple color is from the universal indicator and shows that some of the OH- ions are present b/c purple = pH > 10
Add 1 drop of 3 M HCl with constant stirring
Solution immediately turned pink with more stirring pink color turned orange, green, then blue
Additional drops of 3 M HCl
Solution immediately turned pink; slower change to the blue-green end color
Additional drops of HCl and NaOH
Solution immediately turned pink; color remained pink and solution not cloudy

20. Change in Reactant or Product Concentration
If the system is in equilibrium
adding a reaction component will result in some of it being used up.
removing a reaction component will result in some if it being produced.

21. Change in Volume or Pressure
When gases are involved in an equilibrium, a change in pressure or volume will affect equilibrium:
Higher volume or lower pressure favors the side of the equation with more moles (and vice-versa).

22. Change in Temperature
Is the reaction endothermic or exothermic as written? That matters!
Endothermic: Heats acts like a reactant; adding heat drives a reaction toward products.
Exothermic: Heat acts like a product; adding heat drives a reaction toward reactants.

23. An Endothermic Equilibrium

24. An Exothermic Equilibrium
The Haber Process for producing ammonia from the elements is exothermic.
One would think that cooling down the reactants would result in more product.
However, the activation energy for this reaction is high!
This is the one instance where a system in equilibrium can be affected by a catalyst!

25. Catalysts
Catalysts increase the rate of both the forward and reverse reactions.
Equilibrium is achieved faster, but the equilibrium composition remains unaltered.
Activation energy is lowered, allowing equilibrium to be established at lower temperatures.

26. END DAY 2 LECTURE