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Chapter 8: Reaction Rates and Equilibrium
Problems: 8.1-8.2, 8.10-8.11, 8.23-8.25, 8.30-8.31, 8.34, 8.51-8.58, 8.63-8.64
8.1 SPONTANEOUS AND NONSPONTANEOUS PROCESSES
spontaneous process: takes place naturally, without any apparent cause or stimulus
– e.g. ice melts in your hand, a drop of food coloring will spread throughout a glass of water
nonspontaneous process: only takes place as a result of some cause or stimulus
– e.g. H2O(l) can be converted to H2(g) and O2(g) using electricity,
gasoline burns after being ignited
8.2 REACTION RATES
In a chemical reaction, reactants are converted to products: reactants → products
rate of a chemical reaction: the speed that reactants are converted to products
chemical kinetics: study of the factors that affect the rates of chemical reactions
8.3 MOLECULAR COLLISIONS
Why does a reaction occur?
– Molecules collide with one another
→ Bonds are broken, and
bonds are formed.
– If the bonds that form are
different than the bonds broken
→ A chemical reaction occurred!
CHEM 121: Chapter 8 page 1 of 14
Collision Theory: Molecules must collide to react.
The more often molecules collide
→ more likely the collision will result in bonds breaking and forming (chemical reaction)
Two ways to increase collision frequency
i. Increase the concentration
– more atoms/molecules
→ more collisions that could result in a chemical reaction
→ higher the reaction rate
ii. Increase the temperature
– atoms/molecules are moving faster
→ more collisions that could result in a chemical reaction
→ higher the reaction rate
Example: Which will have more collisions: a sample of N2 molecules at 25°C or a sample at
100°C? Explain.
Two requirements for a reaction to occur:
1. Collision Energy
– For a reaction to occur, molecules have to be moving quickly enough that they can break
and reform bonds when they collide.
– If they're moving too slowly, they merely bounce off each other.
– The minimum energy required for a reaction is called the activation energy (Ea).
(a) When reactants molecules have the activation energy (i.e. collide with the minimum
energy required) → a reaction occurs, and reactants are converted to products
(b) When reactants molecules lack the activation energy (i.e. do not collide with enough
energy) → no reaction occurs, and reactants remain unconverted
CHEM 121: Chapter 8 page 2 of 14
2. Collision Geometry
– Molecules must be in the correct orientation for a reaction to occur
– See the reaction between NO and O3 in the figures below.
correct collision geometry → reaction occurs incorrect collision geometry → no reaction
Note that for the reaction to occur, the N atom in NO must be oriented towards the O3 molecule
when they collide.
– If the O atom in NO is oriented towards the O3 molecule when they collide, no bonds can
form, and the molecules simply bounce off each other.
8.5 FACTORS THAT INFLUENCE REACTION RATES
1. Reactant Concentration
– The higher the concentration = more molecules present
→ more collisions that result in a reaction → higher the reaction rate
2. Temperature
– Increasing the temperature increases reaction rate in 2 ways
i. Collision frequency increases since the reactants are moving faster.
ii. Collision energy is greater, so more molecules have the required activation energy to
react.
3. Catalyst: substance that speeds up a chemical reaction
– generally increase the reaction rate by providing an alternative collision geometry
requirement that lowers the activation energy barrier
– is not consumed in the reaction, so can be used continually
CHEM 121: Chapter 8 page 3 of 14
Metallic Surface Catalysts
– A catalyst can also be a different physical state than the reactants.
– Most often, a solid catalyst is used to increase the rate of a gas-phase or liquid-phase reaction
(e.g. catalytic converters in cars catalyze the oxidation of CO and unburned hydrocarbons to
CO2 and H2O).
– For example, the hydrogenation of ethylene, C2H4(g) + H2(g) → C2H6(g),
occurs much faster on a metal catalyst.
(a) C2H4 and H2 are adsorbed onto the metal surface.
(b) H2 breaks up into H atoms, and one H atom bonds to C2H4 to form a metal-bonded C2H5.
(c) The second H atom bonds to the C2H5.
(d) The resulting C2H6 desorbs from the metal surface.
Example: Consider the reaction between NO and O3 on the previous page. Consider the possibility
of the NO molecule being bound to a metal surface, insuring that the nitrogen atom
always sticking up.
a. Using this metal surface would ______ increase decrease have no effect on
the rate of reaction.
b. Explain why.
CHEM 121: Chapter 8 page 4 of 14
8.4 ENERGY DIAGRAMS (or ENERGY PROFILES OF CHEMICAL REACTIONS)
We can show a reaction in terms of a Reaction Profile (or Energy Diagram), where Energy is
plotted on the y axis the Progress of the Reaction is shown on the x axis.
– Note: The lower the energy of the reaction, the more strongly the atoms/ions are bonded to
each other, and the more stable the compounds.
Consider the figure at the right:
The transition state is the highest-
energy peak on the curve.
The activated complex is the
arrangement of atoms at the
transition state when bonds
between reactants and products are
half-formed and half-broken.
The activation energy (Eact or Ea)
is the energy difference between
reactants and the transition state.
The difference in energy between
the reactants and products is called the heat of the reaction (usually indicated as ∆H or ∆Hrxn).
– This is the amount of energy released or absorbed when the reaction occurs.
Endothermic versus Exothermic Reactions
In an endothermic reaction, the reactants are lower in energy than the products.
– When the reaction occurs, it absorbs heat from the surroundings to occur, so the
surroundings (including us as the observers) experience a drop in temperature.
– Thus, when you monitor a reaction, and the reaction beaker or test tube becomes cold,
the reaction taking place is endothermic.
– We can show heat like a reactant in endothermic reactions.
– For example, ammonium chloride dissolving is an endothermic process:
NH4Cl(s) + heat → NH4Cl(aq)
In an exothermic reaction, the reactants are higher in energy than the products.
– When the reaction occurs, it releases heat to the surroundings, so the surroundings
(including us as the observers) experience a rise in temperature.
– Thus, when you monitor a reaction, and the reaction beaker or test tube becomes hot,
the reaction taking place is exothermic.
– We can show heat like a product in exothermic reactions.
– For example, sodium hydroxide dissolving is an exothermic process:
NaOH(s) → NaOH(aq) + heat
CHEM 121: Chapter 8 page 5 of 14
Example: Consider the following observations of chemical reactions to determine whether the
reaction described is endothermic or exothermic, then include heat as a reactant or
as a product in each chemical equation.
1. When hydrochloric acid, HCl(aq), reacts with aqueous sodium hydroxide, NaOH(aq),
to produce water and a salt, NaCl(aq), the reaction beaker gets hot.
a. The reaction described is ___________. endothermic exothermic
b. Include heat as a reactant or product in this reaction:
HCl(aq) + NaOH(aq) → H2O(l) + NaCl(aq)
2. When barium hydroxide, Ba(OH)2(aq), reacts with aluminum chloride, AlCl3(aq), the
reaction beaker gets so cold, any moisture around the beaker freezes to ice.
a. The reaction described is ___________. endothermic exothermic
b. Include heat as a reactant or product in this reaction:
3 Ba(OH)2(aq) + 2 AlCl3(aq) → 2 Al(OH)3(s) + 3 BaCl2(aq)
Endothermic Reaction Energy Profiles
Ex. 1: Consider the following reaction energy profile, and write a chemical equation to represent
the reaction shown in the reaction energy profile above, including heat as a reactant or
product.
Ex 2. In an endothermic
reaction, the energy of the
reactants is _______ than
the energy of the products.
< > =
Ex 3. If you are holding a beaker
in which an endothermic
reaction occurs, the
beaker will feel _________
following the reaction.
hotter colder the same
CHEM 121: Chapter 8 page 6 of 14
Exothermic Reaction Energy Profiles
Ex. 1: Consider the following reaction energy profile, and write a chemical equation to represent
the reaction shown in the reaction energy profile above, including heat as a reactant or
product.
Ex 2. In an exothermic reaction,
the energy of the reactants
is _______ than the
energy of the products.
< > =
Ex 3. If you are holding a beaker
in which an exothermic
reaction occurs, the beaker
will feel _________
following the reaction.
hotter colder the same
Example: Consider the reaction energy diagram below:
CHEM 121: Chapter 8 page 7 of 14
a. Write a chemical equation to represent the reaction shown in the energy profile above, including
heat as a reactant or product.
b. This reaction is _________. (Circle one) exothermic endothermic
c. The difference in energy between the reactants and products is ______.
(Circle one)
the activation energy (Eact) ∆H the catalyst
d. The activation energy (Eact) of the reaction with a catalyst is _______ compared to the reaction
without a catalyst. (Circle one)
< > =
e. The heat of reaction (∆H) with a catalyst is _______ compared ∆H for the reaction without a
catalyst. (Circle one)
< > =
8.6 CHEMICAL EQUILIBRIUM
A chemical reaction proceeds in the forward direction, with reactants → products.
When enough products are formed, the reaction can also proceed in the reverse direction, so more
reactants are then formed from products → the reaction is reversible.
reactants products
Given enough time, the reaction will reach equilibrium
equilibrium: state where forward and reverse reactions occur at same rate
At equilibrium, since the concentrations of reactants and products are not changing, it may
appear that everything has stopped, but this is not true!
CHEM 121: Chapter 8 page 8 of 14
Traffic Analogy: Consider two island cities connected by two one-way bridges:
– Cars are not allowed on Island #2 after 6pm, and the bridges are closed from 6pm to 6am.
→ When the bridges open at 6am, all of the cars are on Island #1.
Island #1
Island #2
Ex. 1: Provide two explanations for the number of cars on each island not changing.
Ex. 2 One day traffic on the two bridges is equal by noon, so as soon as a car gets off the bridge
at Island #2, another car on Island #2 goes onto the bridge to Island #1.
Question 1: Is the number of cars on Island #1 changing? Yes No
Question 2: Is the number of cars on Island #2 changing? Yes No
Question 3: If the number of cars on each island is not
changing, does that mean traffic has stopped? Yes No
Question 4: Do the number of cars on each island have to Yes No
be equal for them to not be changing?
Thus, in this example the islands have achieved a state of equilibrium,
– The rate of traffic to Island #2 = the rate of traffic to Island #1.
– The # of cars on each island are not changing with time, but they need not be equal to one
another.
Any chemical reaction in a closed vessel will eventually achieve chemical equilibrium—a
state in the concentrations of all reactants and products remain constant with time.
→ At equilibrium, the rates of the forward and reverse reactions are equal.
CHEM 121: Chapter 8 page 9 of 14
Example: Sulfuryl chloride decomposes as follows: SO2Cl2(g) SO2(g) + Cl2(g)
The figures above show closed systems of SO2Cl2, SO2, and Cl2 at 375K.
• Initially, only SO2Cl2 molecules are present.
• When heated, the SO2Cl2 decomposes, and all three molecules are present.
• Given enough time, the system achieves equilibrium.
Ex. 1: Using the figures above, indicate the number of SO2, Cl2, and SO2Cl2 molecules at
equilibrium at 375K.
_______ SO2Cl2 molecules _______ SO2 molecules _______ Cl2 molecules
Ex. 2 : Indicate the number of SO2, Cl2, and SO2Cl2 molecules present 15 minutes after the
equilibrium is initially achieved at the same temperature.
_______ SO2Cl2 molecules _______ SO2 molecules _______ Cl2 molecules
Similarly, at equilibrium,
rate of the forward reaction = rate of the reverse reaction
The concentrations of reactants and products remain constant since reactants are
converted to products at the same rate products are converted to reactants.
However, concentrations of reactants and products do not have to be equal!
CHEM 121: Chapter 8 page 10 of 14
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