chapter 16:

CHAPTER 16:

PRINCIPLES OF REACTIVITY:

Chemical Equilibria

16.0 OBJECTIVES

State the characteristics of a system in dynamic equilibrium as applied to reversible reactions.

Calculate the value of and interpret the meaning of the reaction quotient based on either partial pressures or concentrations.

Understand the factors that will change the values of the reaction quotient.

Use LeChatelier’s Principle to describe changes in reactions based on stresses to the equilibrium system.

Homework

#1 - 9, 11, 13, 15, 37, 39, 45 Basic Equilibrium Calculations

#2 - 17, 19, 21, 23, 41 ICE Problems

#3 - 27, 29, 31, 43, 51, 53, 55, 57 More ICE and Conceptual Equilibrium

#4 - 67, 33, 35, 47, 49 Le’Chatelier’s Principle

No movement observed.

In a reversible chemical change, a static equilibrium is never established.

Static Equilibrium

Dynamic EquilibriumMovement IS observed

Rate of water in-flow = Rate of water out-flow The level of water is constant within the bucket.

In a reversible chemical change:a static equilibrium is never established, but Dynamic Equilibrium is!

Examples of EquilibriaExamples of EquilibriaPhase changes such as HPhase changes such as H22O(g)↔HO(g)↔H22O(liq)O(liq)

ExampleExamples of s of

Chemical Chemical EquilibriaEquilibria

Formation of stalactites and stalagmitesFormation of stalactites and stalagmites

CaCOCaCO3(s) 3(s) +H+H22OO(liq) (liq) + CO+ CO2(g)2(g)↔↔ CaCa2+ 2+ (aq) (aq) + 2 HCO+ 2 HCO3(aq)3(aq)

16.1 NATURE OF THE EQUILIBRIUM STATE

A. Requirements Reaction that forms products to a certain point

(equilibrium!), then reforms reactants

(NOT ALL RXNS will go to completion and only form products)

Equ. will not occur as soon as the reversible reaction begins

Like all good things, it takes time!


B. Characteristics of a Specific System Chem Eqn has Double Headed Arrow ↔↔ Beginning of Rxn Form lots of products (rate laws)

Moments Later Forming both products and reactants (equilibrium)

**Not Necessarily at Equal Rates yet!

Dynamic Equilibrium- “two opposing processes occur at exactly the same raterate”

Reactants Products = Products Reactants


C. Graphs


N2O4 2NO2 Equilibrium Demo Decomposition Reaction

NO2 = Smog (dark brown color)

N2O4 = Propellant (light brown/clear)

**Observe and Write down some observations!

Nitrogen Nitrogen Dioxide Dioxide

EquilibriumEquilibrium

NN22OO44(g) (g) 2 NO2 NO22(g)(g)

Nitrogen Nitrogen Dioxide Dioxide

EquilibriumEquilibrium

NN22OO44(g) (g) 2 NO2 NO22(g)(g)

16.2 THE REACTION QUOTIENT AND EQUILIBRIUM CONSTANT

A. For aA(g) + bB(g) cC(g) + dD(g)

Kc = [C]c [D]d … [A]a [B]b ...

For a given reaction at equilibrium, the ratio of Products over Reactants raised to their coefficients is CONSTANT. That RATIO will always equal K no matter the conc.

REACTION QUOTIENT (“Q”) at equilibrium

EQUILIBRIUM CONSTANT


B. Example 16.1 Write the Kc for a. H2(g) + I2(g) 2 HI (g)

b. 2 SO2(g) + O2(g) 2SO3(g)


C. Kc is

independent of Initial Concentrations

dependent on Temperature Reaction/Stoichiometry

Kc = [C]c [D]d … [A]a [B]b ...

aA(g) + bB(g) cC(g) + dD(g)


D. No term for pure liquids, solvents, or solids

Their concentrations do not change during the reaction

Kc = [C]c [D]d _ [A]a [B]b [H2O]n

aA + bB(s) + nH2O(l) cC + dD


Kc = [C]c [D]d … [A]a [B]b ...

Meaning of K? Usefulness of K?•Dictates weather a reaction if product or reactant favored

A+B A+B

C+DC+D


E. Example 16.2 Write the equilibrium constant expression (mass action expression) for

a. CuO(s) + H2(g) Cu(s) + H2O(g)

b. C(s) + H2O(g) CO(g) + H2(g) c. I2(s) I2(g)

In pairs

16.3 DETERMINING AN EQUILIBRIUM CONSTANT

A. Determine the value of Kc when all equilibrium concentrations are known: Products over Reactants raised to their

coefficients

16.3 DETERMINING AN EQUILIBRIUM CONSTANT 1. Example 16.3 For NH4Cl(s) NH3(g) + HCl(g) at 500oC, at

equilibrium there are 2.0 moles of ammonia, 2.0 moles of hydrochloric acid and 1.0 mole of ammonium chloride present in a 5.0L container. Calculate the equilibrium constant of the system at this temperature.

16.3 DETERMINING AN EQUILIBRIUM CONSTANT 1. Example 16.3 For NH4Cl(s) NH3(g) + HCl(g) at 500oC, at

equilibrium there are 2.0 moles of ammonia, 2.0 moles of hydrochloric acid and 1.0 mole of ammonium chloride present in a 5.0L container. Calculate the equilibrium constant of the system at this temperature.

Kc = [NH3]1 [HCl]1

Kc = [2.0/5.0] [2.0/5.0]

Kc = 0.16

*note: units never given with K!

2. Example 16.4 For the system 2SO3(g) 2SO2(g) + O2(g) at a given temperature the equilibrium concentrations are [SO2] = [O2] = 0.10M and [SO3] = 0.20M. Calculate the Kc at this temperature.


In pairs

2. Example 16.4 For the system 2SO3(g) 2SO2(g) + O2(g) at a given temperature the equilibrium concentrations are [SO2] = [O2] = 0.10M and [SO3] = 0.20M. Calculate the Kc at this temperature.


Kc = [SO2]2 [O2]1

[SO3]2

Kc = [0.10] [0.10]

[0.20]2

Kc = 0.25

1. Example 16.5 For 2HI(g) I2(g) + H2(g). Starting with only HI at a concentration of 0.100M, it is found that at 520oC the equilibrium concentration of H2 is 0.010M. Calculate the equilibrium concentrations of I2 and HI and the Kc for this reaction.


B. Determine the value of Kc when some equilibrium concentrations are missing:

Equation

Initial(M)

Change(M)

Equilibrium(M)

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

0.100 0 0

0.010

- + +

Kc =


Equation

Initial(M) 0.100 0 0

Change(M) - 0.020 +0.010 + 0.010

Equilibrium(M) 0.100-0.02 0.010 0.010 =(0.080)

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

Kc = [I2]1 [H2]1 = [0.010]1 [0.010]1 = [HI]2 [0.080]2

Ex. 16.5

0.016

2. Example 16.6 Starting with the same system as above and with the original concentration of HI as 0.100M, it is found that the equilibrium concentration of HI in 0.074M at a different temperature. Determine the equilibrium concentrations of hydrogen and iodine and the Kc of the reaction.


Equation

Initial(M)

Change(M)

Equilibrium(M)

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

In pairs

2. Example 16.6 Starting with the same system as above and with the original concentration of HI as 0.100M, it is found that the equilibrium concentration of HI in 0.074M at a different temperature. Determine the equilibrium concentrations of hydrogen and iodine and the Kc of the reaction.


Equation

Initial(M)

Change(M)

Equilibrium(M)

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

Kc = [I2]1 [H2]1 =

[HI]2

0.100

0.074

+ X + X

0 0

-2X

X=0.013 X=0.013

0.100 – 2X = 0.074X = 0.013

QUIZ!!!!!!!

0.031=(0.013)(0.013)/(0.074)2

16.4 USING EQUILIBRIUM CONSTANTS IN CALCULATIONS

A. Predicting the direction of shift of a reaction to reach equilibrium 1. For aA(g) + bB(g) cC(g) + dD(g)

Kc = [C]c [D]d …_ [A]a [B]b ...

Q = [C]c [D]d …_ [A]a [B]b ...

In Text Book, Q is introduced in 16.2, pg 663

REACTION QUOTIENT (Q)EQUILIBRIUM CONSTANT (Kc ) is essentially REACTION QUOTIENT (Q) at equilibrium

16.4 USING EQUILIBRIUM CONSTANTS IN CALCULATIONS 2. Relationship between Kc and Q

Q is the ratio of Products/Reactants at ANY given point in time, not necessarily at equilibrium

Q values can change as reaction proceeds

Kc is the equilibrium constant which is established for EQUILIBRIUM CONDITIONS ONLY!

Pg. 665


Pg. 665


Kc = [C]c [D]d …_ [A]a [B]b ...

Q = [C]c [D]d …_ [A]a [B]b ...

16.4 USING EQUILIBRIUM CONSTANTS IN CALCULATIONS 3. Example 16.7 N2(g) + 3H2(g) 2NH3(g) Kc= 5 x 108 at 25oC

If the original concentrations are [N2] = [H2] = 2.0M, which way will the reaction run?

Q =

Compare Q to Kc:


3. Example 16.7 N2(g) + 3H2(g) 2NH3(g) Kc= 5 x 108 at 25oC

If the original concentrations are [N2] = [H2] = 2.0M, which way will the reaction run?

At equilibrium, Q = Kc = 5 x 108

Q = [NH3]2 _ [N2]1 [H2]3

Q = [0]2 _ [2.0]1 [2.0]3

Q< Kc, products favored

16.4 USING EQUILIBRIUM CONSTANTS IN CALCULATIONS 4. Example 16.8 For the system N2O4(g) 2NO2(g) Kc = 0.36 at 100oC:

Predict the direction the reaction will shift to reach equilibrium if the original concentrations are:

a. 0.20 moles of N2O4 in a 4.0L container

b. 0.20 moles of N2O4 and 0.20 moles of NO2 in a 4.0L container

c. 1.00M NO2 and N2O4

a. b. c.

In pairs

16.4 USING EQUILIBRIUM CONSTANTS IN CALCULATIONS 4. Example 16.8 For the system N2O4(g) 2NO2(g) Kc = 0.36 at 100oC:

Predict the direction the reaction will shift to reach equilibrium if the original concentrations are:

a. 0.20 moles of N2O4 in a 4.0L container

b. 0.20 moles of N2O4 and 0.20 moles of NO2 in a 4.0L container

c. 1.00M NO2 and N2O4

a. b. c.

Q<K, products favored Q>K, reactants favoredQ<K, products favored

Q = [0]2 / [.2/4]Q=0

Q = [NO2]2 / [N2O4], Kc = 0.36

Q = [.2/4]2 / [.2/4]Q=0.05

Q = [1.0]2 / 1.0]Q=1

16.4 USING EQUILIBRIUM CONSTANTS IN CALCULATIONS B. Determine the equilibrium concentration of one reactant or product: Example 16.9 For N2(g) + O2(g) 2NO(g) Kc = 1.0 x 10-30 at 25oC

Calculate the equilibrium concentration of NO(g) if at equilibrium the concentration of N2(g) is 0.04M and that of O2 is 0.01M.

16.4 USING EQUILIBRIUM CONSTANTS IN CALCULATIONS B. Determine the equilibrium concentration of one reactant or product: Example 16.9 For N2(g) + O2(g) 2NO(g) Kc = 1.0 x 10-30 at 25oC

Calculate the equilibrium concentration of NO(g) if at equilibrium the concentration of N2(g) is 0.04M and that of O2 is 0.01M.

Kc =1x10-30 = [NO]2 / [N2][O2]

1x10-30 = [NO]2 / [0.04][0.01]

[NO] = 2x10-17

16.4 USING EQUILIBRIUM CONSTANTS IN CALCULATIONS C. Determine the equilibrium concentrations of all reactants and products from

original concentrations and Kc.

Example 16.10 For CO2(g) + H2(g) CO(g) + H2O(g) K c = 0.64 at 900K. Starting with both reactants at a concentration of 0.100M, what are the equilibrium concentrations of all species?

16.4 USING EQUILIBRIUM CONSTANTS IN CALCULATIONS C. Determine the equilibrium concentrations of all reactants and products from

original concentrations and Kc.

Example 16.10 For CO2(g) + H2(g) CO(g) + H2O(g) K c = 0.64 at 900K. Starting with both reactants at a concentration of 0.100M, what are the equilibrium concentrations of all species?

CO2(g) + H2(g) CO(g) + H2O(g)

I 0.1 0.1 0 0

C -x -x +x +x

E 0.1-x 0.1-x x x

K c = 0.64 = [x][x]/[0.1-x][0.1-x]Solve for x: x =0.044M

X = 0.044 M [CO] = [H2O]

16.4 USING EQUILIBRIUM CONSTANTS IN CALCULATIONS Example 16.11 For the same system as above, calculate the equilibrium

concentrations of all species if the reaction system is started with [CO2] = 0.100M and [H2] = 0.200M. CO2(g) + H2(g) CO(g) + H2O(g) K c = 0.64 at 900K

Quad eqn example!

2) Solve for x

3) Use x to find equil concs

1) Set up Kc

16.4 USING EQUILIBRIUM CONSTANTS IN CALCULATIONS Example 16.12 For N2O4(g) 2NO2(g) Kc = 0.36 at 100oC. Starting with a

concentration of 0.100 moles/L for N2O4, what are the equilibrium concentrations of both species?

In pairs

2) Solve for x

3) Use x to find equil concs

1) Set up Kc

16.5 MORE ABOUT BALANCED EQUATIONS AND EQUILIBRIUM CONSTANTS

A. KP

Ratio of Partial Pressures of reactants/products at Equilibrium

KP = PCc PD

d …_ PA

a PBb ...


C. Example 16.13 For N2(g) + 3H2(g) 2NH3(g) Kc = 9.5 at 300oC.

Calculate the KP at this same temperature.

Kp = Kc (RT) n gas


D. Equilibrium constants and different forms of the equationFor N2O4(g) 2NO2(g) Kc = 0.36 at 100oC

1. Kc of 2NO2(g) N2O4(g) Reverse Equation: 1/K

Kc forward rxn = Kc reverse rxn =


2. Kc of 4NO2(g) 2N2O4(g) Multiply by coefficent: Kcoefficient

3. Kc of 1/2N2O4(g) NO2(g) See above

4. Summation of Equations Rule Multiply K values for all equations


5. Example 16.14

Given: 3/2O2(g) O3(g) K = 2.5 x 10-29 State the K for:

a. 3O2(g) 2O3(g) b. 2O3(g) 3O2(g)


6. Example 16.15 Given at 500K: H2(g) + Br2(g) 2HBr(g) Kp = 7.9 x 1011

H2(g) 2H(g) Kp = 4.8 x 10-41

Br2(g) 2Br(g) Kp = 2.2 x 10-15

Calculate the Kp for H(g) + Br(g) HBr(g) at the same temperature

HW # 1, 2 and 3 Due!!!QUIZ!!!!

16.6 DISTURBING A CHEMICAL EQUILIBRIUM: Le Chatelier's Principle

A. Statement

“When a system is stressed (changing conc., temp, pressure, or volume), the system will respond by attaining new equilibrium conditions

that counteract the change”


B. Adding of removing a reactant or product Equilibrium shifts in the direction opposite of the

stress (ex. add products shifts to reactants)

Add Reactant Shifts toward products Remove Reactant Shifts toward reactants Add Product Shifts toward reactants Remove Product Shifts toward products


C. Example 16.16 For 2HI(g) H2(g) + I2(g) Kc = 0.016 at 520oC Starting at the following concentrations at equilibrium [H2] = [I2] = 0.010M and [HI] = 0.080M.. After reaching the above equilibrium, enough HI is added to raise its concentration of 0.096M. What are the equilibrium concentrations of all species?


D. Change in amounts of solid, pure liquids, or solvents

Has NO effect on equilibrium


E. Change in volume or pressure (for gases) Pressure increased (Volume dec.) Equ. shifts in

the direction producing the smaller number of moles of gas

Pressure is decreased (Volume inc.) Equ. Shifts in the direction producing the larger number of moles of gas

NO CHANGE in moles no effect on an equilibrium


F. Change in temperature Increase Temp Equ. shifts in the direction of the

endothermic rxn

Decrease Temp Equ. shifts in the direction of the exothermic rxn


G. Adding a catalyst Has NO effect on equilibrium


H. Example 16.17 What conditions would be most ideal for shifting the following equilibrium as far to the right as possible? N2(g) + 3H2(g) 2NH3(g)

HW#4 Due!!!

END OF CHAPTER 16

http://www.mhhe.com/physsci/chemistry/animations/chang_7e_esp/kim2s2_5.swf

http://www.freezeray.com/flashFiles/theHaberProcess.htm

http://www.freezeray.com/flashFiles/ammoniaConditions.htm

Warm Up Activity,

NO(g) + O2(g) OONO(g) Fast

NO(g) + OONO(g) 2NO2(g) Slow

overall

R rxn = R slow = k2[NO][OONO]

R1,f = R1,r

k1f [NO][O2]= k1r[OONO]

k1f [NO][O2]= k1r[OONO]

k1f

k1r

k2

Ex15.11

Draw a phase diagram and Label the:

a)Critical pointb)Triple Pointc)Gas Phased)Solid Phasee)Liquid Phasef)Gas/liquid Equilibriumg)Liquid/solid Equilibriumh)Gas/solid Equilibrium

Warm Up Activity, Tues. May 13

Label the:a)critical pointb)triple pointc)Gas phased)Solid phasee)Liquid phasef)Gas/liquid equilibriumg)liquid/solid equilibriumh)Gas/solid equilibrium h

gf

b

a

Kinetics and equilibrium

For simple, first order reversible reaction: A ↔B

Begin with only reactants

Begin with only products

Kinetics, equilibrium, and the reaction coordinate diagram

Since the forward reaction activation energy is less than the reverse reaction activation energy (Ea, fwd < Ea, rev ), the reverse reaction will be slower than the forward reaction when the concentrations of A and B are equal.

In order for this reaction to reach equilibrium, the concentration of B will have to increase beyond that of A up to the point where the rates of the forward and reverse reactions become equal. Qualitatively, since Ea, fwd < Ea, rev in this case, Keq > 1.

We can conclude that the net change in chemical potential energy is related to the equilibrium constant (Keq) for the reaction, whereas the magnitude of the activation energy must be related to the rate constants for the reaction.

For simple, first order reversible reaction: A ↔B

•can derive the relationship between rate constants (kforward and kreverse ) and the equilibrium constant (Keq)

Rate laws for forward and reverse reactions:•Rforward = kfor[A] Rreverse =krev[B]

At equilibrium, •Rforward = Rreverse

kfor[A]eq = krev [B]eq

rearranged: kfor/krev = [B]eq/[A]eq

since [B]eq/[A]eq = Keq

kfor/krev = [B]eq/[A]eq = Keq

•we have shown in this case that: Keq = k1/k−1

Relationship between Rate Constant and Equilibrium Constant for a simple, 1st order reaction

Dr. Nahorniak’s Expectations for the remainder of the yearExpectations for class in general:You are expected and required to:•Follow directions•Be in your seat, quiet, with your materials out ready to begin class when the bell rings.•Remain in your seat and working, listening or sitting quietly until the bell rings at the end of the period.•Complete homework on time and in good faith, giving every problem and every assignment a good, true, honest effort, (IN WRITING/ON PAPER!) referring to your notes, the chapter text, your classmates, your instructor, etc. as necessary for guidance.•Pay attention, take notes, and participate in class. •When time is up, STOP working and pass your papers forward or they will not be accepted.•Neatly put your first name, last name, and period on the top of everything you turn in.

In addition:•Only leave the classroom in the case of an emergency. Use the restroom, eat, and drink before or after class.•No headphones, phones, games, etc., •No studying for other courses or doing other work during chemistry class.

Expectations for Lab:At a minimum, you are expected and required to:•Wear your goggles, Wear closed toed shoes, Not eat or drink in lab•Come prepared, having read through the lab and completed any pre-lab assignments BEFORE class begins.•Have your lab notebook, lab handout, and any other necessary materials at your lab table.•Be respectful of all equipment and materials.•Clean up after yourself.In addition:•No fooling around, playing with equipment, or bothering other groups.•Use your time efficiently and don’t be lazy. Lab time is not “free” time.•Put thought and effort into your lab write ups and be legible (especially conclusions!).

I have read and understand “Dr. Nahorniak’s Expectations for the remainder of the year” and agree to meet these expectations to the best of my ability. ______________________________ __________ Name Period

____________________________ _________Signature Date

Announcements May 15, 2014:

•HW 1,2, and 3 will be due when we finish 16.5 and worth 20 points.•Do NOT wait until the night before to do the homework. Start now by doing the problems from the material we cover as we cover it.

•Probable Quiz following 16.4.

•Extra Credit: -Topics Chosen by Wed. May 21 -Proposals due Fri. May 23 -Due June 9 (be ready to present/etc. anytime between June 9 and the final)

•Qual Lab II: How many need to make it up?

•POGIL Due Monday May 19 at START of class!

chapter 16:

Documents