chapter 13.1-13.4 ap chem chemical equilibrium. 2 chapter 13 table of contents copyright © cengage...

Chapter 13.1-13.4AP Chem

Chemical Equilibrium

Chapter 13

Table of Contents

13.1The Equilibrium Condition

13.2The Equilibrium Constant

13.3Equilibrium Expressions Involving Pressures

13.4Heterogeneous Equilibria

13.5Applications of the Equilibrium Constant

13.6 Solving Equilibrium Problems

13.7 Le Châtelier’s Principle

Chapter 13

Table of Contents

• WEEK OUTLOOK• Monday - Notes 13.1-13.4 with problems w/sheet due

Tuesday - should be able to complete in class today.• *Be sure all reports are turned in and made up TODAY!• Tuesday - Notes 13.5-13.7 with emphasis on Le

Chatelier’s Principle emphasized & problems assigned due Wed. - some time in class to complete

• Kaci & Jonathan - library 2nd floor 7:30 with ACT invent.• Wednesday- Lab• Thursday - CAPS - No class• Friday - Good Friday - No school• Tuesday - April 2nd - ACT Testing 11th graders only.• Test probably next Thursday - just over ch. 13 only.

Chapter 13

Table of Contents

• HANDOUTS - Ch. 13 NMSI Equilibrium Packet (Notes with practice problems)

• HANDOUT - Equilibrium Homework Sheet #1

• TURN IN Kinetics Lab - will go over pre-lab questions• HW: Equilibrium w/s #1 should be done today - due

Mon. for grade• HW: Notes packet #1-6 problems due next week but

keep for studying.• CW: Notes 13.1-13.4 • VOTE ON KINETICS TESTING CH. 12• Iodine Clock Rxn. Simulation Lab - as time permits

Dynamic Nature of Equilibrium

When a system reaches equilibrium, the forward and reverse reactions continue to occur … but at equal rates.

We are usually concerned with the situation after equilibrium is reached.

After equilibrium the concentrations of reactants and products remain constant.

Section 13.1

The Equilibrium Condition

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• The state where the concentrations of all reactants and products remain constant with time.

• On the molecular level, there is frantic activity. Equilibrium is not static, but is a highly dynamic situation.

Section 13.1

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Equilibrium Is:

• Macroscopically static • Microscopically dynamic

Section 13.1

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Changes in Concentration

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

Section 13.1

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• Concentrations reach levels where the rate of the forward reaction equals the rate of the reverse reaction.

Section 13.1

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The Changes with Time in the Rates of Forward and Reverse Reactions

Section 13.1

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Concept Check

Consider an equilibrium mixture in a closed vessel reacting according to the equation:

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

You add more H2O(g) to the flask. How does the concentration of each chemical compare to its original concentration after equilibrium is reestablished? Justify your answer.

Section 13.1

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Concept Check

Consider an equilibrium mixture in a closed vessel reacting according to the equation:

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

You add more H2 to the flask. How does the concentration of each chemical compare to its original concentration after equilibrium is reestablished? Justify your answer.

Section 13.2

Atomic Masses

The Equilibrium Constant

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Consider the following reaction at equilibrium:

jA + kB lC + mD

• A, B, C, and D = chemical species.

• Square brackets = concentrations of species at equilibrium.

• j, k, l, and m = coefficients in the balanced equation.

• K = equilibrium constant (given without units).

[B][A]

[D] [C]K =

Law of Mass ActionLaw of Mass Action

Section 13.2

Atomic Masses

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Conclusions About the Equilibrium Expression

• Equilibrium expression for a reaction is the reciprocal of that for the reaction written in reverse.

• When balanced equation for a reaction is multiplied by a factor of n, the equilibrium expression for the new reaction is the original expression raised to the nth power;

• thus Knew = (Koriginal)n.

• K values are usually written without units.

Conclusions about Equilibrium Conclusions about Equilibrium ExpressionsExpressions

The equilibrium expression for a reaction The equilibrium expression for a reaction is the reciprocal for a reaction written in is the reciprocal for a reaction written in reversereverse

2NO2NO22(g) (g) 2NO(g) + O 2NO(g) + O22(g(g))

2NO(g) + O2NO(g) + O22(g) (g) 2NO2NO22(g)(g)

Conclusions about Equilibrium Conclusions about Equilibrium ExpressionsExpressions

When the balanced equation for a reaction When the balanced equation for a reaction is multiplied by a factor is multiplied by a factor nn, the equilibrium , the equilibrium expression for the new reaction is expression for the new reaction is the the original expressionoriginal expression, raised to the , raised to the nthnth power. power.

2NO2NO22(g) (g) 2NO(g) + O 2NO(g) + O22(g(g))

NONO22(g) (g) NO(g) + ½O NO(g) + ½O22(g(g))

Section 13.2

Atomic Masses

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• K always has the same value at a given temperature regardless of the amounts of reactants or products that are present initially.

• For a reaction, at a given temperature, there are many equilibrium positions but only one value for K. Equilibrium position is a set of equilibrium

concentrations.

Section 13.3

The Mole Equilibrium Expressions Involving Pressures

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• K involves concentrations - (also called Kc)

• Kp involves pressures for gases.

Section 13.3

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Example

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

Section 13.3

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Example

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

Equilibrium pressures at a certain temperature:

Section 13.3

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Example

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

Section 13.3

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The Relationship Between K and Kp

Kp = K(RT)Δn

• Δn = sum of the coefficients of the gaseous products minus the sum of the coefficients of the gaseous reactants.

• R = 0.08206 L·atm/mol·K• T = temperature (in kelvin)

Section 13.3

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Example

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

Using the value of Kp (3.9 × 104) from the previous example, calculate the value of K at 35°C.

Section 13.4

Heterogeneous Equilibria

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Homogeneous Equilibria

• Homogeneous equilibria – involve the same phase:

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

HCN(aq) H+(aq) + CN-(aq)

Section 13.4

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• Heterogeneous equilibria – involve more than one phase:

2KClO3(s) 2KCl(s) + 3O2(g)

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

Section 13.4

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• The position of a heterogeneous equilibrium does not depend on the amounts of pure solids or liquids present. The concentrations of pure liquids and solids

are constant.

2KClO3(s) 2KCl(s) + 3O2(g)

Equilibria Involving PureSolids and Liquids

• The equilibrium constant expression does not include terms for pure solid and liquid phases because their concentrations do not change in a reaction.

• Although the amounts of pure solid and liquid phases change during a reaction, these phases remain pure and their concentrations do not change.

[CaO] [CO2]Kc = –––––––––– [CaCO3]

Kc = [CO2]

Example: CaCO3(s) CaO(s) + CO2(g)

Section 13.4

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ASSIGNMENTS - 2/20/14

• W/sheet #1 problems due Monday • This assignment should be able to be completed today

in class but is due Monday.

• Ch. 13 Equilibrium packet - #1-#6 practice problems - answers shown to see if you are doing correctly.

• Prepare for Test on Kinetics.• HW: Read chapter 13 over the next week.

• Ch. 12 Kinetics Test - VOTED for WEDNESDAY - Feb 26

chapter 13.1-13.4 ap chem chemical equilibrium. 2 chapter 13 table of contents copyright © cengage...

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