lecture Mr. Murdoch Unit 9b: Equilibrium, Enthalpy, … U… · Unit 9b: Equilibrium, Enthalpy, and Entropy ... Unit 9b: Equilibrium Systems Unit 9b Vocabulary: 1. ... Website upload

Unit 9: Kinetics, Thermodynamics, & Equilibrium-lecture Regents Chemistry ’14-‘15 Mr. Murdoch

Website upload 2015 Page 1 of 45 Unit 9b: Equilibrium Systems

Unit 9b:

Equilibrium, Enthalpy, and

Entropy

1.

Student Name: _______________________________________

Class Period: ________

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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.



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.



Notes page:



Unit 9b Homework Assignments:

Assignment: Date: Due:



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 = forward reaction When the rate of the

forward reaction equals

the rate of the reverse

reaction, a state of

equilibrium is reached.

C + D + energy A + B = reverse reaction A + B (±energy) C + D Double arrows ()

indicate that BOTH reactions are occurring 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.

Topic: Equilibrium

Objective: What is the role equilibrium has in chemistry?



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



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.

Topic: Equilibrium Properties

Objective: What properties of equilibrium will systems have?



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%



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.

Topic: Three Types of Equilibrium

Objective: What forms of equilibrium are possible in chemistry?



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.

Unsaturated - solute almost all undissolved; reaction almost all

forward

Close to Saturation - solute almost all dissolved; reaction mostly forward, some reverse

Saturated-dissolving rate is equal to precipitate formation

rate; no net change

Solution Equilibriums



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



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.



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

b) 100 K

c) 273 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

b) Heat of fusion

c) Melting point

d) Heat of vaporization



Notes page:



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.



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

c) Unsaturated

d) Concentrated



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

Topic: Reaction Progress

Objective: How do we know if a reaction will continue when started?

Expanding

Universe! Falling is

easier!



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.)

Topic: Enthalpy

Objective: How do we express the potential energy in a system?



(-) sign is

exothermic;

gives off heat;

Red = fire -

(feels hot!)

(+) sign is

endothermic;

absorbs heat;

Blue = ice -

(feels cold!)



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

Topic: Enthalpy

Objective: How do we express the potential energy in a system?

https://www.youtube.com/watch?v=SV7U4yAXL5I



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.

Topic: Entropy

Objective: How do we express the amount of order in a system?

Freezer

Kitchen



Watch Bozeman Chemistry Entropy video - 7:04

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

https://www.youtube.com/watch?v=MALZTPsHSoo



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.

Topic: Spontaneous Reactions

Objective: What factors will drive a reaction to complete 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.

Topic: Nonspontaneous Reactions

Objective: What factors will drive a reaction to stop on its own?



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.

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.

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

Topic: Partially Spontaneous Rx’s

Objective: What combinations of factors will drive a reaction?

H2O

https://www.youtube.com/watch?v=hNSD0YDsPsE



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.

Topic: Spontaneity of Water

Objective: Why does water ice melt above 273 K?



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)

b) N2(g) + O2(g) 2 NO(g)

c) 2 H2(g) + O2(g) 2 H2O(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)

b) H2O(l) H2O(s)

c) 2 H2O(g) 2 H2(g) + O2(g)

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

b) LiBr

c) KNO3

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.





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.

Reaction Spont, Nonspont,

or at Equil

Enthalpy & entropy

Fav or Unfav?

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) + OH-1

(aq) + 44.51 kJ Enthalpy: _____

Entropy: _____

2 C(s) + 2 H2(g) + 52.4 kJ C2H4(g) Enthalpy: _____

Entropy: _____ Cont’d next page



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? _______________


5. When water freezes, is it exothermic or endothermic? _______________


7. Is the change in enthalpy favored or unfavored? _______________




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.

Topic: Changing Equilibrium

Objective: How can we change system equilibrium for our benefit?



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

Topic: Le Chatelier’s Principle

Objective: What does changing an equilibrium do to a system?

https://www.youtube.com/watch?v=g5wNg_dKsYY



Stressor Shift Change on Concentration

Adding

reactant

Increases the number of collisions

between reactant particles, driving the

forward reaction faster

Reactants: Decrease

Products: Increase

Removing

reactant

Decreases the number of collisions

between reactant particles, driving the

reverse reaction faster

Reactants: Increases

Products: Decrease

Adding

product

Increases the number of collisions

between product particles, driving the

reverse reaction faster

Reactants: Increase

Products: Decrease

Removing

product

Decreases the number of collisions

between product particles, driving the

forward reaction faster

Reactants: Decrease

Products: Increase

Increasing

temperature

Favors the endothermic reaction,

shifting the equilibrium away from the

heat to absorb the excess heat energy

Depends on Direction of

shift:

If the shift is towards the

products, then products

will increase and reactants

will decrease

If the shift is towards the

reactants, then reactants

will increase and products

will decrease

Decreasing

temperature

Favors the exothermic reaction, shifting

the equilibrium towards the heat to

release energy

Increasing

pressure

(gases only)

System shifts towards side with fewer

moles of gas to reduce pressure

Decreasing

pressure

(gases only)

System shifts toward side with more

moles of gas to increase pressure



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

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

Water levels find

new equilibrium

Topic: Equilibrium Shift

Objective: How do we determine the possibility of equilibrium shift?



Application of Determining Direction of Equilibrium Shift:

For the equilibrium: N2(g) + 3 H2(g) 2 NH3(g) + heat

Stressor Shift Change in Concentration

Add N2(g) Forwards-add a reactant; shifts to

products

N2(g): decreases

H2(g): decreases

NH3(g): increases

Remove N2(g) Reverse-remove a reactant; shifts to

reactants

N2(g): increases

H2(g): increases

NH3(g): decreases

Add H2(g) Forwards-add a reactant; shifts to

products

N2(g): decreases

H2(g): decreases

NH3(g): increases

Remove H2(g) Reverse-remove a reactant; shifts to

reactants

N2(g): increases

H2(g): increases

NH3(g): decreases

Add NH3(g) Reverse-add a product; shift to

reactants

N2(g): increases

H2(g): increases

NH3(g): decreases

Remove NH3(g) Forwards-remove a product; shift to

products

N2(g): decreases

H2(g): decreases

NH3(g): increases

Increase Temp Reverse-increase Temp; shift away

from heat

N2(g): increases

H2(g): increases

NH3(g): decreases

Decrease Temp Forwards-decrease Temp; shift

towards heat

N2(g): decreases

H2(g): decreases

NH3(g): increases

Increase Press Forwards-increase Press; shifts to

side with fewer moles of gas

N2(g): decreases

H2(g): decreases

NH3(g): increases

Decrease Press Reverse-decrease Press, shifts to

side with more moles of gas

N2(g): increases

H2(g): increases

NH3(g): decreases

Topic: Determining Shift Direction

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)]





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





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)





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) Increasing the pressure

b) Increasing the temperature

c) Decreasing the amount of O2(g)

d) Increasing the amount of SO3(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) Decreasing the reaction pressure

b) Decreasing the reaction temperature

c) Increasing the concentration of N2(g)

d) 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) Adding a catalyst

b) Adding more heat to the system

c) Decreasing the concentration of X

d) Increasing the concentration of Y



Notes page:





Changing Equilibrium 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)]: _______________ [H2(g)]: _______________

Stressor #2: increase in temperature Direction of shift: _______________


[N2(g)]: _______________ [NH3(g)]: _______________

Stressor #3: increase in pressure 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: _______________


[N2(g)]: _______________ [NO(g)]: _______________

Stressor #2: decrease in temperature Direction of shift: _______________


[O2(g)]: _______________ [NO(g)]: _______________

Stressor #3: increase in pressure 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): _______________

Cont’d next page



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



Notes page:

lecture Mr. Murdoch Unit 9b: Equilibrium, Enthalpy, … U… · Unit 9b: Equilibrium, Enthalpy, and Entropy ... Unit 9b: Equilibrium Systems Unit 9b Vocabulary: 1. ... Website upload

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