SCH4U1
UNIT 4  Chemical Systems and Equilibrium
STSE Focus:
E1.1 analyse the optimal conditions for a specific chemical process related to the
principles of equilibrium that takes place in nature or is used in industry (e.g., the
production of sulfuric acid, electrolyte balance in the human body, sedimentation in
water systems)
Sample issue: The principle of dynamic equilibrium is used in industrial processes to
maximize the concentration of products and minimize leftover reactants. Industrial
chemists determine ideal pressure and temperature conditions, and proper catalysts, so
that fewer materials and less energy are used.
Sample questions: Why are low temperature conditions not used with exothermic
reactions? How do chemicals dissolved in human blood help maintain a blood pH level
between 7.2 and 7.4?
E1.2 assess the impact of chemical equilibrium processes on various biological,
biochemical, and technological systems (e.g., remediation in areas of heavy metal
contamination, development of gallstones, use of buffering in medications, use of
barium sulfate in medical diagnosis)
Sample issue: Heavy metals such as copper, lead, and zinc can accumulate to toxic levels
in the human body. A process called chelation, which causes a chemical reaction
involving an equilibrium shift, removes the metals from the body before permanent organ
damage occurs.
Sample questions: Why are headache tablets buffered? Why is barium sulfate safe to use
for X-rays of the digestive system even though barium ions are poisonous? How do
kidney stones form?
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SCH4U1
UNIT 4 – Chemical Systems In Equilibrium (Chapter 7)
*Opening Reflection Question: How does it feel to walk up the down escalator or run on a treadmill?
What remains constant?________________________
 What opposing processes occur at the same rate? _______________________________
*In both of these case, a state of _____________ _______________ has been achieved – a situation where
at least one property remains ___________ while opposing processes occur at the __________ rate.
7.1 Dynamic Equilibrium in Chemical Systems
 Many chemical reactions proceed in ________directions. (i.e., do not go to completion).
 These types of reactions are symbolized using ___________ arrows. Double arrows represent
________________ reactions or physical changes.
*Examples:
i) carbonic acid (a weak acid):
;
ii) the Haber process (production of ammonia)
 In a chemical system that achieves equilibrium, both the _____________reaction (i.e. the left-to-right
reaction) and the ________________reaction (i.e. the right-to-left reaction) proceed ________________
(this is what the term “dynamic” is meant to convey) and at the __________ rate.
*Four Conditions that Apply to all Equilibrium Systems:
1. The exact moment of equilibrium happens when the rate of the forward reaction __________ the rate of
the reverse reaction. If this is not the case, the system has _______achieved a state of equilibrium.
2. When a chemical system is in equilibrium, ________________ properties (those we can see, e.g. colour)
remain _______________. Furthermore, the ______________, _____________, and ____ of the system at
equilibrium all remain ______________.
3. Equilibrium can only be achieved in a ____________ system at a constant pressure. A closed system is a
one that may exchange ____________ but not ____________ with its surroundings.
4. Equilibrium can be approached from __________ direction (i.e. from the forward direction or from the
reverse direction). *What does this mean? *Consider the following equilibrium: N2O4(g) <===> 2NO2(g)
2
*Examples:
A. Physical Processes:
i) _____________ Equilibrium: H2O(l)  H2O(g)
ii) _______________ Equilibrium: I2(s)  I2(aq) ; CO2(g) CO2(aq) ; CaSO4(s)  Ca2+(aq) + SO42-(aq)
B. Chemical Processes
iii) _________________ Equilibrium: H2(g) + I2(g)  2HI(g)
iv) _________________Equilibrium:
N2O4(g)  2NO2(g)
;
2H2O(l)  2H2(g + O2(g)
;
CaCO3(s)  CaO(s) + CO2(g)
v) ________-_________ Equilibrium: HC2H3O2(aq) + H2O(l)  C2H3O2-(aq) + H3O+(aq)
*Percent Reaction at Chemical Equilibrium:
%Rxn = [actual product yield / theoretical yield] x 100%
Consider: H2(g) + I2(g)  2HI(g)
@ 448C (Starting data: See p. 431)
Equ’m [HI]
(mmol/L)
System
1
2
3
1.56
2.10
2.50
Max. Possible [HI]
(mmol/L)
Percent Reaction (%)
2.00
2.70
3.20
78.0
77.8
78.1
Reactions fall loosely into three categories:
1.
_______________ strongly favoured  (%Rxn << 1%) (i.e. Virtually no reaction.)
2.
_______________ strongly favoured  (%Rxn > 99%)
(i.e. Rxn is quantitative – virtually all limiting reactant consumed)
3.
__________________ conditions achieved  (1% < %Rxn < 99%)
(i) %Rxn < 50% = _____________ favoured
(ii) %Rxn >50% = ______________ favoured
*CALCULATING EQUILIBRIUM CONCENTRATIONS (USING AN “ICE” TABLE)
I = ______________ __________________ of reactant and products(before reaction)
C = _____________ ___ _________________ of reactants and products between the start and the point at
which equ’m is achieved
E = _____________ of reactants and products at ___________________
e.g. Consider:
H2(g) +
Initial [ ]
Final [ ]
F2(g)
1.00 M 1.00 M
?
0.24 M

2HF(g)
0M
?
H2(g)
+
F2(g)

2HF(g)
[ ]i (mol/L)
[ ] (mol/L)
[ ]eq
Therefore:
*Try pp. 437-438 #6,7 (top) ; #1, 3, 4, 5, 8, 9
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7.2 EQUILIBRIUM LAW IN CHEMICAL REACTIONS
The Equilibrium Constant, K
H2(g) + I2(g)  2HI(g)
e.g. Consider the following equilibrium system:
*Equilibrium concentrations (mol/L) @ 440C
Trial
[H2]
1
0.0222
2
0.0350
3
0.0150
4
0.0442
K = _________________
[I2]
0.0222
0.0450
0.0135
0.0442
[HI]
0.156
0.280
0.100
0.311
[HI]2/ [H2][I2]
*Equilibrium Law Expression:
aA + bB ↔ cC + dD
K=
Where:
 A, B, C, D are chemical entities in gas or aqueous phases.
 a, b, c, d = ______________ in balanced chemical equation
 K is the ________________ constant
*Variances in equilibrium ______________ do not affect the value of K. The equilibrium constant is
affected by changes in_____________________.
e.g. Write the Equilibrium Law Expressions for the following systems:
(i) N2(g) + 3H2(g) ↔ 2NH3(g)
(ii) 2NO(g) + O2(g) ↔ 2NO2(g)
K=
K=
Equilibrium Constant of the Reverse Reaction (K′)
K′ =
1
K
(i.e. K′ = the _________________ of K)
Heterogeneous Equilibrium
Def’n  equilibrium in which reactants and products are in ___________ than one phase
*The concentration of reactants and products that are pure ____________ or ____________ are constant
and ____ _______ appear in the equilibrium expression!
e.g.
*1 mol of NaHCO3(s) occupies a volume of 38.9 cm3
*2 mol of NaHCO3(s) occupies a volume of 77.8 cm3
 The ratio of moles to volume (i.e. the molar concentration) remains ____________.
e.g. Write the equilibrium law expression for the hydrolysis of H2O in a closed container.
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2H2O(l) <====> 2H2(g) + O2(g)
K=
[ H 2( g ) ]2 [O2( g ) ]
[ H 2 O(l ) ]
… becomes 
*The [H2O(l)] is incorporated into the value of the equilibrium constant, K.
*The Meaning of the Equilibrium Constant, Kc: (see. p. 448)
As chemists, we are interested in the products of our reactions, so in terms of Kc, we are interested in
increasing the value of Kc, and thus increasing the extent of the reaction.
K >> 1
K>1
K≈1
K<1
K << 1
Reaction goes to ____________________ (i.e. it is _____________________)
Equilibrium position favours _________________
Reactant and product concentrations are __________________
React position favours __________________
Reaction essentially does ___________ occur
*Try: pp. 442-445 #1-7 ; *More Practice pp. 448-449 #1-9
7.3 Qualitative Changes in Equilibrium Systems
http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/lechv17.swf
*Le Chatelier's Principle: If a "____________" is applied to a system at equilibrium, the equilibrium
condition is _________; a net reaction (“shift”) occurs in that direction which tends to ___________ the
"stress," and a ___________ equilibrium is obtained.
initial equilibrium state  “non-equilibrium” state     new equilibrium
--------------------------------------------------------------------------------------------------------------------------------“shift” to right = ______________ reaction predominates until equilibrium is re-established (system
counteracts the “stress” by increasing the concentration of _____________)
"Stress"
" Shift" to left or right
“shift” to left = _____________ reaction predominates until equilibrium is re-established (system
counteracts the “stress” by increasing the concentration of ____________)
*Variables Affecting Chemical Equilibria
Variable
Concentration
Type of Change
Increase
Decrease
Temperature
Increase
Decrease
Increase
(↓ in pressure)
Decrease
(↑ in pressure)
Variables that do not affect chemical equilibria:
 the presence of ________________
 the presence of ________ gases (e.g.
Volume
Response to System
Shifts to ____________ some of the added reactant
or product
Shifts to ____________ some of the removed
reactant or product
Shifts to ___________ some of the added
__________energy (i.e. shifts in the
_______________direction)
Shifts to _____________ some of the removed
thermal energy (i.e. shifts in the _____________
direction)
Shifts toward the side with the ____________ total
amount of gaseous entities
Shifts toward the side with the ____________ total
amount of gaseous entities
*Practice Activity  Le Chatelier’s Principle
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1. Concentration Effects
*Increasing the concentration of a reactant causes a shift to the __________.
*Decreasing the concentration of a reactant causes a shift to the _________.
*Increasing the concentration of a product causes a shift to the __________.
*Decreasing the concentration of a product causes a shift to the __________.
*The shift continues until a______ equilibrium is established.
 In which direction does the equilibrium shift as a result of the change to each homogeneous equilibrium
system?
(a) 2Cl2(g) + O2(g) <===> 2Cl2O(g)
*Adding Cl2(g)
= shift to the ____________
*Removing Cl2O(g) = shift to the ____________
*Adding Cl2O(g)
= shift to the ____________
(b) 2NO2(g) <===> N2(g) + 2O2(g)
*Removing N2(g) = shift to the ___________
*Removing NO2(g) = shift to the ___________
*Adding O2(g)
= shift to the ___________
2. Temperature Effects
*Increasing the temperature of an equilibrium system causes a shift in the ____________
direction.
*Decreasing the temperature of an equilibrium system causes a shift in the ___________
direction.
 For each of the following reversible reactions, determine whether the forward reaction is favoured by
high temperatures or low temperatures.
(a) N2O4(g) <===> 2NO2(g)
∆H = +59 kJ
or: N2O4(g) + 59 kJ <===> 2NO2(g)
*Forward = ____________, therefore ________ temperature will favour forward reaction.
(b) 2ICl(g) <===> I2(g) + Cl2(g)
or:
∆H = -35 kJ
*Forward = ___________, therefore _________ temperature will favour forward reaction.
3. Volume/Pressure Effects
*Decreasing the volume (______________ the pressure) of an equilibrium system favours the reaction that
produces the __________ total amount of gaseous entities.
*Increasing the volume (______________ the pressure) of an equilibrium system favours the reaction that
produces the __________ total amount of gaseous entities.
 In each of the following equilibrium systems, the volume of the reaction vessel is decreased (pressure is
_________). What is the effect (if any) of the position of equilibrium? (i.e. What is the direction of the
equ’m shift?)
(a) 4HCl(g) + O2(g) <===> 2Cl2(g) + 2H2O(g)
*Gaseous entities left side = ______
*Gaseous entities right side = ______
*Therefore, ↑ in pressure = shift to ___________.
(b) 2H2S(g) + CH4(g) <===> 4H2(g) + CS2(g)
*Gaseous entities left side = ______
*Gaseous entities right side = ______
*Therefore, ↑ in pressure = shift to ___________.
(c) N2(g) + O2(g) <===> 2NO(g)
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*Gaseous entities left side = ______
*Gaseous entities right side = ______
*Therefore, ↑ in pressure = ____ shift.
Catalysts
*Catalysts ___ _____ affect the final position of equilibria.
*Catalysts do, however, ____________ the time required to reach the equilibrium position. That is,
equilibrium is established much more rapidly in the presence of a catalyst.
*The forward and reverse reaction rates are increased _____________.
Addition of Inert Gases
*The addition of noble gases, for example, does not affect the equilibrium position.
Homework:
pp. 457-458 #1-8
*More Practice  p. 459 #1-6
7.5 Quantitative Changes in Equilibrium
*The Reaction Quotient, Q
 The reaction quotient, Q, is calculated in order to determine if a system has established _____________
equilibrium.
 Calculate Q exactly the same way you would calculate the equilibrium constant, K.
*The Three Possibilities:
1. Q = K ; *In this case, the system has established _________________.
2. Q < K ; *System needs to shift to the _________ in order to establish equilibrium.
3. Q > K; *System needs to shift to the ________ in order to establish equilibrium.
*pp. 465-480 #1-10 ; More Practice p. 481 #1-8
7.6 Solution Equilibrium (Heterogeneous Equ’m) and the Solubility Product Constant, Ksp
*Knowing the concentration of ions in aqueous solution is important in medicine and chemical analysis.
Solution Equilibrium
*When crystals are first placed in a solvent, many particles may leave the surface of the crystal and go
into solution. As the number of solute particles in the solution increases, some of the dissolved particles
return to the surface of the crystal. When the solution is ______________, the number of particles
leaving the crystal surface ___________ the number returning to the surface [i.e. the rate of
______________ (the process of dissolving) equals the rate of _______________]. At this instance,
___________ equilibrium has been established.
*Recall:
A _____________ solution is not capable of dissolving more ions.
 The ______________of a substance is the maximum amount of solute that will dissolve in a given
quantity of solvent at saturation. (e.g. Solubility of NaCl = _______ g per 100mL of water at 20C)
The Solubility Product Constant, Ksp
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*Ksp values are typically listed for ionic compounds with ______ solubility only. The solubility of highly
soluble substances are usually given in mol/L or g/100 mL.
Example 1: Consider a heterogeneous equilibrium between sparingly soluble copper(I) chloride, CuCl(s),
and its dissolved Cu+(aq) and Cl-(aq) ions.
CuCl(s)  Cu+(aq) + Cl-(aq)
Ksp = [
(aq)][
(saturated solution)
(aq)]
Ksp = 1.7 x 10-7 at 25C (From data table – p. 802.)
Ex. 2: Write the Ksp expression for an equilibrium system between calcium phosphate and its dissolved
ions.
______________________________________________________________________________________
*Solubility Equilibrium Calculations (4 Types):
Type 1: Calculating the Ksp from Solubility Data:
Example: Calculate the solubility product constant for lead(II) chloride, if 50.0 mL of a saturated solution
of lead(II) chloride was found to contain 0.2207 g of lead(II) chloride dissolved in it.
*First, write the equation for the dissolving of lead(II) chloride; then, write the equilibrium law expression
for the dissolving process. (i.e. the Ksp expression)
*Secondly, convert the amount of dissolved lead(II) chloride into moles per liter.
*Third, construct an “ICE” table.
<=========>
+
I
C
E
*Fourth, substitute the equilibrium concentrations into the equilibrium expression and solve for the Ksp.
Type 2: Calculating the Solubility of an Ionic Compound in Pure Water from its K sp
8
Example: Estimate the solubility of silver chromate, Ag2CrO4, in pure water if the solubility product
constant for silver chromate is 1.1 x 10-12.
*First, write the equation and the equilibrium (Ksp) expression.
*Next, construct an “ICE” table.
 Let "x" be the number of moles of silver chromate that dissolves in every liter of solution (its solubility).
<=========>
+
I
C
E
*Lastly, substitute the equilibrium amounts and the Ksp into the equilibrium expression and solve for x.
Type 3: Calculations Involving The Common Ion Effect:
*Theory The Common Ion Effect (An Application of Le Chatelier’s Principle)
Suppose we establish the equilibrium:
CaCO3(s)  Ca2+(aq) + CO32-(aq)
*Adding Ca2+ by dissolving CaCl2 will cause the equilibrium to shift to the _________.
*This will cause some CaCO3 to ____________ and less CaCO3 will be left in solution when equilibrium is
finally re-established.
*There are two sources of Ca2+ :
1. the ___________ that was added; and,
2. the ___________ that remains in solution when equilibrium is reestablished.
*Since Ca2+ is common to both salts, it is called the __________ ____.
*The addition of the common ion ___________ the solubility of the CaCO3.
*This lowering of the solubility by the addition of a common ion is called the common ion effect.
Ex.: What is the molar solubility of PbCl2 in a 0.10 mol/L NaCl solution? For PbCl2, Ksp = 1.2 x 10-5.
The equilibrium is: PbCl2(s)  Pb2+(aq) + 2Cl-(aq) ; Ksp =
*Common Ion = _____
*Firstly, determine the initial concentration of the common ion, Cl-. (The NaCl dissolves completely.)
NaCl(aq)
%
100



Na+(aq)
+
Cl-(aq)
I
C
“E”
*ICE table for the actual equilibrium:
9
PbCl2(s)
<=========>
Pb2+(aq)
+
2Cl-(aq)
I
C
E
*Calculate the solubility of PbCl2 in 0.10 M NaCl solution.
[Pb2+][Cl-]2 = Ksp
*What is the solubility of lead(II) chloride in pure water?
Summary:
Solubility of PbCl2 in 0.10 M NaCl solution = ______________ M;
Solubility of PbCl2 in pure water = _______________ M
Type 4: Double Displacement Reactions and Predicting Precipitation
Overview:
- calculate the trial ion product, Q (similar to reaction quotient)
Three Possibilities:
1) Q = Ksp  _____ precipitate formed (saturated sol’n)
2) Q < Ksp  _____ precipitate formed (unsaturated)
3) Q > Ksp  precipitate formed (supersaturated)
Example:
*25.0 mL of 0.010 mol/L silver nitrate is mixed with 50.0 mL of 0.0050 mol/L potassium chloride. Predict
whether a precipitate forms.
STEPS:
1. a) Overall equation? b) Net-ionic equation? c) equilibrium equation? d) Ksp expression? e) Q expression?
2. Use dilution equation…to determine the final concentrations of each species (ions of potential
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precipitate) once mixed. C2 =
C1V1
V2
3. Calculate the reaction quotient, Q, and compare to the value of the equilibrium constant, Ksp. Will a
precipitate form?
Try: pp. 486-492 #1-11 ; More Practice: p. 493 #1-14
7.7 Thermodynamics and Equilibrium
*Spontaneous (favourable) changes  one that, given the necessary _______________ energy, proceeds
_______________ continuous outside assistance
*What Conditions are Conducive to Favourable Changes?
1. Enthalpy and Favourable Changes:
 When a change __________ the energy of a system, events tend to occur spontaneously. In chemical
reactions, _______________ reactions tend to be spontaneous.
*However, not every exothermic reaction is spontaneous and not every endothermic reaction is nonspontaneous. The important point is that an energy decrease works as one factor in favour of spontaneity.
2. Entropy and Favourable Changes:
*Entropy, S,  a measure of the _____________ or _______________ of a system
 An ordered arrangement of particles (atoms, ions or molecules) has __________ entropy (______
disorder) than the same number of particles in random arrangements. The greater the _________ of
movement, the larger the entropy.
[ΔS = Sproducts - Sreactants] *Any event that is accompanied by an ___________ in the entropy of the system
tend to occur spontaneously. (*Second Law of Thermodynamics = The entropy of the universe is
constantly increasing.)
*Entropy and Physical Changes:
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In general, a system will experience an increase in entropy (S > 0) if:
a) the ____________ of a gaseous system increases,
b) the ___________________ of a system increases,
c) the _____________ state of a system changes from solid to liquid or gas, or liquid to gas
(Sgas > Sliquid > Ssolid); or,
d) A solute dissolves in a solvent.
*Determine whether the following reactions/changes show an increase or decrease in entropy. (Hint: 
disorder = +ve S ;  disorder = -ve S)
1. 2KClO3(s)  2KCl(s) + 3O2(g)
2. H2O(l)  H2O(s)
3. N2(g) + 3H2(g)  2NH3(g)
4. NaCl(s)  Na+(aq) + Cl-(aq)
5. KCl(s)  KCl(l)
6. CO2(s)  CO2(g)
7. Ag+(aq) + Cl-(aq)  AgCl(s)
_________
_________
_________
_________
_________
_________
_________
Gibbs Free Energy  the amount of energy available to do work
*For a reaction to be spontaneous, the sign of G has to be __________. To calculate Gibbs Free Energy
use: G =H - TS
G = change in free energy; H = enthalpy change; T = temperature in Kelvin ; S = entropy change
 When G is negative, the ____________ reaction is spontaneous
 When G is positive, the ____________ reaction is spontanous
 When G is zero, the reaction has established ______________.
Complete the following summary table:
Spontaneity
H
S
G
+
Always _________ _________ spontaneous
+
Always _________ __________ spontaneous
Depends on T
Spontaneous at ___ temperatures, when TS<H
+
+
Depends on T
Spontaneous at _____ temperatures, when TS>H
*Remember  [negative, negative, low, less] and [positive, positive, high, greater]
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