chemistry_collision theory, until gibbs free theory

this is for the grade11 stem students

13. FACTORS AFFECTING
REACTION RATES
Chemistry Unit

15. 1.

16. 2.

20. Each group will investigate one factor
affecting reaction rates and create a
presentation explaining their assigned
factor with real-life examples.
Group 1: Nature of Reactants
Group 2: Concentration
Group 3: Temperature
Group 4: Surface Area

21. Your task is to analyze how the reaction rate is affected by
different factors and justify your answers
Scenario 1: Why does cutting potatoes into smaller pieces
help them cook faster?
Scenario 2: Why does a higher concentration of bleach clean
a stain faster?
Scenario 3: Why do engines heat up faster when running
compared to when they are idle?
Scenario 4: Why do iron nails rust faster in salty water than in
fresh water?
Scenario 5: Why does powdered sugar dissolve more quickly
in water than a sugar cube?
Discuss these scenarios and write down your answers.

23. Zero-Order Reaction: Rate = k
The reaction rate does not
depend on the reactant
concentration.
Even if you increase or
decrease [A], the reaction
proceeds at the same speed.

24. First-Order Reaction: Rate = k[A]
The reaction rate depends
directly on the concentration of
reactant A.
If [A] is doubled, the reaction
doubles in speed.

25. Second-Order Reaction: Rate =
k[A]²
The reaction rate depends on the
square of the reactant
concentration.
If [A] is doubled, the reaction rate
increases 4 times.

33. Thermodynamics
the study of energy changes
and flow of energy from one
system to another

34. ice melting
above 0°C
object falling
to earth

35. object rising
from earth
We can make an
object rise from the
earth by picking it
up.

36. Tell whether the following processes is
spontaneous or non-spontaneous.

37. Spontaneity in chemistry refers to whether a
process or reaction occurs on its own
without continuous external influence. A
spontaneous process happens naturally
under given conditions, while a non-
spontaneous process requires energy input
to proceed.

38. SPONTANEOUS
PROCESSES
a physical or chemical change that
occurs by itself.
- These processes occur without
requiring an outside force.

40. Temperature, pressure, and
energy changes are key factors
that determine whether a
process is spontaneousor not

41. A process is more likely to occur on its own if it
releases energy and requires less input. Higher
temperatures can help certain processes
happen, while lower temperatures can slow
them down. Increased pressure can affect
reactions, especially those involving gases, by
making them less likely to proceed naturally.

42. Enthalpy affects spontaneity by
determining whether a process releases
or absorbs energy. If a process releases
heat (exothermic, ΔH < 0), it is more likely
to occur naturally. If it absorbs heat
(endothermic, ΔH > 0), it may need
external energy to proceed.

43. 1. Define spontaneity and enthalpy.
2. Categorize given processes as spontaneous
or non-spontaneous.
3. Give one example of an exothermic and an
endothermic process.

44. Enthalpy is the heat content of a
system, and it tells us whether a
reaction releases or absorbs heat. If it
releases heat, it’s exothermic; if it
absorbs heat, it’s endothermic.

47. Entropy Changes (ΔS):
Positive ΔS: The system becomes more
disordered (e.g., solid melting into a liquid,
liquid evaporating into a gas).
Negative ΔS: The system becomes more
ordered (e.g., gas condensing into a liquid,
liquid freezing into a solid).

54. GIBBS FREE ENERGY

61. The decomposition of hydrogen peroxide (H₂O₂) into water and oxygen is given by:
2H2O2(aq)→2H2O(l)+O2(g)
Given the following data at 298 K:
ΔH° = -196 kJ
ΔS° = 125 J/K
Questions:
1. Convert ΔS° into kJ/K.
2. Calculate ΔG using the Gibbs Free Energy equation
3. Calculate ΔG using the Gibbs Free Energy equation: ΔG=ΔH−TΔS
4. Determine if the reaction is spontaneous or non-spontaneous at 298 K.
5. Predict what happens to ΔG if the temperature increases. Will the reaction become
less spontaneous? Explain.

62. CHEMICAL
EQUILIBRIUM

63. Chemical equilibrium
occurs when the rate of the forward reaction
equals the rate of the reverse reaction,
meaning the concentrations of reactants and
products remain constant over time.
However, this does not mean the reactions
stop—both the forward and reverse reactions
continue happening at the same rate

64. Key Features of Chemical Equilibrium:
✅ Dynamic Process – The reaction is still happening, but no
overall change is observed.
✅ Constant Concentrations – The amounts of reactants and
products stay the same.
✅ Reversible Reactions – Equilibrium only occurs in reactions
that can go both forward and backward.
✅ Depends on Conditions – Factors like temperature,
pressure, and concentration can shift equilibrium.

65. Homogeneous Equilibrium happens when all reactants
and products are in the same phase—either all gases or all
aqueous solutions.
N2(g)+3H2(g)⇌2NH3(g)
Heterogeneous Equilibrium happens when reactants and
products are in different phases (solid, liquid, or gas).
CaCO3(s)⇌CaO(s)+CO2(g)

66. The equilibrium constant (Kc) is a value that expresses the
ratio of product concentrations to reactant concentrations at
equilibrium, raised to their respective coefficients in the
balanced equation. It helps determine whether the
equilibrium favors the reactants or products.
General Formula for Kc:
For a reversible reaction:
aA+bB⇌cC+dD

67. The equilibrium constant expression is:
Kc=[C]^c[D]^d/[A]^a[B]^b
where:
[C] and [D] = concentrations of products at equilibrium
[A] and [B] = concentrations of reactants at equilibrium
a, b, c, d = coefficients from the balanced equation

68. EXAMPLE:
N2(g)+3H2(g)⇌2NH3(g)
The equilibrium constant expression is:
Kc=[NH3]^2/[N2][H2]^3

71. EQUILIBRIUM
CONSTANT &
QUOTIENT

73. If Kc < 1 - reactants are favored at equilibrium
If Kc > 1 -products are favored at equilibrium

74. Given the K values at 25°C, predict the extent of each reaction by
indicating whether it is
reactant-favored or product-favored.

_______________ a. N2(g) + 3H2(g) ⇌ 2NH3(g) K = 3.5 x 10^8
_______________ b. CaCO3(s) ⇌ Ca+2(aq) + CO3 –2 (aq) K = 3.8 x 10^–9

75. Example Calculation:
H₂ + I₂ ⇌ 2HI
[H₂] = 0.2M
[I₂] = 0.3M,
[HI] = 1.5M

76. The reaction quotient (Q) is a measure of the
relative amounts of products and reactants
during a reaction at a given time. The expression
for the reaction quotient, Q, has the same form
as the equilibrium constant, but it involves
specific values that are not necessarily
equilibrium concentrations

77. Reaction Quotient (Q) vs. Equilibrium Constant
(K)
Concept:
Q < K: Reaction shifts forward (more products).
Q = K: Reaction is at equilibrium.
Q > K: Reaction shifts backward (more reactants).

79. N2O4⇌2NO2
[N2O4]=0.50M
[NO2]=0.20M
If initially:
Find the reaction quotient (Q)

80. YOUR TASK:
1. Calculate Q using the given concentrations.
2. Compare Q and K to determine if the reaction
shifts left (toward reactants) or right (toward
products).
3. Explain what happens to the amounts of
reactants and products as the reaction progresses.

81. SO2(g)+NO2(g)⇌SO3(g)+NO(g)
At a certain temperature, the equilibrium constant is
K=5.0
The initial concentrations are:
[SO2]=0.4M,[NO2]=0.2M,[SO3]=0.1M,[NO]=0.3M
Write the expression for the reaction quotient (Q).
Calculate Q by substituting the given
concentrations into the expression.

82. Calculate the equilibrium constant and predict the
direction of a reaction using Q,
PCl5⇄PCl3+Cl2
The following concentrations are measured:
[PCl5]=0.8M, [PCl3]=0.4M, [Cl2]=0.2M
If the equilibrium constant is K=0.5, determine if the
reaction will shift left, shift right, or remain at
equilibrium.

83. LE CHATELIER’S PRINCIPLE
(CONCENTRATION AND PRESSURE)

84. The French chemist Henry-Louis Le Chatelier discovered
that when stress is applied to a system in equilibrium, its
tendency is to shift in a direction that best reduces the
stress so that equilibrium will be re-established. Le
Chatelier’s Principle emphasizes that disturbing an
equilibrium due to changes in condition may cause the
equilibrium to shift either to the right or to the left. Le
Chatelier’s Principle allows qualitative predictions about
the response of an equilibrium system to various stress or
disturbances.

85. Different Types of Stress that can be
Applied to a System.

86. Effect of Changing the Concentration
Consider the reaction:
N2(g) + 3H2(g) ⇌ 2NH3(g)

87. Stress
Direction of Shift
Increase concentration of reactant Forward
Increase concentration of product Reverse
Decrease concentration (or remove)
reactant
Reverse
Decrease concentration (or remove)
product
Forward

88. Effect of Changing the Pressure and
Volume
Consider the formation of ammonia:
N2(g) + 3H2(g) ⇌ 2NH3(g)
Number of moles: 1 3 2

89. Only gases are affected significantly by changes in
pressure and volume because gases are free to
expand and compress in accordance with Boyle’s
Law. An increase in pressure will result in a decrease
in volume and vice versa. However, liquids and
solids are not compressible, so their volumes are
unaffected by pressure.

90. Stress
Direction of Shift
Volume decrease, pressure increase Toward smaller number of moles of gas
Volume increase, pressure decrease Toward larger number of moles of gas

91. In contrast, the decomposition of hydrogen iodide,
2HI(g) ⇌ H2(g) + I2(g)
is unaffected by pressure. The number of moles of gaseous
product and reactant are identical. If the product and
reactant side contain the same number of moles of
gaseous substances, the addition and removal of pressure
will not affect the system

92. Consider the following equilibrium:
2SO2(g) + O2(g) ⇌ 2SO3(g)
How will each of the following changes in concentration affect an
equilibrium mixture of the three gases? Will the direction of the
equilibrium shift forward, reverse or no change?
__________ 1. O2 is added
__________ 2. SO2 is removed
__________ 3. Concentration of SO3 is decreased
__________ 4. Concentration of O2 is decreased
__________ 5. SO2 is added

93. Consider the following equilibrium:
2SO2(g) + O2(g) ⇌ 2SO3(g)
How will each of the following changes in pressure affect an
equilibrium mixture of the three gases? Will the direction of the
equilibrium shift forward, reverse or no change?
__________ 1. decreasing the total pressure
__________ 2. the volume of the reaction vessel is doubled
__________ 3. The total pressure of the system is increased by
adding a noble gas
__________ 4. The volume of the reaction vessel is decreased

94. Predict whether the equilibrium for the following
reactions would shift forward, reverse,
or no change when the pressure is increased.
________________ 1. 2CO(g) + O2(g) ⇌ 2CO2(g)
________________ 2. 2NO(g) ⇌ N2(g) + O2(g)
________________ 3. N2O4(g) ⇌ 2NO2(g)
________________ 4. Ni(s) + 4CO(g) ⇌ Ni(CO)4(g)
________________ 5. N2(g) + 3H2(g) ⇌ 2NH3

96. Cup Condition Initial Temperature (°C) Final Temperature (°C) Temperature Change (°C)
Cold (Ice Water Bath)      
Room Temperature      
Warm (Warm Water Bath)    

97. The reaction that proceeds and produces heat as a by-
product is an exothermic reaction. A reaction that requires
heat for the reaction to proceed is called an endothermic
reaction. The change in enthalpy △H of a reaction shows if the
system is endothermic or exothermic. A negative value means
the system is giving off heat, hence, exothermic; while a
positive value means it is using up heat and therefore,
endothermic
Effect of Changing the Temperature

98. Endothermic: Reactants + Heat ⇌ Products
Exothermic : Reactants ⇌ Products + Heat
Here is another way to distinguish endothermic and
exothermic reactions:
Endothermic: Reactants ⇌ Products
△H = positive (△H>0)

Exothermic : Reactants ⇌ Products
△H = negative (△H<0)

99. According to Le Chatelier’s Principle, increasing the temperature of an
equilibrium system would favor the reaction that will absorb the added
heat, which is the endothermic direction. Cooling the reaction or lowering
the temperature would favor the direction that produces heat, or the
exothermic reaction. A change in temperature causes a change in the value
of the equilibrium constant. Consider the equilibrium system:

N2(g) + 3H2(g) ⇌ 2NH3(g) + heat

100. The forward direction is exothermic while the reverse
reaction is endothermic. Using Le Chatelier’s Principle,
increasing the temperature would cause the equilibrium to
shift to the left (i.e. reverse), the endothermic direction.
Cooling the system would favor the exothermic direction,
the forward direction, thus forming more NH3 and
resulting in a K which is higher than the original.

101. Co(H2O)6 + 2(aq) + Cl–1 (aq) ⇌ CoCl4–2 (aq)) + 6H2O(l)
△H>0
Since △H>0 (positive), the forward reaction is
endothermic while the reverse reaction is exothermic.
When the solution is heated, the equilibrium shifts to
the right, the endothermic direction. Lowering the
temperature will favor the reverse direction, the
exothermic direction

102. The addition of a catalyst does not affect the
equilibrium state. It only hastens or delays the
rate at which the system achieves equilibrium.
The presence of a catalyst does not affect
the equilibrium concentration and constant
rather the time it takes to achieve equilibrium
Effect of Catalyst

103. 1. 2SO2(g) + O2(g) ⇌ 2SO3(g)
Ho = –198 kJ/mol
2. PCl5(g) ⇌ PCl3(g) + Cl2(g)
Ho = 92.5 kJ/mol
3. H2(g) + I2(g) ⇌ 2HI(g)
Ho = –9.45 kJ/mol
4. 2NOCl(g) ⇌ 2NO(g) + Cl2(g)
Ho = 75 kJ/mol
5. C(s) + H2O(g) ⇌ CO(g) + H2(g)
Ho = 131 kJ/mol
What would be the effect of increasing the temperature on each of the
following systems at equilibrium? Will the equilibrium shift forward or
reverse?

104. In this reaction:
CO2(g) + H2(g) + heat ⇌ CO (g) + H2O (g)

A. Is heat absorbed or released by the
forward reaction?
B. Is heat absorbed or released by the
reverse reaction?