Name: _________________________
Chemistry 30
Unit 3:
Equilibrium and Acid-Base
April 2019
Key Concepts:
The following concepts are developed in this unit and may also be
addressed in other units or in other courses. The intended level and scope of treatment is defined by the outcomes
chemical equilibrium systems
Brønsted–Lowry acids and bases
reversibility of reactions
Le Chatelier’s principle
conjugate pairs of acids and bases
equilibrium law expression
amphiprotic substances
equilibrium constants Kc , Kw , Ka , Kb
buffers
acid-base equilibrium
indicators
titration curves
1
Part A: Chemical EQUILIBRIUM
Chemical Systems
Assumptions of chemical systems we often make …
all chemical reactions are fast
they are all spontaneous
they are all quantitative (go to completion; 100% products and no limiting reactant(s) left)
they are stoichiometric (we can predict the amounts of products, since they go to completion)
In actual fact, many chemical reactions do not go to completion.
A state of dynamic equilibrium exists.
Microscopic properties – are continually changing (at the molecular level there is a great deal of change)
Recap on Systems
Open System – matter and energy are exchanged
Closed System – no matter is exchanged, but energy can go in and out
Isolated System – no matter or energy is exchanged
How does this cartoon represent dynamic equilibrium? What does the fence represent? Could the piles of soil be the same size?
Different sizes?
Macroscopic properties – are constant (examples of macroscopic properties include observable properties, such as temperature, colour and pH
Dynamic equilibrium
There is a balance between the forward and reverse reactions that are occurring at the same rate so the observable properties are constant.
How do we show dynamic equilibrium?
Example of an equilibrium equation:
2 4(aq) 2(aq) 2 (aq) 4( )s
Na SO CaCl NaCl CaSO
Collision theory explains that reverse reaction can occur (products can react to re-form the initial reactants). The final state of the reaction is a competition between the forward and reverse reactions
A system at Dynamic Equilibrium:
The [Reactants] and [Products] will not change once equilibrium is reached, even though the forward and reverse reactions are happening simultaneously
The macroscopic properties are constant, but a great deal of change at the microscopic/molecular level
Requirements to stay at equilibrium:
1) Temperature must remain constant, (if the Temp changes, Kc changes) 2) The system must remain closed.
Both physical and chemical changes can reach equilibrium.
Here are three examples of types of equilibrium situations 1. Solubility equilibrium - Saturated solutions with excess solid.
Example: saturated salt water (at a certain temp)
( )s 2 ( )l (aq) (aq) 2 ( )l
NaCl H O Na Cl H O
What remains constant? Temperature, Concentration ([Na+] & [Cl-]), mass of NaCl(s). (Shape of crystals changes constantly)
2. Phase / state equilibrium
Example: mixture of ice and water at 0°C:
H2O(s )↔H2O(l )
What remains constant: temperature; water volume; mass ice (shape of ice is changing) 3. Chemical reactions at constant temperature
Examples:
1. NO(g ) + 12 O2( g ) ⃗at 25oC NO2 ( g) ΔH=57 kJ
What remains constant?: temperature; pressure
2. Carbonated beverages: H CO2 3(aq) H O + CO2 (l) 2 g
All of the above are examples of a dynamic equilibrium which means there is a balance between the forward and reverse processes (at the same rate)
homogeneous expression : same state (aqueous, gas or liquid state…solid state is not useful in equilibrium )
heterogeneous expression : different state
An equilibrium law expression constant is the product of the molar concentrations of the products for a chemical reaction, each raised to the power of its respective coefficient in the equation, divided by the product of the molar concentrations of the reactants, each raised to the power of its respective coefficient in the equation.
RULES:
1. Concentrations are always expressed in moles per litre.
2. Product concentrations are always placed in the numerator of the equilibrium law expression.
3. Reactant concentrations are always placed in the denominator of the equilibrium law expression.
4. The powers to which concentrations are raised are always determined by the coefficients in the balanced equation for the reaction.
How to Determine an Equilibrium Law Expression and Calculate an Equilibrium Constant Example:
The following reaction occurs in a 500 mL container:
at equilibrium, there is 0.70 mol of C2H2, 0.80 mol of O2, 0.90 mol of CO2, 0.50 mol of H2O in a sealed 500mL container
Write out the equilibrium law expression and …
Then determine equilibrium constant, Kc (first: calculate the molar concentration of each reactant and product).
Equilibrium Law Expression Rules 1. What must ALWAYS be included?
A. Gases
B. Aqueous solutions
2. What must NEVER be included?
A. Solids
3. What is SOMETIMES included?
A. Liquids
2 2( ) 2( ) 2( ) 2 ( )
2 C H g 5 O g 4 CO g 2H O g
If more than one liquid is in the equilibrium equation, they are ALL included in the expression
If only one liquid is in the equilibrium equation, it is NOT included in the expression.
EQUILIBRIUM QUESTIONS
1. List three properties of chemical systems at equilibrium.
a. _______________________________________
b. _______________________________________
c. _______________________________________
2. Why is chemical equilibrium considered to be dynamic?
3. What is meant when we say that the “macroscopic properties” of a system at equilibrium are constant”?
4. Define the following:
a. Constant:
b. Dynamic:
c. Static
d. Rate:
Reminders:
Only concentrations of gases and substances in solution are written in an equilibrium law expression.
A liquid alone in an equation must be excluded from the expression but several liquids in the same equation must be included.
5. Write equilibrium law expressions for each of the following reactions:
Equilibrium Equation Equilibrium Law Expression, Kc
(a) 4NH3(g) + 3O2(g) 2N2(g) + 6H2O(g)
Kc
(b) 6Ca(s) + 2NH3(g) 3CaH2(s) + Ca3N2(s)
(c) 2Ag2CO3(s) 4Ag(s) + 2CO2(g) + O2(g)
(d) NaCl(aq) + AgNO3(aq) AgCl(s) + NaNO3(aq)
(e) CO(g) + H2O(g) CO2(g) + H2(g)
(f) N2H4(g) + H2O(l) N2H5+
(aq) + OH-(aq)
(g) 2NaHCO3(s) Na2CO3(s) + H2O(g) + CO2(g)
(h) 2H2(g) + O2(g) 2H2O(l)
(i) C6H6(l) + Br2(l) C6H5Br(l) + HBr(g)
(j) C2H5OH(l) + C2H5COOH(l) C2H5COOC2H5(l) + H2O(l)
Problems Involving Calculating the Equilibrium Constant, Kc
6. At a temperature of 350oC, the equilibrium concentrations for the reaction N2(g) + 3H2(g) 2NH3(g)
[N2] = 0.885 mol/L [H2] = 0.665 mol/L [NH3] = 1.230 mol/L Calculate the value of the equilibrium constant for this reaction at 350oC.
7. The unique flavor of fruits and berries is due to the presence of esters, chemical compounds that are easily synthesized in the laboratory. For example ethyl propanoate is responsible for the flavour characteristic to pineapple. It is produced by the reaction
C2H5OH(l) + C2H5COOH(l) C2H5COOC2H5(l) + H2O(l)Kc = 2.2
What must be the concentration of water be for a mixture with [C2H5COOC2H5] = 2.0 mol/L, [C2H5COOH] = 0.10 mol/L, [C2H5OH] = 5.0 mol/L to be at equilibrium?
8. 2SO2(g) + O2(g) 2SO3(g)
The above reaction occurred in a 0.500 L flask at constant temperature. At equilibrium, the flask contained 3.00 moles of SO3(g), 2.25 moles of SO2(g), and 100.0 g of O2(g). Calculate the equilibrium constant for the reaction at flask temperature.
9. The equilibrium constant for the hypothetical reaction below is 10
A(aq) + B(aq) C(aq) + D(aq) Kc = 10 What would be the equilibrium constant for the reverse reaction?
C(aq) + D(aq) A(aq) + B(aq) Kc = ?
10. For the reaction below Kc is equal to 2.1 x 103 at a particular temperature.
H2(g) + F2(g) 2HF(g)
When the system is analyzed at equilibrium at this temperature, the concentrations of H2(g) and F2(g) are both found to be 0.0021 mol/L. What is the concentration of HF(g) in the equilibrium system under these conditions ?
11. What is the only change we can impose on an equilibrium system that will result in a different equilibrium constant?
12. Complete the following table: H2(g) + I2(g) 2HI(g)
Equilibrium Conditions for Hydrogen-Iodine System at 448oC Equilibrium
Conditions
[H2(g)]
(mol/L)
[I2(g)]
(mol/L)
[HI(g)]
(mol/L)
A 0.40 0.50 1.00
B 0.16 8.5
C 0.333 3.750
D 3.56 22.4
E 4.0 9.0
13. Write a balanced equilibrium equation for a totally gaseous equilibrium system that would lead to the following expressions for the equilibrium constant.
a.
2 ( )
2( ) 2 2( )
g c
g g
K HCN
H C N
b.
2
4( ) 2 ( )
4
2( ) 2( )
g g
c
g g
CH H S
K CS H
c.
2 ( ) 2
( ) 2( )
g c
g g
K NOBr
NO Br
d.
2
2( ) ( )
2 ( )
g g
g
Br NO
K NOBr
14. Multiple Choice question:
Given the following Kc expression:
Kc = [Cl-(aq)][I2(aq)]
[OCl-(aq)][H+(aq)]2[I-(aq)]2
The reaction represents the balanced equation that would lead to the Kc expression is?
a. OCl-(aq) + 2 H+(aq) + 2 I-(aq) ↔ Cl-(aq) + I2(aq)
b. Cl-(aq) + I2(aq) ↔ OCl-(aq) + 2 H+(aq) + 2 I-(aq)
c. OCl-(aq) + 2 H+(aq) + 2 I-(aq) ↔ Cl-(aq) + I2(aq) + H2O(l)
d. Cl-(aq) + I2(aq) + H2O(l) ↔ OCl-(aq) + 2 H+(aq) + 2 I-(aq)
15. Consider the following system at equilibrium:
PCl5(g) PCl3(g) + Cl2(g) Kc = 4.2 x 10-2
If the equilibrium concentration of PCl5(g) is 12 mol/L and that of Cl2(g) is 49 mol/L, what is the equilibrium concentration of PCl3(g) ?
16. Which of the following changes would affect the value of a system's equilibrium constant?
(a) addition of a reactant or product (b) increase in the total pressure
(c) addition of an inert gas (d) decrease in the temperature
17. For the hypothetical reaction: 2X(g) + 3Y(g) 2Z(g) the Kc = 0.145 at 45.0oC.
What is the value of Kc (equilibrium constant) for this equilibrium equation?
2Z(g) 2X(g) + 3Y(g)?
18. The following results were collected for two experiments involving the reaction at 600oC between gaseous sulfur dioxide and oxygen to form gaseous sulfur trioxide.
2SO2(g) + O2(g) 2SO3(g)
Experiment 1 Experiment 2
Initial Equilibrium Initial Equilibrium
[SO2] = 2.00 mol/L [SO2] = 1.50 mol/L [SO2] = 0.500 mol/L [SO2] = 0.590 mol/L [O2] = 1.50 mol/L [O2] = 1.25 mol/L [O2] = 0 mol/L [O2] = 0.0445 mol/L [SO3] = 3.00 mol/L [SO3] = 3.50 mol/L [SO3] = 0.350 mol/L [SO3] = 0.260 mol/L
a) Calculate the equilibrium constant (Kc) for experiment 1 and experiment 2.
b) Why are the values for Kc the same?
19. The equilibrium constant, Kc, for the following reaction is 82 at 450oC.
N2(g) + 3H2(g) 2NH3(g)
Calculate the number of moles of nitrogen present at equilibrium when a 8.00L vessel showed the presence of 4.0 moles of hydrogen and 5.0 moles of ammonia.
Ways to Communicate Equilibrium Position
There are 4 ways to talk about equilibrium position:
1. Words (see columns in table below)
2. Arrows (this can include percent reaction) (see last column in table below) 3. Equilibrium expressions and the equilibrium constant (Kc)
4. Graphs
Law of Equilibrium and the Magnitude of Kc
What’s Favoured? Concentration of products and reactants at
equilibrium
Arrows and percent reaction Products are favoured
Or
Forward reaction is favoured [products] > [reactants]
Reactants are favoured Or
Reverse reaction is favoured [reactants] > [products]
Reaction is quantitative More than 99% are products, almost no reactants
Equilibrium can be quantitative to reactant side if no appreciable product is produced.
Remember that in a chemical equilibrium, there’s a constant ratio between the concentration of products and the concentration of reactants. The equilibrium constant, K, is the ratio of products to reactants.
Therefore, the larger Kc the more products are present at equilibrium.
Conversely, the smaller Kc the more reactants are present at equilibrium.
• If K >> 1, then products could dominate at equilibrium and equilibrium favours the products / lies to the right.
• If K << 1, then reactants could dominate at equilibrium and the equilibrium favours the reactants / lies to the left.
• If K≈1, then the concentration of products and reactants could be the same concentration.
Examples:
Ok, this is just a math question…no chem knowledge needed….
1. Put these numbers in numeric order:
1) 2.2x10-5 2) 1.4x10-4 3) 3.5x10-5 4) 1.5x10-4 5) 2.2x103 6) 2.8x10-6
2. Place these Kc values in order from most favouring products to least favouring products
1) Kc = 2.2x10-8 2) Kc = 2.1x10-8 3) Kc = 3.5x10-7 4) Kc = 1.7x102 5) Kc = 5.9x10-5 6) Kc = 5.82x10-3
Answer: Answer:
Limitations of Equilibrium Law Expressions and Equilibrium constants
They do not provide information on rate of reaction, only relative amounts of products and reactants at equilibrium.
Kc is affected by temperature
Percent reaction has a different value for every change in initial concentrations of reactants (often only used for 0.1mol/L solutions)
Equilibrium graphs
1. x-axis = time or reaction coordinate
2. y-axis = concentration of reactants and products (usually in mol/L) 3. The concentration of reactants is decreasing and the concentration of
products increasing as the rate of increase and decrease are proportional to coefficients of equation
4. Equilibrium is reached if all lines are parallel
Interpreting Equilibrium Graphs - What can we be expected to determine from the
graph?
a) What time was equilibrium established?
b) What is favoured (products or reactants)? (Which one do you make more of?)
c) What is value Kc?
20. Rank the following equilibrium constants in order from the one most favoring the products to the one least favoring the products:
K1 = 6.4 x 10-3 K2 = 3.9 x 10-2 K3 = 9.2 x 10-4 K4 = 3.0 x 10-1 K5 = 4.2 x 10-2
Trivia Time Eggs and Equilibrium: Feeding chickens carbonated water?
Eggshells are made of calcium carbonate (CaCO3(s) ), which is made from carbon dioxide (CO2), a product of cellular respiration. The net equation is
When chickens become hot, they pant, which decreases the concentration of carbon dioxide in the blood. To offset the stress, the equilibrium will shift in the reverse direction and decrease the amount of calcium carbonate available to make eggshells. This yields eggs with thin shells that break easily. Ted Odom, a graduate student at the University of Illinois, found that giving chickens carbonated water to drink will shift equilibrium in the forward direction and minimize the effects of panting on warm days. This allows farmers to minimize the effects without having to install expensive air conditioning in chicken coops (van Kessel, et al. 457).
Fritz Haber - "father of chemical warfare” for his work developing and deploying chlorine and other poisonous gases during World War I. Considered a war criminal by many. For some time after the war, he was afraid that charges would be pressed against him.
He is reported to have said, “During peace time a scientist belongs to the world, but during war time he belongs to his country”
Haber hated the war but hoped that in developing the gases he would help to bring it to a speedy end by breaking the deadlock of trench warfare. His wife, however, denounced his work as a perversion of science.
After a violent argument with Haber in 1915, she committed suicide.
And yet….he gets a Nobel Prize in Chemistry….. Because of the “Haber Process” that allows us to make nitrogen bases fertilizers….doubles the worlds food supply.
The famous process : N2(g) + 3H2(g) 2NH3(g) + 92 kJ
Use the following information to answer the next question:
= N2
= H2
= NH3
10
9
8 7
6
5
4
3
2
1
0
0 4 8 12 16 20 24 28 32 36 40
Time (min)
21. In terms of the chemical equilibrium represented in the above graph:
Legend Formation of ammonia
Concentratio n
mol/L
(a) Write out the balanced chemical equation
(b) Determine what time equilibrium was reached.
(c) Indicate whether the reactants or products are favored.
(d) Calculate the equilibrium constant
22. The Haber-Bosch process for the industrial production of ammonia involves the equilibrium.
N2(g) + 3H2(g) 2NH3(g) + 92 kJ
In a laboratory experiment designed to study this equilibrium, a chemical engineer injects 0.20 mol of N2(g) and 0.60 mol of H2(g) into a 1.0 L flask at 500oC. She records her analysis of the contents of the flask at 5 minute intervals in the table shown.
Time(min) [N2(g)] [H2(g)] [NH3(g)]
0 0.20 0.60 0.00
5 0.14 0.42 0.12
10 0.11 0.33 0.18
15 0.10 0.30 0.20
20 0.10 0.30 0.20
25 0.10 0.30 0.20
Analyze the data. Your responses should include:
a plot of the concentrations of N2(g), H2(g), and NH3(g) versus time on the grid provided, including an appropriate title
the time required to establish equilibrium
the equilibrium constant for the reaction
two methods, other than increasing the amounts of reactants, that could increase the yield of ammonia
Manipulating Equilibrium Systems
For any closed chemical system the tendency is to move toward a state of equilibrium.
Le Chatelier’s Principle
When a system at equilibrium is disturbed by a change in a property of the system it adjusts in a way that opposes that change.
If stress is applied to a system at equilibrium, the system changes at relieve that stress.
The application of the principle is a 3 step process:
1. An initial equilibrium state 2. A shifting, non-equilibrium state 3. A new equilibrium state
Stresses that shift equilibriums:
Adding/removing reactants/products
Changing pressure
Changing volume
Changing the temperature Shifts to Equilibrium:
Right shift: the concentration of the product increases because the forward reaction is momentarily favoured
Left shift: the concentration of the reactant increases because the reverse reaction is momentarily favoured
A. Concentration Changes
If we add more of a substance, the system will shift to consume some of this added substance.
o Therefore … if we add product, we shift to the reactant side; if we add reactant, we shift to the product side
If we remove some substance, the system will shift to produce more of this removed substance.
o Therefore … if we remove product, we shift to the product side; if we remove reactant, we shift to the reactant side
Then the system will achieve a new equilibrium.
Kc is unaffected.
Example: 3NO2( )g H O2 ( )g 2HNO3(aq) NO( )g
i. If the concentration of NO(g) is increased, the system will shift to the ________________________ and the concentration of the products will _____________ and the concentration of the reactants will
______________.
ii. If the concentration of ___________ is decreased, the system will shift to the ________________________
and the concentration of the products will _____________ and the concentration of the reactants will ______________.
B. Energy/Temperature Changes
If we add heat the system will shift to remove some of the added heat.
o Therefore if we heat an endothermic reaction shifts to the products o …. If we heat an exothermic reaction shifts to the reactants
If we remove heat (cool the system) it will shift to produce more heat.
o Therefore if we cool an endothermic reaction shifts to the reactants o …. If we cool an exothermic reaction shifts to the products
the value of Kc will be affected (Kc will either increase or decrease depending on whether the products or reactants concentration increased)
Example: Consider the Haber process: N2( )g 3H2( )g 2 NH3( )g H
i. If the system is heated, the system will shift to the ________________________ and the concentration of the products will _____________ and the concentration of the reactants will ______________. Kc will __________________...why?
ii. If the system is cooled, the system will shift to the ________________________ and the concentration of the products will _____________ and the concentration of the reactants will ______________. Kc
will __________________...why?
C. Volume/Pressure Changes of Gases
For gaseous systems, changing the volume changes the concentration of every gas involved in the system.
To predict the effect of changing pressure you need to see how many moles of gas are on the product side and how many moles of gas are on the reactant side.
Pressure and Volume – A Review
Pressure and volume are inversely related. (Boyle’s Law) As pressure goes up, volume goes down and vice versa. Also remember that with ideal gases, the size of the gas molecule doesn't really matter.
Only the spaces between the gas molecules count.
So the volume of N2(g) = H2(g) = NH3(g) even though they are different masses
if we increase volume (decrease pressure) the system will shift towards the side with the most amount of moles
if we decrease volume (increase pressure) the system will shift towards the side with the least amount of moles
Kc is unaffected.
Example:
2( )g 3 2( )g 2 3( )g
N H NH energy
i. If the pressure is increased, the system will shift to the ________________________
ii. If the pressure is decreased, the system will shift to the ________________________
Can you think of a scenario when there would be no shift even though pressure / volume has been changed?
D. Addition of a Catalyst or Inert Gas do NOT cause equilibrium shifts
Add a catalyst: Catalysts allow a system to reach equilibrium faster but, do not change the equilibrium constant, or change equilibrium concentrations.
Add an Inert gas: Inert gases are noble gases. When we add them to an equilibrium mixture, they cause no shift. They do not change equilibrium concentrations, or equilibrium constants.
Le Chatelier summary
So what can actually change K?
Only Temperature changes K!
When concentrations are changed, equilibrium is restored through chemical shifts.
Altering the pressure may seem to alter equilibrium, but in truth it does not. This is due to partial pressures (University concept and chem 20 but we don’t really care about that now do we?).
A few In class Demos: Testing LeChatelier’s Principle
#1. The following equilibrium is observed:
22 2 2
4( )al 6 ( )al 6( ) 4 ( )al
CoCl H O Co H O al Cl energy
blue pink
Problem: How does changing the temperature affect this equilibrium system?
Prediction
(answer the problem based on LeChatelier Principle)
Initial (describe solution color, etc) and Final Observations (describe the final solution color, etc)
Heat the system by placing in hot water Cool the system by placing in ice water
Problem: How does changing the concentration affect this equilibrium system?
Prediction
(answer the problem based on LeChatelier Principle)
Final Observations
Increase Cl( )al by adding NaCl( )s
Be very specific as to your final observations here.
Decrease Cl( )al by adding AgNO3( )s
Be specific as to your final observations.
What substance is produced?
#2.
N2O4( g ) ↔2 NO2( g)
(colorless) (red−brown)
When cooled, the tube turns more ________________, which means it shifted to the ________
When heated, the tube turns more ________________, which means it shifted to the ________
This means the above equilibrium system is _____________________ (endo or exo thermic?)
Le Chatelier Principle Questions
23. Consider the equilibrium equation:
2CrO42-
(aq) + 2H+(aq) Cr2O72-
(aq) + H2O(l) (a) What would happen if you added HCl(aq)to the above reaction ?
(b) What would happen if you added Cr2O72-
(aq) to the above reaction?
24. Consider the reaction:
PCl5(g) + 94 kJ PCl3(g) + Cl2(g) Explain the effect the following changes would have upon the equilibrium mixture:
increased temperature
addition of Cl2(g)
addition of PCl5(g)
presence of a catalyst
EQUILIBRIUM DEMONSTRATION: When a few drops of Fe(NO3)3(aq) solution are added to KSCN(aq) solution the following equilibrium is established:
Fe3+(aq + SCN-(aq) FeSCN2+(aq)
(pale yellow) (colorless) (red)
25. For each of the following situations, predict the shift in equilibrium:
a. A crystal of KSCN(s)is added to the equilibrium mixture
b. A few drops of dilute NaOH(aq)are added to the equilibrium mixture.
If CoCl2.6H2O(s)is dissolved in a saturated NaCl(aq)solution, the following equilibrium is established: Co(H2O)62+(aq) + 4Cl-(aq) + heat CoCl42-(aq) + 6H2O(l)
pink blue
26. The equilibrium mixture (pink at room temperature) is warmed gently over a Bunsen burner flame. Predict any colour changes.
27. Consider the equilibrium
H2PO4-(aq) + HSO3-(aq) H2SO3(aq) + HPO42-(aq) + 23 kJ An imposed change that would cause the equilibrium to shift to the left would be
A. increasing the concentration of HPO42-(aq) B. decreasing the temperature
C. decreasing the concentration of H2SO3(aq) D. increasing the concentration of H2PO4-(aq)
In an equilibrium system involving gases, the total pressure in the system can be changed by increasing or decreasing the volume of the system. The formation of ammonia is represented by the following equation:
N2(g) + 3H2(g) 2NH3(g)
+ <---->
4 volumes 2 volumes
If the pressure in the reaction vessel is increased (decreased volume), the system will try to decrease the pressure by decreasing the number of molecules inside the container.
This means a shift to the right.
Use the equilibrium example above to answer the next 6 questions:
28. In which direction will this system shift if pressure is increased by decreasing the size of the reaction vessel?
29. Which direction will this system shift if pressure is decreased by increasing the size of the reaction vessel?
30. What will happen to this system if a xenon gas is added?
31. How does a change in pressure affect a system with the same number of molecules on both sides of the equation?
32. How does the addition of a catalyst affect a system already at equilibrium?
33. How does the addition of a catalyst affect a system not yet at equilibrium?
34. Methane, CH4, reacts with limited oxygen and releases heat according to the equation 2CH4(g) + O2(g) 2CO(g) + 4H2(g) + heat
Predict the direction of the equilibrium shift for each of the following stresses (a) [CH4] increases
(b) [CO] increases (c) [O2] decreases (d) [H2] decreases
(e) temperature increases (f) pressure decreases (g) volume decreases
35. Traces of formaldehyde in smog are responsible for an eye-burning sensation.
Formaldehyde, CH2O, results from the reaction of ozone and the hydrocarbon pollutant ethylene, C2H4, as follows:
2C2H4(g) + 2O3(g) 4CH2O(g) + O2(g) + heat Predict the direction of equilibrium shift for each of the following stresses:
(a) [C2H4] increases (b) [O3] increases (c) [O2] decreases
(d) temperature increases (e) pressure decreases (f) volume decreases (g) addition of a catalyst
36. Consider the equilibrium situation:
CO(g) + H2O(g) + heat CO2(g) + H2(g)
For each of the following changes initiated on the above system, indicate the effect on the position of the equilibrium. (shift left, shift right, no effect)
(a) Heat the equilibrium mixture ______________
(b) Add CO(g) to the equilibrium mixture ______________
(c) Add H2(g) to the equilibrium mixture ______________
(d) Remove CO2(g) from the equilibrium mixture ______________
(e) Add a catalyst to the equilibrium mixture ______________
(f) Increase the size of the reaction flask ______________
37. What effect does the addition of heat have on each of the following equilibrium reactions?
Indicate the direction in which the equilibrium shifts.
(a) 3Fe(s) + 4H2O(l) + 151 kJ Fe3O4(s) + 4H2(g) ________
(b) 4NH3(g) + 5O2(g) 4NO(g) + 6H2O(l) + 909 kJ _______
38. For the following reactions indicate the effect on the position of the equilibrium of:
i lowering the temperature ii adding nitrogen iii increasing the pressure
(a) N2(g) + 3H2(g) 2NH3(g) ∆H = negative i ____________________
ii ____________________
iii____________________
(b) N2(g) + O2(g) 2NO(g) ∆H = positive i ___________________
ii ____________________
iii____________________
39.
You are given the following chemical system at equilibrium:
CO2(g) + H2(g) CO(g) + H2O(g) ∆H = 42 kJ Complete the following table using Le Chatelier's Principle.
Stress Direction of Shift Change in [CO] Change in Kc
Decrease volume of container Add heat
Remove H2O Add CO2
Put reaction vessel in ice Add a catalyst
Increase volume of container Remove hydrogen
Use the following information to answer the next three questions I SO2(g) + Cl2(g) SO2Cl2(g) + energy
II CH4(g) + H2O(g) + energy CO(g) + 3H2(g)
III 2SO2(g) + O2(g) 2SO3(g) + energy
IV 3O2(g) + energy 2O3(g)
V H2(g) + I2(g) + energy 2HI(g)
40. Which of the systems would shift towards the products with the addition of heat
A. I, III C. II, IV, V
B. II D. V
41. The systems that would shift left with an increase in pressure would be ?
A. I, II, IV C. IV
B. II D. V
42. The system whose equilibrium would not shift because of any change in pressure is:
A. I C. IV
B. III D. V
Graphing Le Chatelier's Principle
Example: The following graph shows the concentrations of certain substances as various stresses are exerted on the system. Points A, B, C and D could be one of the stresses listed in the key.
Cl2(g) + 2 O2(g) ⇌ 2ClO2(g) + heat
Answer: A = _____ B = _____ C = _____ D = _____
Key:
1. Addition of heat 2. Cooling the system 3. Increase the pressure 4. Decrease the
pressure
5. Removal of chlorine 6. Removal of oxygen 7. Addition of a catalyst 8. Addition of oxygen
Le Chatelier Graphs: A graph can be used to represent the changes that occur when a system at equilibrium is upset.
I. Energy Change Effects on the Hydrogen-Nitrogen-Ammonia System
N2(g) + 3H2(g) 2NH3(g) + 92 kJ
A B C
N2(g)
H2(g)
Concentration
NH3(g)
Reaction Progress
A. The system is at equilibrium. The forward and reverse reaction rates are equal. The concentration of each substance remains constant.
B. Heat energy is added to the system. The system shifts according to Le Chatelier’s Rule.
C. A new equilibrium is established
II. Pressure Change Effects on the Hydrogen-Nitrogen-Ammonia System 43. Fill in the blanks:
The __________(production/decomposition) of ammonia results in a decrease in the total number of moles in a closed system and therefore a __________ (decrease/increase) in the total pressure of the system.
When the pressure of an ammonia system is increased, the equilibrium will shift to _______
(decrease/increase) the total pressure and will cause the _________ (production/decomposition) of ammonia to relieve the stress.
II. Pressure / Volume changes on the Hydrogen-Nitrogen-Ammonia System N2(g) + 3H2(g) 2NH3(g) + 92 KJ
A B C
Note: The pressure is increased
N2(g) by decreasing the volume of
the container Concentration H2(g)
NH3(g)
A. The system is at equilibrium.
B. The pressure of the system is increased by immediately decreasing the container volume. All concentrations immediately increase, proportional to their mole ratios. The system tries to relieve the increase in pressure by making fewer molecules. This means a shift to the product side.
C. The system reaches a new equilibrium
Questions:
44. Does an increase in pressure always force an equilibrium to produce more products ? Explain.
45. What characteristic of the reaction determines whether the equilibrium will shift with an increase in pressure?
46. Why does an increase in pressure not shift an equilibrium system like
H2(g) + I2(g) 2HI(g)
III. Concentration Change Effects on the Nitrogen-Hydrogen-Ammonia System N2(g) + 3H2(g) 2NH3(g) + 92 kJ
A B C
N2(g) Concentration
H2(g)
NH3(g)
Reaction Progress
A. The system is at equilibrium
B. An amount of H2(g) is added to the system. The concentration of H2(g) immediately
increases. The system tries to use up the extra H2(g) by shifting to produce more product.
C. The system comes to a new equilibrium.
Question:
47. Nitric acid can be produced according to the following equation:
3NO2(g) + H2O(l) 2HNO3(aq) + NO(g)
Draw an equilibrium graph to show the effects on the equilibrium of increasing the [NO2(g)].
NO2 HNO3 NO
