Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Unit 9b:
Equilibrium, Enthalpy, and
Entropy
1.
Student Name: _______________________________________
Class Period: ________
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Unit 9b: Equilibrium Systems
Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
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Unit 9b: Equilibrium Systems
Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Unit 9b Vocabulary:
1. Activated Complex: The species that are formed and decomposed
during the mechanism, and is also called the intermediate.
2. Activation Energy: The energy that must be added to allow the
reactants to complete the reaction and form the activated complex.
3. Catalyst: A chemical that is added to a reaction to eliminate steps in
the mechanism and increase the reaction rate and decrease the
activation energy without itself being consumed by the reaction.
4. Effective Collision: A collision between reactant particles that results
in a chemical reaction taking place.
5. Enthalpy: The total amount of potential energy stored in a substance.
6. Endothermic: A reaction that absorbs and stores energy from the
surrounding environment.
7. Entropy: A system’s state of disorder. Entropy increases as
temperature increases. Entropy increases as a substance goes from
solid to liquid to gas.
8. Equilibrium: A system where the rate of forward change is equal to
the rate of reverse change. At equilibrium there is no net change.
9. Exothermic: A reaction that releases stored energy into the
surrounding environment.
10. Favored: A change in a thermodynamic property that contributes
towards the reaction being spontaneous.
11. Free Energy: The total amount of energy available in a system to do
work. Free Energy is a combination of both enthalpy and entropy.
12. Heat of Reaction: The net gain or loss of potential energy during a
chemical reaction.
13. Inhibitor: A chemical that is added to a reaction to add steps to the
mechanism to decrease the reaction rate and increase the activation
energy without itself being consumed by the reaction.
14. Kinetics: The study of reaction mechanisms and reaction rates.
15. Nonspontaneous: A reaction that requires a constant input of energy
to occur, or the reaction will reverse or stop.
16. Reaction Rate: The amount of reactant consumed in a given unit of
time.
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Unit 9b: Equilibrium Systems
Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
17. Spontaneous: A reaction that continues independently once started.
18. Thermodynamics: The study of heat flow during physical and
chemical changes.
19. Unfavored: A change in a thermodynamic property that contributes
towards the reaction being nonspontaneous.
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Unit 9b: Equilibrium Systems
Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Notes page:
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Unit 9b: Equilibrium Systems
Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Unit 9b Homework Assignments:
Assignment:
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Date:
Due:
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Unit 9b: Equilibrium Systems
Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Topic:
Regents Chemistry ’14-‘15
Mr. Murdoch
Equilibrium
Objective: What is the role equilibrium has in chemistry?
Equilibrium:
Equilibrium is a continuous state of the rate of balance between two
opposing changes. In a state of equilibrium the rate of the forward
change is equal to the rate of the reverse change.
Most chemical reactions are reversible:
A + B  C + D + energy
C + D + energy  A + B
A + B  (±energy)  C + D
= forward reaction When the rate of the
= reverse reaction forward reaction equals
Double arrows () the rate of the reverse
indicate that BOTH
reaction, a state of
reactions are occurring
equilibrium is reached.
at the same time.
If you ride up a moving escalator, you are moving at the rate that the
escalator is moving upwards. However, if you turn around and start to
walk DOWN the up escalator, and you match the escalator’s rate (up)
but in the opposite direction (down), to someone watching you it looks
as if you are not moving. However, you are still expending energy
trying to go to the bottom, and the escalator is expending energy trying
to carry you uphill. If anything was to upset the process (power failure
to the escalator; you trip and fall, etc.), the equilibrium would be upset
and you would either make it to the bottom or ride to the top.
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Unit 9b: Equilibrium Systems
Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Equilibrium examples:
Haber Process for ammonia gas:
Forward reaction:
N2(g) + 3 H2(g)  2 NH3(g) + 92 kJ (exo)
Reverse reaction:
2 NH3(g) + 92 kJ  N2(g) + 3 H2(g) (endo)
Equilibrium:
N2(g) + 3 H2(g)  2 NH3(g) + 92 kJ
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Unit 9b: Equilibrium Systems
Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Topic:
Regents Chemistry ’14-‘15
Mr. Murdoch
Equilibrium Properties
Objective: What properties of equilibrium will systems have?
Properties of Systems at Equilibrium:
1. Equilibrium is a dynamic state; think of equilibrium as a continuous
pathway, never achieving a ‘set’ endpoint. Particles of reactants are
reacting and forming products at the same rate that products are
decomposing back into the reactants they came from. Remember that
the system is in continuous motion, though it may look like the
reaction is stagnant.
2. Equilibrium can only be maintained in a closed system. A closed
system neither gains nor loses anything. This includes energy (loss or
gain), adding reactants, or the removal of products.
3. As long as the system is closed, a system at equilibrium will remain
that way forever. Changing ANY condition of equilibrium will alter
the balance of the entire equilibrium (see pgs. 33-40).
4. Equilibrium occurs at different concentrations of product and
reactant. Depending on the nature of the species involved, assuming
we start with the forward reaction, the rate of the reverse reaction will
increase as the product is formed during the forward reaction. When
the forward AND reverse reaction rates are equal, equilibrium is
achieved. This may occur at different concentrations of product and
reactant.
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Unit 9b: Equilibrium Systems
Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Equilibrium Diagram:
 Equilibrium may be reached ANYWHERE along a line that starts at
0% and ends at 100%.
 At any point along the line the percentage of the reaction going
forward (reactants) ADDED to the percentage of the reaction going
backward (products) equals 100%.
(% forward reaction) + (% reverse reaction) = 100%
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Unit 9b: Equilibrium Systems
Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Topic: Three Types of Equilibrium
Objective: What forms of equilibrium are possible in chemistry?
1. Chemical Equilibrium:
i. If the rate of the forward reaction is equal to the rate of the reverse
reaction the reaction has achieved chemical equilibrium.
You have seen the Haber Process for the production of ammonia:
N2(g) + 3 H2(g)  2 NH3(g) + 92 kJ
This reaction produces ammonia (and heat), but some of the
ammonia, NH3(g), produced will decompose during the reaction back
into reactants, N2(g) and 3 H2(g).
ii. When the rate of synthesis (forward reaction) equals the rate of
decomposition (reverse reaction), and no other changes occur, this
system will be at equilibrium.
iii. As stated before, changing ANY component of the system will
change the equilibrium of the entire system.
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Unit 9b: Equilibrium Systems
Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
2. Solution Equilibrium:
i. If a solution becomes saturated, the rate of dissolving equals the rate
of precipitation, and the reaction has achieved solution equilibrium.
NaCl(s)  Na+1(aq) + Cl-1(aq)
ii. When sodium chloride is first placed into pure water, the solid ionic
crystals dissolve. As the concentration of the dissolved ions
increases, some of those dissolved Na+1(aq) and Cl-1(aq) ions will
temporarily rejoin to form a soluble precipitate which almost
immediately dissolves again. Eventually all the ions will be held
apart by the polar water molecules, and no more solid may enter the
solution until some ions come out of solution as precipitate. At this
point the rate of dissolving equals the rate of precipitation, and you
have a SATURATED solution. Additional added solid would not
dissolve, or only as a temporary supersaturated solution.
Solution Equilibriums
Unsaturated - solute almost all Close to Saturation - solute
undissolved; reaction almost all almost all dissolved; reaction
forward
mostly forward, some reverse
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Saturated-dissolving rate is
equal to precipitate formation
rate; no net change
Unit 9b: Equilibrium Systems
Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
3. Physical Equilibrium:
i. If the rate of a forward phase change is equal to the rate of a reverse
phase change, then the system is in Physical (or Phase) Equilibrium.
ii. Physical equilibrium occurs AT the phase change temperature.
Remember that during a phase change, all energy input is going
towards increasing the potential energy of the substance, as there is
no increase in average kinetic energy (temperature) at the phase
change temperature. For water, the boiling (vaporization point) at 1
atm is 373 K. This means if water is maintained in a sealed
container at 1 atm and 373 K, for each water molecule that changes
from liquid to gaseous, another water molecule will change from
gaseous to liquid.
Liquid-Gaseous Equilibrium for Water at 1 atm and 373 K
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Unit 9b: Equilibrium Systems
Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Triple Point Temperature:
An interesting phenomenon of Phase Equilibrium is the concept of the
Triple Point temperature. At a substance’s Triple Point, that substance
can exist in THREE phases simultaneously. Note that the Triple Point is
also a function of pressure.
For the diagram above, note that water is different from most materials
in that as you increase the pressure at the melting-freezing point, the
freezing point of water decreases with increasing pressure. This is one
of the reasons why a sealed can or bottle of a carbonated beverage can
have the liquid contents BELOW normal freezing temperature of water,
but as soon as the container is opened to the atmosphere (pressure drops
rapidly), the liquid may suddenly partially freeze, due to the fact that the
pressure drops much more rapidly than the temperature can increase.
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Unit 9b: Equilibrium Systems
Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Equilibrium Systems Practice Regents Problems: (ungraded)
1. Which statement about a system at equilibrium is true?
a) The forward reaction rate is less than the reverse reaction rate.
b) The forward reaction rate is greater than the reverse reaction rate.
c) The forward reaction rate is equal to the reverse reaction rate.
d) The forward reactions stop and the reverse reactions continue.
2. Given the reaction in water of AgCl(s)  Ag+1(aq) + Cl-1(aq), once equilibrium
is reached, which statement is accurate?
a) The AgCl(s) will be completely consumed.
b) The rates of the forward and reverse reactions are equal.
c) The entropy of the forward reaction will continue to decrease.
d) The concentration of Ag+1(aq) is greater than the concentration of Cl-1(aq).
3. Which type(s) of change, if any, may reach equilibrium?
a) A physical change, only.
b) A chemical change, only.
c) Neither a chemical nor a physical change.
d) Both a chemical and a physical change.
4. In a reversible reaction, a chemical equilibrium is attained when the
a) Concentration of the reactants reaches zero.
b) Concentration of the products remains constant.
c) Rate of the forward reaction is greater than the rate of the reverse reaction.
d) Rate of the reverse reaction is greater than the rate of the forward reaction.
5. Given the reaction of H2O(s)  H2O(l), at which temperature will equilibrium
exist when the atmospheric pressure is equal to 101.3 kPa?
a) 0 K
c) 273 K
b) 100 K
d) 373 K
6. The temperature at which solid and liquid phases of the same type of matter
exist in equilibrium is called its
a) Boiling point
c) Melting point
b) Heat of fusion
d) Heat of vaporization
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Unit 9b: Equilibrium Systems
Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Notes page:
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Unit 9b: Equilibrium Systems
Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Student name: _________________________
Class Period: _______
Please carefully remove this page from your packet to hand in.
Equilibrium Systems homework
1. Which of the following need to be equal at equilibrium?
a) The masses of both products and reactants.
b) The volumes of both products and reactants.
c) The concentrations of both the products and reactants.
d) The rates of formation of both the products and the reactants.
2. A stoppered (sealed) flask contains 20.0 grams of liquid water and 20.0 grams
of water vapor. Does a state of equilibrium of water exist in the flask?
a) Yes, the bottle is stoppered.
b) Yes, the amount of each component is equal.
c) Yes, but only if the rates of evaporation and precipitation are equal.
d) Yes, but only if a) and b) are both true.
e) Yes, but only if a) and c) are both true.
f) Yes, but only if a), b), and c) are all true.
3. A mixture of 50.0 g of water ice and 100.0 g of liquid water is massed and then
kept at a steady 0.0˚C in a closed container. After one hour, you mass the
contents of the container again. What would you predict the resulting masses to
be?
a) The mass of the ice and the mass of the liquid will remain constant.
b) The mass of the ice decreased and the mass of the water increased as the ice
melted.
c) The mass of the ice increased and the mass of the water decreased as the
water froze.
4. The same mixture in question #3 above of 50.0 g of water ice and 100.0 g of
liquid water is massed and then heated to 3.98˚C, the point of maximum density
for water. After one hour, you mass the contents of the container again. What
would you predict the resulting masses to be?
a) The mass of the ice and the mass of the liquid will remain constant.
b) The mass of the ice decreased and the mass of the water increased as the ice
melted.
c) The mass of the ice increased and the mass of the water decreased as the
water froze.
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Unit 9b: Equilibrium Systems
Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Use the information below to answer questions #5 and #6.
When sodium chloride is first added to distilled water, it looks like the solid
crystals disappear into the water. Continue to add crystals, and the rate that
they disappear slows down until eventually you achieve a condition where any
added solid sinks to the bottom of the container.
5. When the added sodium chloride crystals no longer dissolves in the water and
sinks to the bottom of the container, what type of a solution do you have?
a) Moist
b) Saturated
c) Unsaturated
d) Supersaturated
6. In the same conditions as listed above, while the solid crystals sit on the bottom
of the container, what will happen to the crystals in the water?
a) They’ll stay exactly the same.
b) They will change size and only get larger over time.
c) They will change size and only get smaller over time.
d) They will change size, but the total mass of the solids will be constant.
7. The vapor pressure of a liquid at a given average kinetic energy in a sealed
system is measured when the rate of evaporation of the liquid is
a) Less than the rate of condensation.
b) Equal to the rate of condensation.
c) Equal to a zero rate of condensation.
d) Greater than the rate of condensation.
8. A system is said to be in a state of dynamic equilibrium when the
a) Concentration of products is the same as the concentration of reactants.
b) Concentration of products is greater than the concentration of reactants.
c) Rate at which products are formed is the same as the rate at which reactants
are formed.
d) Rate at which products are formed is greater than the rate at which reactants
are formed.
9. A solution that is at equilibrium must be
a) Dilute
b) Saturated
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c) Unsaturated
d) Concentrated
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Unit 9b: Equilibrium Systems
Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Topic:
Regents Chemistry ’14-‘15
Mr. Murdoch
Reaction Progress
Objective: How do we know if a reaction will continue when started?
Spontaneous Chemical Reactions:
 A reaction that continues without additional input once it has initiated
is called a Spontaneous Reaction.
1. Spontaneous reactions are very useful in the industrial world, and
also in chemistry.
2. Once you start a spontaneous reaction it will go to completion until
either the reactants are consumed or it enters a state of equilibrium if
the products are not removed.
3. Spontaneous reactions require a balance between two factors:
i. Enthalpy- the available (potential) energy in a substance
ii. Entropy-the amount of randomness (disorder) in a system
Expanding
Universe!
Falling is
easier!
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Unit 9b: Equilibrium Systems
Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Topic:
Regents Chemistry ’14-‘15
Mr. Murdoch
Enthalpy
Objective: How do we express the potential energy in a system?
Enthalpy:
i. Enthalpy is the heat content (potential energy) of a system.
ii. Nature favors reactions that undergo a decrease in enthalpy.
iii. Let go of a ball in your hand; it falls.
iv. Falling is a spontaneous (no additional energy) decrease in enthalpy
(potential energy). The ball can’t fall again from its starting height.
 A decrease in potential energy is favored in nature, so exothermic
reactions are the most favored (and most common - see Table I).
v. Most exothermic (decreasing enthalpy) reactions are spontaneous,
and complete once started. (Think of a bonfire; it burns as long as it
has fuel.)
vi. Conversely, most endothermic reactions are nonspontaneous, and
require constant input of energy to keep going. (Think of ice; if you
keep your ice in the freezer, you prevent outside energy from getting
to it.)
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Unit 9b: Equilibrium Systems
Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
(-) sign is
exothermic;
gives off heat;
Red = fire (feels hot!)
(+) sign is
endothermic;
absorbs heat;
Blue = ice (feels cold!)
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Unit 9b: Equilibrium Systems
Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Topic:
Regents Chemistry ’14-‘15
Mr. Murdoch
Enthalpy
Objective: How do we express the potential energy in a system?
Enthalpy in common situations:
 Imagine you are on a bike at the top of a hill. If you push off and
pick your feet up, and you could coast down the hill without any
additional energy (well, you should steer!) After the initial ‘push’
(EA), the reaction (you and the bike) are on a spontaneous exothermic
change.
 Now, you are on the bottom of the hill and need to ride back to the
top. You have the same initial ‘push’ (EA), but you need to expend
almost continuous energy to climb the hill. If you stop pedaling, the
reaction (you and the bike) will stop, and probably reverse itself (roll
down the hill.) This is a nonspontaneous endothermic change.
Watch Crash Course Chemistry Enthalpy video - 11:23
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Unit 9b: Equilibrium Systems
Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Topic:
Regents Chemistry ’14-‘15
Mr. Murdoch
Entropy
Objective: How do we express the amount of order in a system?
Entropy:
The randomness (disorder) of a system is called Entropy. Nature favors
reactions that increase entropy.
 As a substance increases in temperature, the substance undergoes
an increase in entropy as well. As each subsequent phase change
occurs, the randomness (disorder) of the particles of that substance
increases.
 In order of LEAST to MOST entropy, the phases are: solid 
liquid gas. Solids are locked in a lattice, and gases have very
random movement controlled only by the confines of their
container. Liquids fall in between.
 As nature favors entropy, nature favors increases in phase.
Kitchen
Freezer
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Unit 9b: Equilibrium Systems
Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Reactants are
a solid(s) AND
a gas(g);
product is a
solid(s) –
Entropy
DECREASED,
and is
UNFAVORED
Reactants are
a solid(s) AND
a gas(g);
product is a
gas(g) –
Entropy
INCREASED,
and is
FAVORED
Watch Bozeman Chemistry Entropy video - 7:04
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Unit 9b: Equilibrium Systems
Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Topic:
Regents Chemistry ’14-‘15
Mr. Murdoch
Spontaneous Reactions
Objective: What factors will drive a reaction to complete on its own?
Spontaneous Reactions:
 For a reaction to occur spontaneously, both Enthalpy AND Entropy
will be considered.
i. Nature favors DECREASING enthalpy (exothermic processes);
ii. Nature favors INCREASING entropy (phase change to less order)
 If both enthalpy and entropy are favored, then the reaction will be
spontaneous.
1. Favored Reactions:
A favored reaction will be spontaneous at all temperatures.
i. This reaction has a ∆H of -84.0 kJ/mole of C2H6(g) produced,
enthalpy decreases, is exothermic, and is favored.
ii. This reaction starts with a solid and a gas, and ends with only gas.
Entropy increases, and is favored.
iii. Both enthalpy and entropy are favored, and this reaction will
ALWAYS be spontaneous at any temperature.
iv. No additional energy after EA will be needed to complete the
reaction.
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Unit 9b: Equilibrium Systems
Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Topic: Nonspontaneous Reactions
Objective: What factors will drive a reaction to stop on its own?
2. Unfavored Reactions:
An unfavored reaction will be nonspontaneous at all temperatures.
An unfavored reaction will require constant energy input to complete.
i. This reaction has a ∆H of +33.2 kJ/mole of NO2(g) produced,
enthalpy increases, is endothermic, and is unfavored.
ii. This reaction starts with a total of three moles of gaseous reactants,
and ends with only two moles of gaseous product. While mass is
conserved, the number of particles has decreased, or entropy
decreased, which is against nature, and therefore unfavored.
iii. Both enthalpy and entropy are unfavored, and this reaction will
ALWAYS be nonspontaneous at any temperature.
iv. After EA is input, continuous energy will be required to maintain
this reaction.
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Unit 9b: Equilibrium Systems
Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Topic: Partially Spontaneous Rx’s
Objective: What combinations of factors will drive a reaction?
3. Lower-Temperature favored Reactions:
If enthalpy is favored, but entropy is unfavored, the reaction will be
spontaneous at lower temperatures.
CO2(g) has a sublimation/deposition temperature of near 195 K,
meaning below that it is in the solid phase.
i. This reaction has a ∆H of -283.0 kJ/mole of CO2(g) produced, so
this reaction is exothermic, which is favored.
ii. This reaction starts with a total of three moles of gas, and ends
with only two moles of gas. Entropy decreases, and is unfavored.
Below 195 K entropy decreases even more (becomes solid), further
unfavoring the reaction.
iii. This reaction will be spontaneous only at temperatures when the
product is in the gaseous phase. Below 195 K the entropy
decreases more, and the reaction will be nonspontaneous.
Additional energy after EA will be needed to complete the
reaction.
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Unit 9b: Equilibrium Systems
Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Topic: Partially Spontaneous Rx’s
Objective: What combinations of factors will drive a reaction?
4. Higher-Temperature favored Reactions:
If enthalpy is unfavored, but entropy is favored, the reaction will be
spontaneous at higher temperatures.
H2O
i. This reaction has a ∆H of +25.69 kJ/mole of NH4NO3(s) when
decomposed in water, so this reaction is endothermic, which is
unfavored.
ii. This reaction starts out as a solid, but dissolves in aqueous
solution. In an aqueous solution, ions may move freely, and
entropy increases, which is favored.
iii. This reaction will be spontaneous only at temperatures when the
product is in the aqueous phase. Below about 273 K the entropy
decreases more (becomes solid), and the reaction will be
nonspontaneous.
Additional energy after EA will be needed to complete the
reaction, mostly to keep the water from freezing.
Watch Bozeman Science Spontaneous Processes video - 7:42
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Unit 9b: Equilibrium Systems
Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Topic:
Regents Chemistry ’14-‘15
Mr. Murdoch
Spontaneity of Water
Objective: Why does water ice melt above 273 K?
5. Spontaneity of the melting of water ice:
The reaction for the melting of water ice to liquid water is:
H2O(s)  H2O(l)
+ 6.01 kJ (334 J/g x 18.0 g/mole)
i. This process has a +∆H (the Heat of Fusion of water), is
endothermic, and enthalpy increases, which is unfavored.
ii. This process starts with a solid, and ends with a liquid, which is a
phase change to less order, so entropy increases, which is
favored.
iii. The combination of increasing enthalpy and increasing entropy
makes for a process that is spontaneous at higher temperatures.
Below 273 K the process is nonspontaneous, but above 273 K the
process will maintain without any additional energy. Once any
ice melts, the kinetic energy in the water is greater than the
kinetic energy in ice, and melting will continue.
*Note: This is DIFFERENT than water right at the freezing
point of 273 K/ 0.0°C. Water AT a temperature of 273 K/ 0.0°C
would equally freeze/melt if all other factors are kept the same.
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Unit 9b: Equilibrium Systems
Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Enthalpy & Entropy Practice Regents Problems: (ungraded)
1. According to Reference Table I, which reaction below spontaneously forms a
compound from its reactants?
a) H2(g) + I2(g)  2 HI(g)
c) 2 H2(g) + O2(g)  2 H2O(g)
b) N2(g) + O2(g)  2 NO(g)
d) N2(g) + 2 O2(g)  2 NO2(g)
2. Which change below is exothermic?
a) Melting of iron
b) Freezing of water
c) Sublimation of iodine
d) Vaporization of ethanol
3. Which reaction below has the greatest increase in entropy?
a) H2O(g)  H2O(l)
c) 2 H2O(g)  2 H2(g) + O2(g)
b) H2O(l)  H2O(s)
d) 2 H2O(l)  2 H2(g) + O2(g)
4. According to Reference Table I, which compound decreases in enthalpy as it
dissolves?
a) NaCl
c) KNO3
b) LiBr
d) NH4NO3
Given the reaction: 2 Na(s) + Cl2(g)  2 NaCl(s)
5. As the reactants form products, the entropy of the chemical system will
a) Increase
b) Decrease
c) Remain the same
6. Which chemical reaction will always be spontaneous?
a) An exothermic reaction in which entropy increases.
b) An exothermic reaction in which entropy decreases.
c) An endothermic reaction in which entropy increases.
d) And endothermic reaction in which entropy decreases.
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Unit 9b: Equilibrium Systems
Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Student name: _________________________
Class Period: _______
Please carefully remove this page from your packet to hand in.
Enthalpy, Entropy, and Reaction Spontaneity Homework:
For the given reactions below, state if entropy increases or decreases, and if the
change is favored or unfavored. 1 pt. ea.
Reaction
Entropy: Inc or Dec? Change: Fav or Unfav?
CO2(s)  CO2(g)
I2(g)  I2(s)
C(s) + O2(g)  CO2(g)
4 Al(s) + 3 O2(g)  2 Al2O3(s)
2 CO(g) + O2(g)  2 CO2(g)
2 H2(g) + O2(g)  2 H2O(l)
For the given reactions below, state if the reactions are spontaneous,
nonspontaneous, or at equilibrium, and also state whether enthalpy and entropy
are favored or unfavored. 1 pt. ea.
Spont, Nonspont, Enthalpy & entropy
Fav or Unfav?
or at Equil
Reaction
4 Al(s) + 3 O2(g)  2 Al2O3(s) + 3351 kJ
Enthalpy: _____
Entropy: _____
2 CO(g) + O2(g)  2 CO2(g) + 566 kJ
Enthalpy: _____
Entropy: _____
NaOH(s)  Na+1(aq) + Cl-1(aq) + 44.51 kJ
Enthalpy: _____
Entropy: _____
2 C(s) + 2 H2(g) + 52.4 kJ  C2H4(g)
Enthalpy: _____
Entropy: _____
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Unit 9b: Equilibrium Systems
Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Answer the following questions about water freezing. 1 pt. ea.
Pure water freezes at Standard Pressure at temperatures of 273 K/0.0°C or below,
as shown by the reaction: H2O(l)  H2O(s) + 6.01 kJ
1. Is this process an increase or decrease in entropy? _______________
2. Explain your answer for question #1 above.
3. Is the change in entropy favored or unfavored? _______________
4. Explain your answer for question #3 above.
5. When water freezes, is it exothermic or endothermic? _______________
6. Explain your answer for question #5 above.
7. Is the change in enthalpy favored or unfavored? _______________
8. Explain your answer for question #7 above.
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Unit 9b: Equilibrium Systems
Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Topic:
Regents Chemistry ’14-‘15
Mr. Murdoch
Changing Equilibrium
Objective: How can we change system equilibrium for our benefit?
Equilibrium Changes:
 Equilibrium systems are dynamic; this means that they are
continuously in some form of change. However, the reaction MUST
be in a closed system, or equilibrium cannot be maintained.
 What if you want to change the equilibrium in a system? A system at
equilibrium would be forming products at the same rate as the
products would be decomposing back into the starting reactants.
 We can change ONE aspect of a system at equilibrium at a time and
force the system to do what WE want; we can control the reaction.
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Unit 9b: Equilibrium Systems
Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Topic:
Regents Chemistry ’14-‘15
Mr. Murdoch
Le Chatelier’s Principle
Objective: What does changing an equilibrium do to a system?
Le Chatelier’s Principle:
“If a dynamic equilibrium is disturbed by changing
the conditions, the position of equilibrium shifts to
counteract the change to reestablish an equilibrium.”
 Le Chatelier’s Principle may be paraphrased to say this:
o If a system at equilibrium has some a stressor (any factor that
changes reaction rate), the equilibrium for that system will shift in a
way that lessens the added stress, in the direction of whichever
reaction rate was increased by the stressor.
o Stressors in chemistry include:
i. Temperature;
ii. Concentration;
iii. Pressure (only for gasses);
iv. Murdoch
 Any stressor introduced may cause a change in the concentrations of
both the reactants AND the product until equilibrium is restored at a
new point.
Watch Crash Course Chemistry Equilibrium video - 10:56
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Unit 9b: Equilibrium Systems
Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Stressor
Regents Chemistry ’14-‘15
Mr. Murdoch
Shift
Increases the number of collisions
between reactant particles, driving the
forward reaction faster
Change on Concentration
Removing
reactant
Decreases the number of collisions
between reactant particles, driving the
reverse reaction faster
Reactants: Increases
Adding
product
Increases the number of collisions
between product particles, driving the
reverse reaction faster
Reactants: Increase
Removing
product
Decreases the number of collisions
between product particles, driving the
forward reaction faster
Reactants: Decrease
Increasing
temperature
Favors the endothermic reaction,
shifting the equilibrium away from the
heat to absorb the excess heat energy
Depends on Direction of
shift:
Adding
reactant
Reactants: Decrease
Products: Increase
Products: Decrease
Products: Decrease
Products: Increase
Favors the exothermic reaction, shifting If the shift is towards the
products, then products
the equilibrium towards the heat to
will increase and reactants
release energy
will decrease
Increasing
System shifts towards side with fewer
pressure
If the shift is towards the
moles of gas to reduce pressure
(gases only)
reactants, then reactants
will increase and products
Decreasing
System shifts toward side with more
will decrease
pressure
mores of gas to increase pressure
(gases only)
Decreasing
temperature
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Unit 9b: Equilibrium Systems
Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Equilibrium Shift
Topic:
Objective: How do we determine the possibility of equilibrium shift?
 Some factors that may affect the rate of a reaction have NO effect on
systems at equilibrium. These factors include:
i. Catalysts
ii. Inhibitors
iii. Surface area
 The above factors allow a system to achieve equilibrium faster, but
once equilibrium is established, these factors affect both reactions
equally. The equilibrium would not change then.
Water levels at
equilibrium
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Add water to left
container; raises
left level
Water levels find
new equilibrium;
higher levels
Remove added
water from left
container; lowers
left water level
Page 36 of 45
Water levels find
new equilibrium
Unit 9b: Equilibrium Systems
Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Topic: Determining Shift Direction
Objective: How do we determine the direction of equilibrium shift?
Application of Determining Direction of Equilibrium Shift:
For the equilibrium: N2(g) + 3 H2(g)  2 NH3(g) + heat
Stressor
Shift
Add N2(g)
Forwards-add a reactant; shifts to
products
Remove N2(g)
Reverse-remove a reactant; shifts to
reactants
Add H2(g)
Forwards-add a reactant; shifts to
products
Remove H2(g)
Reverse-remove a reactant; shifts to
reactants
Add NH3(g)
Reverse-add a product; shift to
reactants
Remove NH3(g)
Forwards-remove a product; shift to
products
Increase Temp
Reverse-increase Temp; shift away
from heat
Decrease Temp
Forwards-decrease Temp; shift
towards heat
Increase Press
Forwards-increase Press; shifts to
side with fewer moles of gas
Decrease Press
Reverse-decrease Press, shifts to
side with more moles of gas
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Change in Concentration
N2(g): decreases
H2(g): decreases
NH3(g): increases
N2(g): increases
H2(g): increases
NH3(g): decreases
N2(g): decreases
H2(g): decreases
NH3(g): increases
N2(g): increases
H2(g): increases
NH3(g): decreases
N2(g): increases
H2(g): increases
NH3(g): decreases
N2(g): decreases
H2(g): decreases
NH3(g): increases
N2(g): increases
H2(g): increases
NH3(g): decreases
N2(g): decreases
H2(g): decreases
NH3(g): increases
N2(g): decreases
H2(g): decreases
NH3(g): increases
N2(g): increases
H2(g): increases
NH3(g): decreases
Unit 9b: Equilibrium Systems
Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Topic:
Regents Chemistry ’14-‘15
Mr. Murdoch
Equilibrium Shift
Objective: How do we determine the direction of equilibrium shift?
For the equilibrium: KNO3(s) + 34.89 kJ  K
+1
(aq) +
NO3-1(aq)
1. What happens to the concentration of K+1(aq) when temperature is
increased?
i. Stressor: increased temperature
ii. Shift: away from heat input (forward)
iii. Change in Concentration: Since the shift is towards K+1, the
concentration of K+1(aq) increases (along with [NO3-1(aq)]
2. What happens to the concentration of NO3-1(aq) when the temperature
is decreased?
i. Stressor: decreasing temperature
ii. Shift: towards heat input (reverse)
iii. Change in Concentration: Since the shift is away from NO3-1(aq),
the concentration of NO3-1(aq) decreases (along with [K+1(aq)]
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Unit 9b: Equilibrium Systems
Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Topic:
Regents Chemistry ’14-‘15
Mr. Murdoch
Equilibrium Shift
Objective: How do we determine the direction of equilibrium shift?
For the equilibrium: 2 CO(g) + O2(g)  2 CO2(g) + 566 kJ
1. What happens to the concentration of CO2(g) when CO(g) is added to
the equilibrium system?
i. Stressor: increasing concentration of a reactant
ii. Shift: away from reactant (forward)
iii. Change in Concentration: Since the shift is towards CO2(g), the
concentration of CO2(g) increases
2. What happens to the concentration of O2(g) when CO2(g) is removed
from the equilibrium system?
i. Stressor: decreasing concentration of a product
ii. Shift: towards product (forward)
iii. Change in Concentration: Since the shift is away from O2(g), the
concentration of O2(g) decreases
3. What happens to the concentration of CO(g) when pressure is
increased?
i. Stressor: increasing pressure
ii. Shift: towards side with fewer moles of gas (forwards)
iii. Change in Concentration: Since the shift is away from CO(g), the
concentration of CO(g) decreases
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Unit 9b: Equilibrium Systems
Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Topic:
Regents Chemistry ’14-‘15
Mr. Murdoch
Equilibrium Shift
Objective: How do we determine the direction of equilibrium shift?
For the equilibrium: N2(g) + 2 O2(g) + 66.4 kJ  2 NO(g)
1. State five (5) things that can be done to the equilibrium that will
result in an increase of the concentration of NO(g).
Desired shift: to make more NO(g), you must shift towards NO(g), so
shift equilibrium forwards
 How can we drive the equilibrium forwards?
a. Add N2(g): (adding reactant drives the reaction forward)
b. Add O2(g): (adding reactant drives the reaction forward)
c. Remove NO(g): (removing product drives the reaction forwards)
d. Increase Temperature: (adding heat makes the reaction shift
away from heat)
e. Increase Pressure: (adding pressure shifts the equilibrium
towards the side with fewer moles of gas)
For the equilibrium: NaCl(s) + 3.88 kJ  Na+1(aq) + Cl-1(aq)
2. State four (4) things that can be done to increase the concentration of
NaCl(s).
Desired shift: to make more NaCl(s), you must shift towards NaCl(s),
so shift equilibrium reverse
 How can we drive the equilibrium in reverse?
a. Remove NaCl: (removing a reactant makes the reaction shift in
reverse)
b. Remove heat: (removing heat makes the reaction shift towards
the heat)
c. Increase Na+1(aq): (adding product with Na+1 ions drives the
reaction in reverse)
d. Increase Cl-1(aq): (adding a product with Cl-1 ions drive the
reaction in reverse)
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Unit 9b: Equilibrium Systems
Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Changing Equilibrium Practice Regents Problems: (ungraded)
1. Given the reaction at equilibrium:
2 SO2(g) + O2(g)  2 SO3(g) + heat
Which change will shift the equilibrium to the right?
a)
b)
c)
d)
Increasing the pressure
Increasing the temperature
Decreasing the amount of O2(g)
Increasing the amount of SO2(g)
2. Given the reaction at equilibrium:
N2(g) + O2(g) + 182.6 kJ  2 NO(g)
Which change would cause an immediate increase in the rate of the forward
reaction?
a)
b)
c)
d)
Decreasing the reaction pressure
Decreasing the reaction temperature
Increasing the concentration of N2(g)
Increasing the concentration of NO(g)
3. Given the equilibrium reaction:
X + Y  2 Z + heat
The concentration of the product may be increased by
a)
b)
c)
d)
Adding a catalyst
Adding more heat to the system
Decreasing the concentration of X
Increasing the concentration of Y
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Unit 9b: Equilibrium Systems
Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Notes page:
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Unit 9b: Equilibrium Systems
Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Student name: _________________________
Class Period: _______
Please carefully remove this page from your packet to hand in.
Naming and Writing Binary and Ternary Formulas Homework:
For each of the following systems at equilibrium, predict the effect of a given change on the
concentration of each of the specific substances.
Write I if the concentration increases, D if the concentration
decreases, and R if the concentration remains the same.
1. 2 NH3(g) + heat  N2(g) + 3 H2(g)
Stressor #1: increase in [N2(g)]
Direction of shift: _______________
What is the resulting effect on the concentration of:
[NH3(g)]: _______________
Stressor #2: increase in temperature
[H2(g)]: _______________
Direction of shift: _______________
What is the resulting effect on the concentration of:
[N2(g)]: _______________
Stressor #3: increase in pressure
[NH3(g)]: _______________
Direction of shift: _______________
What is the resulting effect on the number of moles of N2(g): _______________
What is the resulting effect on the number of moles of NH3(g): _______________
2. 2 NO(g)  N2(g) + O2(g) + heat
Stressor #1: decrease in [O2(g)]
Direction of shift: _______________
What is the resulting effect on the concentration of:
[N2(g)]: _______________
Stressor #2: decrease in temperature
[NO(g)]: _______________
Direction of shift: _______________
What is the resulting effect on the concentration of:
[O2(g)]: _______________
Stressor #3: increase in pressure
[NO(g)]: _______________
Direction of shift: _______________
What is the resulting effect on the number of moles of O2(g): _______________
What is the resulting effect on the number of moles of NO(g): _______________
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Unit 9b: Equilibrium Systems
Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Complete the following questions by circling the correct answer.
Given the equilibrium: N2(g) + 3 H2(g)  2 NH3(g) + heat
1. If N2(g) is added to the system at equilibrium, in which direction will the
equilibrium shift?
Forward
Reverse
2. If H2(g) is removed from the system at equilibrium, in which direction will the
equilibrium shift?
Forward
Reverse
3. If NH3(g) is added to the system at equilibrium, in which direction will the
equilibrium shift?
Forward
Reverse
4. If the temperature is decreased in the system at equilibrium, in which direction
will the equilibrium shift?
Forward
Reverse
5. If the pressure is increased in the system at equilibrium, in which direction will
the equilibrium shift?
Forward
Reverse
6. If H2(g) is removed from the system at equilibrium, what will happen to the
concentrations of:
N2(g):
Increase
Decrease
Remain the same
NH3(g):
Increase
Decrease
Remain the same
7. If NH3(g) is removed from the system at equilibrium, what will happen to the
concentrations of:
N2(g):
Increase
Decrease
Remain the same
H2(g):
Increase
Decrease
Remain the same
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Unit 9b: Equilibrium Systems
Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture
Regents Chemistry ’14-‘15
Mr. Murdoch
Notes page:
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Unit 9b: Equilibrium Systems