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Why do chemical reactions occur?
Spontaneous process -- a process that, once it is initiated,
proceeds on its own without any external influence
Spontaneous vs. non-spontaneous processes
Examples:
Spontaneous process -- a process that, once it is initiated,
proceeds on its own without any external influence
• Rolling a ball downhill
• Combustion of methane
CH4(g) + 2 O2(g)
Non-spontaneous process -- a process that takes place
only in the presence of continuous external influence
CO2(g) + 2 H2O(g)
Note that spontaneous does not necessarily mean instantaneous
-- some spontaneous processes proceed very slowly
• Transition of carbon from diamond form to graphite form
C(diamond)
C(graphite)
This process is spontaneous, but very slow
Non-spontaneous process -- a process that takes place only
in the presence of continuous external influence
Processes that decrease the energy of a system tend to occur
spontaneously
Examples:
• Rolling a ball uphill
•
Examples:
Electrolysis of water
2 H 2O
electricity
Rolling a ball downhill
2 H2
+
O2
Note that the reverse of a spontaneous process is always
non-spontaneous
• Transition of carbon from graphite form to diamond form
C(graphite)
Energy change and spontaneous processes
C(diamond)
(gravitational potential
energy decreases)
Decompressing a spring
(elastic potential
energy decreases)
!height
Energy change and spontaneous processes
Processes that decrease the energy of a system tend to occur
spontaneously
Exothermic reactions (which release heat because the chemical
potential energy of the products is lower than that of the reactants)
tend to occur spontaneously
Example: Combustion of methane
CH4(g) + 2 O2(g)
Energy change and spontaneous processes
Processes that increase the energy of a system tend to be
non-spontaneous
Examples:
Rolling a ball uphill
!height
(gravitational potential
energy increases)
CO2(g) + 2 H2O(g) + 890 kJ
Compressing a spring
Bond energy of reactants > Bond energy of products
Heat is
released
Energy change and spontaneous processes
Processes that increase the energy of a system tend to be
non-spontaneous
Endothermic reactions (which absorb heat because the
chemical potential energy of the products is greater than that of
the reactants) tend to be non-spontaneous
(elastic potential
energy increases)
Energy change and spontaneous processes
• Exothermic reactions tend to occur spontaneously
• Endothermic reactions tend to be non-spontaneous
But some exothermic reactions are non-spontaneous,
and some endothermic reactions are spontaneous
Example: Crystallization of ammonium nitrate from aqueous solution
Example: Electrolysis of water
2 H2O (l) + 286 kJ
Heat is
absorbed
2 H2 (g)
+
O2 (g)
Bond energy of reactants < Bond energy of products
NH4+(aq) + NO3–(aq)
NH4NO3(s) + 25.7 kJ
Bond energy of reactants > Bond energy of products
Heat is
released
!H is negative (exothermic) -- but this reaction is not spontaneous
Energy change and spontaneous processes
• Exothermic reactions tend to occur spontaneously
• Endothermic reactions tend to be non-spontaneous
But some exothermic reactions are non-spontaneous,
and some endothermic reactions are spontaneous
Energy change and spontaneous processes
How can an endothermic process be spontaneous?
How can an exothermic process be non-spontaneous?
We have to consider another property in addition to the
change in energy associated with the reaction
Example: Reaction between barium hydroxide and ammonium chloride
Ba(OH)2 (s) + 2 NH4Cl (s) + 63.6 kJ
Heat is
absorbed
BaCl2 (s) + 2 NH3 (aq) + 2 H2 O (l)
Bond energy of reactants < Bond energy of products
!H is positive (endothermic) -- but this reaction is spontaneous
entropy (S): The amount of disorder in a system
(on a molecular level)
Entropy increases as the disorder (or randomness) of
the particles in a system increases
Examples of physical /chemical processes
that increase the entropy of a system
Examples of physical /chemical processes
that decrease the entropy of a system
Processes that increase entropy tend to be spontaneous
Processes that decrease entropy tend to be non-spontaneous
Phase changes
Phase changes
• solid to liquid
• liquid to gas
• gas to liquid
• liquid to solid
Dissolution of a solid
Crystallization of a solid
Reactions that increase the total number of gas molecules
Reactions that decrease the total number of gas molecules
2 SO3(g)
2 SO2(g) + O2(g)
N2(g) + 3 H2(g)
2 NH3(g)
Changes in energy and enthalpy
Spontaneous or non-spontaneous?
Whether or not a reaction occurs spontaneously depends
on both the change in energy and the change in entropy
At standard temperature and pressure (25 °C, 1 atm), the reaction
between carbon (graphite) and oxygen to form carbon dioxide has
the following values for enthalpy and entropy change:
Energy
C(s, graphite) + O2(g)
!H is negative (energy decreases): Favorable
!H is positive (energy increases): Unfavorable
Entropy
!S is positive (entropy increases): Favorable
!S is negative (entropy decreases): Unfavorable
CO2(g)
! H = –393.5 kJ/mol
!H is negative: Favorable
! S = +2.91 J/mol K
!S is positive: Favorable
Is this reaction spontaneous under these conditions?
Reaction is spontaneous
Spontaneous or non-spontaneous?
The following reaction is endothermic:
6 CO2 (g) + 6 H2O (g)
C6H12O6 (s)
+
6 O2 (g)
Spontaneous or non-spontaneous?
At standard temperature and pressure (25 °C, 1 atm), the
decomposition of calcium carbonate has the following values for
enthalpy and entropy change:
CaCO3 (s)
CaO (s)
+
CO2 (g)
Is this reaction spontaneous?
! H = +178.3 kJ/mol
!H is positive: Unfavorable
! H > 0 (endothermic) !H is positive: Unfavorable
! S = +160.5 J/mol K !S is positive: Favorable
! S < 0 (number of gas molecules decreases)
!S is negative: Unfavorable
Reaction is non-spontaneous
Is this reaction spontaneous under these conditions?
????
Free energy change
Spontaneous or non-spontaneous?
A quantity called free energy change (! G ) can be calculated
to determine whether a reaction is spontaneous or non-spontaneous
At standard temperature and pressure (25 °C, 1 atm), the
decomposition of calcium carbonate has the following values for
enthalpy and entropy change:
!G = !H – T ! S
Free energy
change
Enthalpy change
(heat of reaction)
Entropy change
Temperature (Kelvin)
CaCO3 (s)
CaO (s)
+
CO2 (g)
! H = +178.3 kJ/mol
! S = +160.5 J/mol K = +0.1605 kJ/mol K
T = 25°C + 273 = 298 K
When ! G is negative: Reaction is spontaneous
When ! G is positive: Reaction is non-spontaneous
Is this reaction spontaneous under these conditions?
Calculate free energy change:
!G = !H – T ! S
! G = +178.3 kJ/mol – (298 K)(+0.1605 kJ/mol K)
! G = +130.5 kJ/mol
Free energy change is temperature dependent
At a temperature of 2000 K and pressure of 1 atm, the
decomposition of calcium carbonate has the following values for
enthalpy and energy change:
CaCO3 (s)
CaO (s)
+
CO2 (g)
! H = +178.3 kJ/mol
! S = +160.5 J /mol K = +0.1605 kJ/mol K
T = 2000 K
Is this reaction spontaneous under these conditions?
Calculate free energy change:
!G = !H – T ! S
! G = +178.3 kJ/mol – (2000 K)(+0.1605 kJ/mol K)
! G = –142.7 kJ/mol
! G negative: spontaneous
! G positive: non-spontaneous
Reaction rates
Collision theory of reactions
Collision theory of reactions
A chemical reaction does not take place if:
A chemical reaction occurs when:
• Collisions between molecules have sufficient energy to
break the bonds in the reactants
• Collisions between molecules do not have sufficient
energy to break the bonds in the reactants
• Molecules are not properly aligned
• Molecules collide with the proper orientation
Insufficient energy
N2
O2
NO
NO
Wrong orientation
• Bonds between atoms in reactant molecules (N2 and O2) are broken
• New bonds are formed between atoms in product molecules (NO)
Bonds between atoms in reactant molecules (N2 and O2)
are not broken -- no new products are formed
Activation energy
Activation energy
The activation energy of a reaction is the minimum amount
of energy required for the reaction to occur upon proper
collision between reactant molecules
Both exothermic and endothermic reactions require activation
energy to initiate the reaction
-- often supplied as a spark or flame
After the activation energy is provided, an exothermic reaction
generates enough heat to keep the reaction going
In an endothermic reaction, heat must be continuously
supplied to keep the reaction going (even after the activation
energy is provided)
Exothermic reaction
Endothermic reaction
Rate of reaction
Rate of reaction
The rate of a chemical reaction is equal to:
• the speed at which reactants are consumed
• the speed at which products are produced
Example: Consider the following decomposition reaction:
2 N2O5
4 NO2 + O2
Initial N2O5 concentration = 0.01756 mol/L
Temperature = 45°C
Change in concentration
of reactant/product
Time (min)
O2 concentration (mol/L)
0.00
0.00000
Change in time
20.0
0.00412
Reaction rate =
Over the first 20 minutes of the reaction, the average rate is:
(0.00412 mol/L – 0.00000 mol/L)
(20.0 min - 0.00 min)
=
2.06 x 10-4 mol / L min
Effect of temperature on reaction rate
Reaction rate increases when temperature increases
Atoms/molecules move faster at higher temperatures
• Increase in frequency of collisions
• Increase in energy of collisions
Effect of concentration on reaction rate
Reaction rate increases when the concentration of
reactants increases
• increase in frequency of collisions
HEAT
T1 < T2
C1 < C2
Temperature = T1
Temperature = T2
A catalyst is a substance that increases the rate of a chemical
reaction
Non-catalyzed reaction
O
N
O
N
O
N
O
N
O
N
O
N
O
N
O
Reaction in the presence of a catalyst (platinum)
N
O
N
O
N
O
N
O
Pt
Concentration = C2
Effect of a catalyst on reaction rate
Catalysts
N
Concentration = C1
The presence of a catalyst increases the reaction rate
• The catalyst reduces the activation energy
• The catalyst is not consumed during the reaction
Questions
What is the effect on the rate of reaction?
(I) increases, (D) decreases, or (N) no change
• Increasing the temperature
• Removing some of the reactants
• Adding a catalyst
• Placing the reaction flask in ice
• Increasing the concentration of a reactant
Homework assignment
Chapter 7 Problems:
7.30, 7.31, 7.32, 7.33, 7.34, 7.37, 7.38, 7.40, 7.42, 7.43,
7.46, 7.47, 7.48, 7.50