Chemical Equilibrium I. A State of Dynamic Balance 1N 2 (g) 3H 2(g) +2NH 3(g) ΔG 0 = -33.1 kJ The...

Chemical Equilibrium

I. A State of Dynamic Balance

1N2 (g) 3H2(g)+ 2NH3(g) ΔG0 = -33.1 kJ The reaction isspontaneous

under standardconditions

The concentrationsof the reactants

decrease at first…

…while the concentrationof the product increases

But, then, before thereactants are used up,

all concentrationsbecome constant


I. A State of Dynamic Balance-when a ________ results in the almost ________ conversion of ________ to ________, the ________ is said to go to __________, but _____ _________ ___ ____ go to __________, most _________ are __________

reactioncomplete

reactants productsreactioncompletion most reactionsdo not completionreactions reversible

1N2 (g) 3H2(g)+ 2NH3(g)

1N2 (g) 3H2(g)+ 2NH3(g)

1N2 (g) 3H2(g)+ 2NH3(g)

At first, only the reactantsare present, so only the

forward reaction can occur

1N2 (g) 3H2(g)+ 2NH3(g)


I. A State of Dynamic Balance-as soon as the ________ ________ begins, the ____________ of the _________ go _____, and the _________ _____ goes _____ as the number of __________ per unit ____ goes _____

forward reactionconcentrations

reactants downreaction rate down

collisionstime down

As soon as the productsbegin forming, the forwardreaction rate slows and the

reverse reaction begins

1N2 (g) 3H2(g)+ 2NH3(g)


I. A State of Dynamic Balance-as the _________ proceeds, the ____ of the ________ _________ continues to ________ and the ____ of the ________ ________ continues to ________ until the two _____ are _____, and the system has reached a state of ________ __________

reactionrate forward reaction

decrease ratereverse reaction

increase ratesequal

chemical equilbrium1N2 (g) 3H2(g)+ 2NH3(g)


I. A State of Dynamic Balance-at ___________, the ____________ of the ________ and ________ are not _____, but _______, because the ____ of _________ of the ________ is _____ to the ____ of _________ of the ________

equilibrium concentrationsreactants products

equal constantrate formation products

equal rate formationreactants

The Golden Gate Bridgeconnects San Francisco

to Sausalito.

If all other roads leadingin and out of the two cities

were closed…

…and the number of vehiclescrossing the bridge per hourin one direction equaled thenumber of vehicles crossing

the bridge in the oppositedirection…

What is true of the numberof vehicles in each city throughout

the day?

Are there the same number ofvehicles in each city?


II. Equilibrium Expressions and Constants

Cato Maximilian Guldberg1836-1902

Peter Waage1833-1900

-while _____ chemical systems have little tendency to _____, and _____ chemical systems _____ readily and ___ to __________, _____ chemical systems reach a _____ of __________, leaving varying amounts of ________ ____________

somereact somereact

go completion moststate equilibrium

reactantunconsumed

-in 1864, Norwegian chemists ______ and _________ proposed the _______ ___________________, which states, at a given ___________, a chemical system may reach a _____ in which a particular _____ of _______ and _______ ____________ has a _______ value

WaageGuldberg Law of

Chemical Equilibriumtemperature

stateratio reactant

product concentrations constant



-the _______ ________ for a _______ at __________ can be written ______________________________, where __ and __ are ________, __ and __ are ________, __, __, __, and __ are the ___________ in the ________ ________, and the __________ _______ __________ is

general equation reactionequilibrium

aA + bB cC + dDA B reactants C

D products a b c dcoefficients

balancedequation equilibriumconstant expression

Keq =[C]c[D]d

[A]a[B]b

-___________ ________ with ___ values __ __ contain more ________ than ________ at ___________, while __________ ________ with ___ values __ __ contain more ________ than ________ at __________

equilibrium mixtures Keq

> 1 productsreactants equilibrium

equilibrium mixtures Keq

< 1 reactantsproducts equilibrium



Write the equilibrium constant expression for the homogeneous equilibrium for the synthesis of ammonia from nitrogen and hydrogen.

1N2 (g) 3H2(g)+ 2NH3(g)

Keq =[C]c[D]d

[A]a[B]b

Keq = [C]c

[A]a[B]b= [NH3]c

[N2]a[H2]b= [NH3]2

[N2]1[H2]3

The equilibrium is homogeneousbecause all the reactants and

products are in the same physicalstate (gas)



Write the equilibrium constant expression for the equilibrium for the synthesis of Hydrogen iodide from iodine and hydrogen.

1H2 (g) 1I2(g)+ 2HI(g)

Keq =[C]c[D]d

[A]a[B]b

Keq = [C]c

[A]a[B]b= [HI]c

[H2]a[I2]b= [HI]2

[H2]1 [I2]1





Write the equilibrium constant expression for the equilibrium for the decomposition of Dinitrogen tetroxide into Nitrogen dioxide.

1N2O4 (g) 2NO2(g)

Keq =[C]c[D]d

[A]a[B]b

Keq = [C]c

[A]a=

[N2O4]a

[NO2]c=

[N2O4]1

[NO2]2





Write the equilibrium constant expression for the equilibrium for the reaction of Carbon monoxide and Hydrogen which produces methane (Tetrahydrogen monocarbide) and water.

3H2 (g) 1CH4(g)

Keq =[C]c[D]d

[A]a[B]b

Keq = [C]c

[A]a= [CH4]c

[CO]a=

1CO(g) + 1H2O (g)+

[D]d

[B]b

[H2O]d

[H2]b

[CH4]1

[CO]1

[H2O]1

[H2]3



Write the equilibrium constant expression for the equilibrium for the decomposition of Dihydrogen monosulfide into diatomic hydrogen and diatomic sulfur.

2H2S (g) 2H2 (g)

Keq =[C]c[D]d

[A]a[B]b

Keq = [C]c

[A]a= [H2]c

[H2S]a=

1S2 (g)+

[D]d [S2]d [H2]2 [S2]1

[H2S]2



-_________ in which all ________ and ________ are in the same ________ _____ are ____________, but ________ with _________ and ________ in _____ than ___ ________ _____ result in _____________ _________

equilibria reactantsproducts physicalstate homogeneousreactions reactantsproducts more one physicalstate heterogeneousequilibriagaseous

ethanol

liquid ethanol

1C2H5OH (l) 1C2H5OH (g)

Keq =[C]c[D]d

[A]a[B]b

Keq =[C2H5OH (g)]1

[C2H5OH (l)]1

Keq = [C2H5OH (g)]1



-since ______ and _____ ________ and ________ don’t change ___________, (which is really their ______), if the ___________ remains ________, then in the ___________ _______ __________ for a ____________ ___________, the ___________ ________ only depends on the ______________ of the ________ and ________ in the _______ state of matter

liquid solid reactantsproducts concentration

densitytemperature constant

equilibrium constantexpression heterogeneousequilibrium equilibriumconstantconcentrations reactantsproducts gaseous



Write the equilibrium constant expression for the heterogeneous equilibrium for the decomposition of Sodium Hydrogen carbonate into Sodium carbonate, Carbon dioxide, and water.

2NaHCO3 (s) 1Na2CO3 (s)

Keq =[C]c[D]d

[A]a

[E]e

Keq = [D]d = [CO2]d [H2O]e = [CO2]1[H2O]1

The equilibrium is heterogeneousbecause the reactants and

products are in different physicalstates (gas and solid)

+ 1CO2 (g) 1H2O (g)+

[E]e



Write the equilibrium constant expression for the heterogeneous equilibrium for the decomposition of Calcium carbonate into Calcium oxide and Carbon dioxide.

1CaCO3 (s) 1CaO (s)

Keq =[C]c[D]d

[A]a

Keq = [D]d = [CO2]d = [CO2]1

The equilibrium is heterogeneousbecause the reactants and

products are in different physicalstates (gas and solid)

+ 1CO2 (g)



Write the complete, balanced thermochemical equation and equilibrium constant expression for the homogeneous equilibrium for the reaction of hydrazine (Tetrahydrogen dinitride) and Nitrogen dioxide, which produces nitrogen and water.

2NO2 (g) 3N2 (g)

Keq =

2N2H4 (g) + 4H2O (g)+

[N2]3

[N2H4]2

[H2O]4

[NO2]2



Write the complete, balanced thermochemical equation and equilibrium constant expression for the homogeneous equilibrium for the reaction of Sulfur trioxide and Carbon dioxide, which produces Carbon disulfide and oxygen.

1CO2 (g) 1CS2 (g)

Keq =

2SO3 (g) + 4O2 (g)+

[CS2]1

[SO3]2

[O2]4

[CO2]1



Write the complete, balanced thermochemical equation and equilibrium constant expression for the heterogeneous equilibrium for the reaction of monatomic Sulfur and fluorine gas, which produces Sulfur tetrafluoride gas and Sulfur hexafluoride gas.

5F2 (g) 1SF4 (g)

Keq =

2S (s) + 1SF6 (g)+

[SF4]1 [SF6]1

[F2]5



Write the complete, balanced thermochemical equation and equilibrium constant expression for the heterogeneous equilibrium for the reaction of magnatite (Fe3O4) and hydrogen gas, which produces iron and water vapor.

4H2 (g) 3Fe (s)

Keq =

1Fe3O4 (s) + 4H2O (g)+

[H2O]4

[H2]4



Write the equilibrium constant expression for the homogeneous equilibrium for the synthesis of ammonia and calculate the value of Keq when [NH3] = 0.933 M, [N2] = 0.533 M, and [H2] = 1.600 M.

1N2 (g) 3H2(g)+ 2NH3(g)

Keq = [NH3]2

[N2]1[H2]3

Keq = [0.933]2

[0.533]1 [1.600]3

Keq = 0.399



Write the equilibrium constant expression for the homogeneous equilibrium for the decomposition of Sulfur trioxide into Sulfur dioxide and oxygen gas, and calculate the value of Keq when [SO3] = 0.0160 M, [SO2] = 0.00560 M, and [O2] = 0.00210 M.

2SO3 (g) 1O2 (g)+2SO2 (g)

Keq = [SO2]2[O2]1

[SO3]2

Keq = [0.00560]2 [0.00210]1

[0.0160]2

Keq = 2.58 x 10-4


III. Le Châtelier’s Principle

A. Safety:

1. Hypothesis: What is the effect of temperature on equilibrium?

2. Prediction:

3. Gather Data:

The surfaces of the hot plates and the water will be hot enough to cause burns. Use caution. Cobalt(II) chloride hexahydrate is toxic, with an LD50 = 80mg/kg Avoid ingestion (don’t eat or drink it). Wash handsthoroughly with soap and water before leaving lab. Ethanol is extremely flammable. No open flame.

B. Procedure:

1. Pick up a sheet of white construction paper and an artist’s paintbrush.

http://www.safetyemporium.com/?03032

3. Gather Data:

B. Procedure:

4. Use the solution to paint a winter scene on your white construction paper, including a pink-colored field of snow.

3. Add 10 mL of ethanol to the test tube, cap, and shake vigorously until CoCl2·6H2O dissolves. If the solution is not light pink, add water dropwise until it turns light pink.



2. With a partner, using a top-loading electronic balance, mass 0.3 grams of CoCl2·6H2O, crush it into a fine powder using a mortar and pestle, and place it in a test tube.

5. To simulate the coming of spring, warm your painting over the hotplate in the fume hood. Record your observations.

4. Analyze Data:



4Cl- (aq) 1CoCl4

2-(aq) + 6H2O (l)1Co(H2O)6

2+ (aq) +

Hexahydrate Co2+ ion (pink) chloride ion Tetrachlorocobaltate ion (blue)

+ heat

4Cl- (aq) 1CoCl4

2-(aq) + 6H2O (l)1Co(H2O)6

2+ (aq) +


5. Draw Conclusions:



Henry-Louis Le Châtelier1850-1936

-in 1888, ________________________ discovered that there are ways to _______ _________ in order to make _________ more __________

Henry-Louis Le Châtelier

control equilibriareactions productive

-____________________ states that if a ______ (like a ______ in __________) is applied to a system at __________, the system _____ in the ________ that _______ the _____

Le Châtelier’s Principlestress changetemperature

equilibrium shiftsdirection relieves

stress

-________ that reach __________ instead of going to __________ do not ________ as much

reactions equilibriumcompletion

produce


III. Le Châtelier’s Principle -stressors that cause a shift in equilibrium

A. Changes in Concentration

Write the equilibrium constant expression for the equilibrium for the reaction of Carbon monoxide and Hydrogen to produce methane and water. Then, calculate the Keq value when [CO] = 0.30000 M, [H2] = 0.10000 M, and [CH4] = 0.05900 M, and [H2O] = 0.02000 M.

3H2 (g) 1CH4(g)

Keq =

1CO(g) + 1H2O (g)+

[CH4]1

[CO]1

[H2O]1

[H2]3

Keq = [0.05900]1

[0.30000]1

[0.02000]1

[0.10000]3= 3.933 = first equilibrium

position



-_________ the ____________ of ___ _________ the _______ of _________ between ___ and ___, _________ the _____ of the _______ _______


3H2 (g) 1CH4(g)1CO(g) + 1H2O (g)+

increasing concentrationCO increases numbercollisions COH2 increasing rateforward reaction

-the system responds to the ______ of the addition of _______ by forming more _______ to bring the system back into equilbrium

stressreactantproduct




Keq = [CH4]1

[CO]1

[H2O]1

[H2]3

Keq = [0.06648]1

[0.99254]1

[0.02746]1

[0.07762]3= 3.933

3H2 (g) 1CH4(g)1CO(g) + 1H2O (g)+0.99254 M 0.07762 M 0.06648 M 0.02746 M

= second equilibriumposition



A. Changes in Concentration -_________ the ____________ of a ________ causes __________ to _____ to the ____ to _______ the ____ of formation of ______

increasing concentrationreactant equilbriumshift right increaserate product

3H2 (g) 1CH4(g)1CO(g) + 1H2O (g)+

3H2 (g) 1CH4(g)1CO(g) + 1H2O (g)+

-_________ the ____________ of a ________ causes __________ to _____ to the ____ to _______ the ____ of formation of ______

decreasing concentrationproduct equilbriumshift right increaserate product




4Cl- (aq) 1CoCl4

2-(aq) + 6H2O (l)1Co(H2O)6

2+ (aq) +


Predict what should happen to the following equilibrium if hydrogen bonding due to the addition of acetone binds water and effectively removes it from the products.



A. Changes in Volume -_________ the ______ of the _______ container, according to ______, ________ the ________, which in turn ________ the _____ of _________ between the ________ of the ________, _________ the _____ of the ________ _______

3H2 (g)

1CH4(g)

1CO(g) +

1H2O (g)+

decreasing volume reactionBoyle

increases pressureincreases rate collision

particlesreactants increasing rateforward reaction

-the _____ in the _________ causes the _____ on the system to be _______ as for every __ _____ of _______ _______ _________, only __ _____ of _______ _______ are _________, which, according to ________, occupies __ the ______, which _________ the ________

shift equilibriumstress relieved

4 moles gaseous reactantconsumed 2 moles gaseousproduct produced

Avogadro ½volume decreases pressure



Use Le Châtelier’s Principle to predict how each of these changes would affect the ammonia equilibrium system.

1N2 (g) 3H2(g)+ 2NH3(g)

a. removing hydrogen from the system __________________________

b. adding ammonia to the system _______________________________

1N2 (g) 3H2(g)+ 2NH3(g)

equilibrium shifts to the left

1N2 (g) 3H2(g)+ 2NH3(g)

equilibrium shifts to the left



Use Le Châtelier’s Principle to predict how each of these changes would affect the ammonia equilibrium system.

1N2 (g) 3H2(g)+ 2NH3(g)

c. adding hydrogen to the system _______________________________

1N2 (g) 3H2(g)+ 2NH3(g)

equilibrium shifts to the right



How would decreasing the volume of the reaction container affect each of these equilibria?

equilibrium shifts to the right2SO2 (g) 1O2(g)+ 2SO3(g)a. _________________________

stress has no effect on equilibrium1H2 (g) 1Cl2(g)+ 2HCl(g)b. _____________________________

equilibrium shifts to the left2NOBr(g) 1Br2(g)+2NO(g)c. _________________________



A. Changes in Temperature -while _______ in _____________ and ________ in _______ cause ______ in _________, they ___ ___ _______ the __________ _______, but a ______ in ___________ causes ______ in both the __________ ________ and the __________ _______

changes concentrationchanges volume

shifts equilibria donot change equilibriumconstant changetemperature change

equilibrium positionequilibrium constant



A. Changes in Temperature

ΔH0 = -206.5 kJ

3H2 (g)

1CH4(g)

1CO(g) +

1H2O (g)+

-since the _______ for making _______ has a _______ ____, the ________ _______ is _________, and the _______ _______ is __________, so ____ can be thought of as a _______ in the ________ _______ and a _______ in the _______ _______

reactionmethane negative ΔH0

forward reaction exothermicreverse reaction

endothermic heat

reactantforward reactionproduct

reverse reaction+ heat



A. Changes in Temperature -_________ the __________ is like _______ more _______ to the _______ in which _____ acts as a _______ and is _____ ___, in this case, the __________ _______ _______

increasing temperatureadding reactant

reaction heatused up

endothermic reversereaction

3H2 (g) 1CH4(g)1CO(g) + 1H2O (g)+ + heat

reactant

-__________ shifts to the _____, _________ the ___________ of _______ because _______ is a _______ in the _______ _______

equilibrium leftdecreasing concentrationmethane methanereactant reverse reaction



A. Changes in Temperature -_________ the __________ is like ________ _______ from the _______ in which _____ acts as a _______, in this case, the __________ _______ _______

decreasing temperatureremoving reactant

reaction heat

endothermic reverse reaction

3H2 (g) 1CH4(g)1CO(g) + 1H2O (g)+ + heat

reactant

-__________ shifts to the _____, _________ the ___________ of _______ because _______ is a _______ in the _______ _______

equilibrium rightincreasing concentrationmethane methaneproduct forward reaction



In the following equilibrium, would you raise or lower the temperature to get the following results?

1C2H2 (g) 1H2O(g)+ 1CH3CHO(g)

a. increase the amount of CH3CHO______________________________lower the temperature

ΔH0 = -151 kJ

1C2H2 (g) 1H2O(g)+ 1CH3CHO(g) + heat

b. decrease the amount of C2H2 ________________________________lower the temperature




In the following equilibrium, would you raise or lower the temperature to get the following results?

1C2H2 (g) 1H2O(g)+ 1CH3CHO(g)

c. increase the amount of H2O _________________________________raise the temperature

ΔH0 = -151 kJ




In the following equilibrium, what effect does decreasing the volume of the reaction vessel have?

1CO(g) 1Fe3O4(s)+ 1CO2 (g)

______________________________________________________________________________________________________________________________________________________________________________

None. The solids do not change their concentrations, and the numberof moles of gaseous reactant equals the number of moles of gaseousproduct

3FeO(s)+

1CO (g) 1Cl2(g)+ 1COCl2(g) ΔH0 = -151 kJ

In the following equilibrium, what effect does simultaneously increasing the temperature and the pressure have?

______________________________________________________________________________________________________________________________________________________________________________

Cannot predict. An increase in temperature causes a shift in the equilibrium to the left, while an increase in pressure causes a shift in equilibrium to the right.

+ heat



A. Safety:

1. Hypothesis: What is the effect of a change in concentration of reactants and a change in temperature on equilibrium?

2. Prediction:

3. Gather Data:

The surfaces of the hot plate will be hot enough to cause burns. Use caution. Cobalt(II) chloride hexahydrate is toxic, with an LD50 = 80mg/kg Avoid ingestion (don’t eat or drink it). Wash handsthoroughly with soap and water before leaving lab. Concentrated Hydrochloric acid is extremely corrosive. Avoid contact with eyes, skin, and clothing. Goggles, aprons, and gloves mandatory.

3. Gather Data:

B. Procedure:

3. Add water dropwise to the test tube until a color change occurs. Record color. ______________

2. Add 3 mL (60 drops) of concentrated HCl to the test tube. Record color. _____________



1. With a partner, measure out 2 mL of 0.1 M CoCl2 solution into a test tube. Record initial color. __________

4. Add 2 mL of 0.1 M CoCl2 solution to another test tube. Add concentrated HCl dropwise until the solution turns purple. If the solution turns blue, add water until it turns purple.

3. Gather Data:

B. Procedure:

6. Place the test tube in a hot water bath. Record color. ___________



5. Place the test tube in an ice bath. Record color. ________

4. Analyze Data:

4Cl- (aq) 1CoCl4

2-(aq) + 6H2O (l)1Co(H2O)6

2+ (aq) +


A. The equation for the reversible reaction in this experiment is:

+ heat



4. Analyze Data:

A. Use the equation to explain the colors of the solution in steps 1, 2, and 3 In Step 1, the solution is initially a pink color, because the reaction

arrives at an equilibrium in which the concentration of the pink-colored 1Co(H2O)6

2+(aq) is at a higher concentration than the blue-colored

1CoCl42-

(aq). In Step 2, the addition of HCl increases the concentration of Cl-, shifting the equilibrium to the right to favor the formation of the blue 1CoCl4

2-(aq), so the solution turns blue. In Step 3, the increase in

concentration of water shifts the equilibrium left, re-establishing a newequilbrium where the concentration of 1CoCl4

2-(aq) is higher than it was

orginally, so the purple color shows more of a balance of pink and blue.



4. Analyze Data:

B. Explain how the equilibrium shifts when heat energy is added or removed.

In Step 5, since heat acts like a product in the exothermic reverse reaction, removing heat by lowering the temperature causes the equilibrium to shift to the left, increasing the rate of the reverse reaction and causing the solution to turn pink. In Step 6, since heat acts like a reactant in the endothermic forward reaction, adding heat by increasing the temperature causes the equilibrium to shift to the right, increasing the rate of the forward reaction and causing the solution to turn blue.

5. Draw Conclusions:


IV. Using Equilibrium Constants -when a ________ has a _____ ___, the __________ _______ contains _____ ________ than ________ at __________

reaction largeKeq equilibrium mixture

more productsreactants equilibrium

-when a ________ has a _____ ___, the __________ _______ contains _____ ________ than ________ at __________

reaction smallKeq equilibrium mixture

more reactantsproducts equilibrium

A. Calculating Equilibrium Concentrations -__________ ________ can also be used to ________ the __________ ____________ of any ________ in the _______

equilibriumconstants

calculateequilibriumconcentration

substancereaction

Keq = [CH4]1

[CO]1

[H2O]1

[H2]3

3.933 = [CH4]1

[0.850]1

[0.286]1

[1.333]3

= 27.7 M

3H2 (g) 1CH4(g)1CO(g) + 1H2O (g)+0.850 M 1.333 M ? M 0.286 M


IV. Using Equilibrium Constants

A. Calculating Equilibrium Concentrations

At 1200 K, the Keq for the following reaction equals 3.933. What is the concentration of the methane produced, if [CO] = 0.850 M, [H2] = 1.333 M, and [H2O] = 0.286 M?

[CH4]

Keq = [H2]2 [S2]1

[H2S]2

2.27 x 10-3 = [H2]2 [0.0540]1

[0.184]2

= 0.0377 M

2H2S (g) 2H2(g) 1S2 (g)+0.184 M ? M 0.0540 M




At 1405 K, the Keq for the following reaction equals 2.27 x 10-3. What is the concentration of the Hydrogen gas produced, if [S2] = 0.0540 M, and [H2S] = 0.184 M?

[H2]

Keq =[CO]1

[CH3OH]1

[H2]2

10.5 =[CO]1

[1.32]1

[0.933]2

= 0.144 M

2H2 (g)1CO(g) + 1CH3OH (g)

? M 0.933 M 1.32 M




If Keq for the following reaction equals 10.5, what is the equilibrium concentration of Carbon monoxide, if [H2] = 0.933 M, and [CH3OH] = 1.32 M?

[CO]



A. Calculating Equilibrium Concentrations from Initial Concentrations Using ICE (Initial, Change, Equilibrium)

If the Keq for the following reaction equals 64.0, what are the equilibrium concentrations of I2, H2, and HI, if [I2]0 = 0.200 M, [H2]0 = 0.200 M and [HI] = 0.000 M?

1H2 (g) 2HI(g)1I2(g) +? M ? M ? M

[H2] [I2] [HI]

Initial

Change

Equilibrium

0.200 0.200 0.000

-1x -1x +2x

0.200 - 1x 0.200 - 1x 2x




1H2 (g) 2HI(g)1I2(g) +

[H2] [I2] [HI]

Initial

Change

Equilibrium

0.200 0.200 0.000

-1x -1x +2x

0.200 - 1x 0.200 - 1x 2x

Keq =[I2]1

[HI]2

[H2]1

64.0 =[0.200 – 1x]1

[2x]2

[0.200 – 1x]1

8.00 =[0.200 – 1x]1

[2x] x = 0.160

[H2] = 0.040 M

[I2] = 0.040 M

[HI] = 0.320 M




If the Keq for the following reaction equals 16.0, what are the equilibrium concentrations of PCl3, Cl2, and PCl5, if [PCl5]0 = 1.00 M?

1Cl2 (g)1PCl5(g) 1PCl3(g) +? M ? M? M

[PCl3] [Cl2] [PCl5]

Initial

Change

Equilibrium

0.00 0.00 1.00

+1x +1x -1x

0.00 + 1x 0.00 + 1x 1.00 – 1x




[PCl3] [Cl2] [PCl5]

Initial

Change

Equilibrium

0.00 0.00 1.00

+1x +1x -1x

0.00 + 1x 0.00 + 1x 1.00 – 1x

Keq = [PCl3]1

[PCl5]1

[Cl2]1

16.0 = [x]1

[1.00 - x]1

[x]1

x2 = 16.0 – 16.0x

x2 + 16.0x - 16.0 = 0

ax2 + bx + c = 0

x = -b ± √b2 – 4ac

2a

1Cl2 (g)1PCl5(g) 1PCl3(g) +




x = -b ± √b2 – 4ac

2a

x = -16.0 ± √(-16.0)2 – 4(1.00)(-16.0)

2(1.00)

x = 0.950 (but not -17.0)[PCl3] = 0.950 M

[Cl2] = 0.950 M

[PCl5] = 0.05 M

x2 + 16.0x - 16.0 = 0

ax2 + bx + c = 0




If the Keq for the following reaction equals 0.680, what are the equilibrium concentrations of COCl2, CO, and Cl2, if [CO]0 = 0.500 M and [Cl2]0 = 1.00 M?

1Cl2 (g)1CO (g) 1COCl2 (g)+? M? M ? M

[COCl2] [CO] [Cl2]

Initial

Change

Equilibrium

0.00 0.500 1.00

+1x -1x -1x

0.00 + 1x 0.500 - 1x 1.00 – 1x




[COCl2] [CO] [Cl2]

Initial

Change

Equilibrium

0.00 0.500 1.00

+1x -1x -1x

0.00 + 1x 0.500 - 1x 1.00 – 1x

Keq =[Cl2]1[CO]1

[COCl2]1

0.680 =[1.00 - x]1[0.500 - x]1

[x]1

x = 0.340 - 0.340x – 0.680x + x2

0.680x2 - 2.02x + 0.340 = 0

ax2 + bx + c = 0

1Cl2 (g)1CO (g) 1COCl2 (g)+




x = -b ± √b2 – 4ac

2a

x = 2.02 ± √(-2.02)2 – 4(0.680)(0.340)

2(0.680)

x = 2.79 or 0.176[COCl2] = 0.176 M

[CO] = 0.324 M

[Cl2] = 0.82 M

x = 0.176

0.680x2 - 2.02x + 0.340 = 0

ax2 + bx + c = 0




If the Keq for the following reaction equals 36.0, what are the equilibrium concentrations of H2, Br2, and HBr, if [H2]0 = 0.250 M and [Br2]0 = 0.250 M?

1Br2 (g) 2HBr(g)1H2(g) +? M ? M ? M

[H2] [Br2] [HBr]

Initial

Change

Equilibrium

0.250 0.250 0.000

-1x -1x +2x

0.250 - 1x 0.250 - 1x 0.000 + 2x




1Br2 (g) 2HBr (g)1H2 (g) +

Keq =[H2]1

[HBr]2

[Br2]1

36.0 =[0.250 - x]1

[2x]2

[0.250 - x]1

[H2] [Br2] [HBr]

Initial

Change

Equilibrium

0.250 0.250 0.000

-1x -1x +2x

0.250 - 1x 0.250 - 1x 0.000 + 2x

6.00 =(0.250 – x)

2x

x = 0.188




32x2 - 18.0x + 2.25 = 0

ax2 + bx + c = 0

x = -b ± √b2 – 4ac

2a

x = 18.0 ± √(-18.0)2 – 4(32)(2.25)

2(32)

x = 0.375 or 0.188[H2] = 0.062 M

[Br2] = 0.062 M

[HBr] = 0.376 M




If the Keq for the following reaction equals 20.0, what are the equilibrium concentrations of H2, Cl2, and HCl, if [H2]0 = 1.00 M and [Cl2]0 = 2.00 M?

1Cl2 (g) 2HCl(g)1H2(g) +? M ? M ? M

[H2] [Cl2] [HCl]

Initial

Change

Equilibrium

1.00 2.00 0.00

-1x -1x +2x

1.00 - 1x 2.00 - 1x 0.00 + 2x




1Cl2 (g) 2HCl (g)1H2 (g) +

Keq =[H2]1

[HCl]2

[Cl2]1

20.0 =[1.00 - x]1

[2x]2

[2.00 - x]1

[H2] [Cl2] [HCl]

Initial

Change

Equilibrium

1.00 2.00 0.00

-1x -1x +2x

1.00 - 1x 2.00 - 1x 0.00 + 2x

4.00x2 = 40.0 – 60.0x + 20.0x2

16.0x2 - 60.0x + 40.0 = 0

ax2 + bx + c = 0




16.0x2 - 60.0x + 40.0 = 0

ax2 + bx + c = 0

x = -b ± √b2 – 4ac

2a

x = 60.0 ± √(-60.0)2 – 4(16.0)(40.0)

2(16.0)

x = 2.88 or 0.869[H2] = 0.13 M

[Cl2] = 1.13 M

[HCl] = 1.74 M


V. Solubility Equilibria -like a few _________ _________ that go to _________, upon __________, some ______ __________ _________ completely into _____

chemical reactionscompletion dissolving

ionic compounds dissociateions

1NaCl (s) 1Na+ (aq) + 1Cl-

(aq)

-some _____ __________, however, are only ________ _______, and quickly reach a ________ __________

1BaSO4 (s) 1Ba2+ (aq) + 1SO4

2- (aq)

ionic compoundssparingly soluble

solubility equilibrium

Keq = [SO42-]1[Ba2+

]1

[BaSO4]1


V. Solubility Equilibria -in the __________ _______ __________, ______ ______ is a _____, so the _______ is _______, and can be combined with the ___ value to form the ________ _______ _______

Keq = [SO42-]1[Ba2+

]1[BaSO4]1x

equilibrium constant expressionBarium sulfate solid [BaSO4]

constantKeq solubility productconstant

Ksp = [SO42-]1[Ba2+

]1

Write the solubility constant expression for the following solubility equilibrium:

1Mg(OH)2 (s) 1Mg2+ (aq) + 2OH-

(aq)

Ksp = [OH-]2[Mg2+ ]1


V. Solubility Equilibria

A. Calculating Solubilities from Solubility Product Constants

What is the solubility, in M, of Silver iodide at 298 K?

1AgI (s) 1Ag+ (aq) + 1I-

(aq)

Ksp = [I-]1[Ag+ ]1

8.5 x 10-17 = [I-]1[Ag+ ]1

=8.5 x 10-17√ s2√

s = 9.2 x 10-9 M

8.5 x 10-17 = s2



A. Calculating Solubilities from Solubility Product Constants

What is the solubility, in M, of Copper(II) carbonate at 298 K?

1CuCO3 (s) 1Cu2+ (aq) + 1CO3

2- (aq)

Ksp = [CO32-]1[Cu2+

]1

2.5 x 10-10 = [CO32-]1[Cu2+

]1

=2.5 x 10-10√ s2√

s = 1.6 x 10-5 M

2.5 x 10-10 = s2



B. Calculating Ion Concentration from Ksp

What is [OH-] at 298 K in a saturated solution of Mg(OH)2 at equilibrium? 1Mg(OH)2 (s) 1Mg2+

(aq) + 2OH- (aq)

Ksp = [OH-]2[Mg2+ ]1

let x = [Mg2+ ]

so 2x = [OH-]

5.6 x 10-12 = (x)(2x)2

5.6 x 10-12 = 4x3

1.4 x 10-12 = x3

√1.4 x 10-123 = x

1.1 x 10-4 = x2.2 x 10-4 = 2x = [OH-]




What is [Ag+] at 298 K in a saturated solution of AgBr at equilibrium?

1AgBr (s) 1Ag+ (aq) + 1Br-

(aq)

Ksp = [Br-]1[Ag+ ]1

let x = [Ag+ ]

so x = [Br-]

5.4 x 10-13 = x2

√5.4 x 10-13 = x

7.3 x 10-7 = x = [Ag+]




What is [F-] at 298 K in a saturated solution of CaF2 at equilibrium?

1CaF2 (s) 1Ca2+ (aq) + 2F-

(aq)

Ksp = [F-]2[Ca2+ ]1

3.5 x 10-11 = (x)(2x)2

3.5 x 10-11 = 4x3

8.8 x 10-12 = x3

√8.8 x 10-123 = x

2.1 x 10-4 = x

4.2 x 10-4 = 2x = [F-]



C. Predicting Precipitates -besides being used to calculate the _________ of an _____ _________ and the ___________ of ____ in a _________ _______, ___ values can be used to _______ if a _________ will form if ___ _____ __________ are mixed

solubility ionic compoundconcentration ions

saturated solution Ksp

predict precipitatetwo ionic compounds

Predict whether PbCl2 will form as a precipitate if 100 mL of 0.0100 M NaCl is added to 100 mL of 0.0200 M Pb(NO3)2:

-the ____________ of the ______ ________ allow you to _______ the ____________ of ____ and ___ ions in the _____ _________, which when _________ together, determine the ___ _______, or ___

concentrations initialsolutions calculateconcentrations Pb2+ Cl-

mixed solutionsmultiplied

ion product Qsp



C. Predicting Precipitates

Predict whether PbCl2 will form as a precipitate if 100 mL of 0.0100 M NaCl is added to 100 mL of 0.0200 M Pb(NO3)2:

1PbCl2 (s) 1Pb2+ (aq) + 2Cl-

(aq)

Qsp = [Cl-]2[Pb2+]1

[Pb2+] = 0.0200 M

2

= 0.0100 M

[Cl-] = 0.0100 M

2

= 0.00500 M

Qsp = [0.00500]2[0.0100]1

Qsp = 2.50 x 10-7 < 1.7 x 10-5 = Ksp



C. Predicting Precipitates -if the ___ is ___ the ___, the _______ is __________, and a _________ ____ ___ ____, and if the ___ is ___ the ___, the _______ is _________ and ___ ______ will occur, but if ___ is ___ the ___, a __________ will form, reducing the ___ ___________ until ___ ___ ___, and the system arrives at __________ and the _______ becomes ________

Qsp < Ksp

solution unsaturatedprecipitate will not form

Qsp = Ksp solutionsaturated no change

Qsp > Ksp

precipitateion concentration Qsp = Ksp

equilibrium solutionsaturated

Qsp = 2.50 x 10-7 < 1.7 x 10-5 = Ksp

No precipitate should form




Predict whether Ag2SO4 will form as a precipitate if 500 mL of 0.010 M AgNO3 is added to 500 mL of 0.25 M K2SO4:

1Ag2SO4 (s) 2Ag+ (aq) + 1SO4

2- (aq)

Qsp = [SO42-]1[Ag+]2

[Ag+] = 0.010 M

2

= 0.0050 M

[SO42-] = 0.25 M

2

= 0.012 M

Qsp = [0.012]1[0.0050]2

Qsp = 3.0 x 10-7 < 1.2 x 10-5 = Ksp

No precipitate should form




Predict whether a precipitate will form if 200 mL of 0.20 M MgCl2 is added to 200 mL of 0.0025 M NaOH:

1Mg(OH)2 (s) 1Mg2+ (aq) + 2OH-

(aq)

Qsp = [OH-]2[Mg2+]1

[Mg2+] = 0.20 M

2

= 0.10 M

[OH-] = 0.0025 M

2

= 0.0012 M

Qsp = [0.0012]2[0.10]1

Qsp = 1.4 x 10-7 > 5.6 x 10-12 = Ksp

A precipitate of Mg(OH)2 should form



D. Common Ion Effect -the ________ of _______ in _____ is ________ mol/L, which means that you can ________ ________ of _______ in ____ L of _____ _____, but _________ of _______ will ____ _______ in ____ L of a ______ solution of _______, because of the ________ ___ ______

solubility PbCrO4 water4.8 x 10-7

dissolve 4.8 x 10-7

PbCrO4 1.00 pure water4.8 x 10-7 PbCrO4 not

dissolve 1.00 0.10 MK2CrO4

common ion effect

1PbCrO4 (s) 1Pb2+ (aq) + 1CrO4

2- (aq)

Ksp = [CrO42-]1[Pb2+]1 = 2.3 x 10-13

-since the _______ of the ____________ of both ____ is _____ to a _______, (the _________ _______ _______), if _______ goes __, _____ must go _____

product concentrationsions equal constant

solubility product constant[CrO4

2-] up [Pb2+] down



D. Common Ion Effect -adding a _______ to an __________ that contains a ________ ___ _______ the ________ of a _________ containing that ___, or, according to _____________ ________, stresses the __________ and causes the _______ to _____ the __________ in the _______ that _______ the ______

solution equilibriumcommon ion lowers

solubility substanceion

Le Châtelier’s Principleequilibrium

system shift equilibriumdirection relieves stress

1PbCrO4 (s) 1Pb2+ (aq) + 1CrO4

2- (aq)


A. Safety:

1. Hypothesis: How Do Solubility Product Constants Compare?

2. Prediction:

3. Gather Data:

Silver nitrate stains skin and clothing and is highly toxic, with an LD50 = 50mg/kg Avoid ingestion (don’t eat or drink it). Wash hands thoroughly with soap and water before leaving lab. Goggles mandatory.



B. Procedure:

1. Using a pipette, place 10 drops of AgNO3 solution into test well A1 of a 20-well microplate. Place 10 more drops of the same solution in test well A2.

3. Gather Data:

B. Procedure:

3. To test well A2 only, add 10 drops of Na2S solution. Record observations______________________________

2. Add 10 drops of NaCl solution to both test well A1 and test well A2. Record observations___________________



4. Compare the contents of test wells A1 and A2. Record observations_____________________________________

4. Analyze Data:

A. Write the complete thermochemical equation for the reaction that occurred in Step 2.



B. Write the net ionic equation for the reaction in Step 2.

C. Write the equation for the solubility equilibrium that was established in test wells A1 and A2 during Step 2.

D. Write the solubility constant expression for the equilibrium established in test wells A1 and A2 during Step 2.

E. Write the equation for the solubility equilibrium that was established in test well A2 during Step 4.

4. Analyze Data:

F. Match the chemical formula of each precipitate with its color.



G. Compare the two Ksp values for the two precipitates. Infer which is the more soluble.

H. Use Le Châtelier’s Principle to explain how the addition of Na2S in Step 4 affected the equilibrium in test well A2.

4. Analyze Data:



I. Calculate the molar solubilities of both precipitates in the experiment. Which of the precipitates is more soluble?

4. Analyze Data:

A. Write the complete thermochemical equation for the reaction that occurred in Step 2.



B. Write the net ionic equation for the reaction in Step 2.

C. Write the equation for the solubility equilibrium that was established in test wells A1 and A2 during Step 2.

D. Write the solubility constant expression for the equilibrium established in test wells A1 and A2 during Step 2.

E. Write the equation for the solubility equilibrium that was established in test well A2 during Step 4.

AgNO3 (aq) + NaCl (aq) → AgCl (s) + NaNO3 (aq)

Ag+ (aq) + Cl- (aq) AgCl (s)

Ksp = [Ag+]1 [Cl-]1 = 1.8 x 10-10

1AgCl (s) 1Ag+ (aq) + 1Cl- (aq)

1Ag2S (s) 2Ag+ (aq) + S2- (aq)

4. Analyze Data:

F. Match the chemical formula of each precipitate with its color.



G. Compare the two Ksp values for the two precipitates. Infer which is the more soluble.

H. Use Le Châtelier’s Principle to explain how the addition of Na2S in Step 4 affected the equilibrium in test well A2.

AgCl (s) is white; Ag2S (s) is black

Ksp for AgCl (s) = 1.8 x 10-10

Ksp for Ag2S (s) = 8 x 10-48; AgCl is more soluble

The addition of S2- to the equilibrium removes Ag+ from the equilibrium, causing the system to shift to the right in favor of the formation of Ag+, so at the same time the black precipitate Ag2S is forming, the white AgCl is dissolving to relieve the stress on the equilibrium.

1AgCl (s) 1Ag+ (aq) + 1Cl- (aq)

4. Analyze Data:



I. Calculate the molar solubilities of both precipitates in the experiment. Which of the precipitates is more soluble?

1AgCl (s) 1Ag+ (aq) + 1Cl-

(aq)

Ksp = [Cl-]1[Ag+ ]1

1.8 x 10-10 = [Cl-]1[Ag+ ]1

=1.8 x 10-10√ s2√

s = 1.3 x 10-5 M

1.8 x 10-10 = s2

1Ag2S (s) 2Ag+ (aq) + 1S2-

(aq)

Ksp = [S2-]1[Ag+ ]2

8 x 10-48 = [S2-]1[Ag+ ]2

=

s = 1 x 10-16 M

8 x 10-48 = 4s3

2 x 10-48√3 s3√3

AgCl is more soluble

Chemical Equilibrium I. A State of Dynamic Balance 1N 2 (g) 3H 2(g) +2NH 3(g) ΔG 0 = -33.1 kJ The...

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Transcript of Chemical Equilibrium I. A State of Dynamic Balance 1N 2 (g) 3H 2(g) +2NH 3(g) ΔG 0 = -33.1 kJ The...