Honors Unit 11 Notes - Part I: Chemical Kinetics

Name: _______________________________________________

Honors Unit 11 Notes Part I: Chemical Kinetics

OBJECTIVES

Derive rate expressions from experimental data.

Interpret potential energy diagrams.

Understand the concept of reaction mechanism, including the rate determining step and overall reaction. Identify intermediates and catalysts.

Rates of Chemical Reactions

Kinetics — the study of ________________________________ and the _________________

(the way the reaction proceeds)

o Only kinetics will tell you ______________________ the reaction happens!

o Thermodynamics will tell you if the reaction is product or reactant favored

A rate is any __________________________________________________________________

o Example:

Reaction Rate – change in ___________________________ of a reactant or product over time

o Units:

Types of Rates:

o _________________________ = rate at time zero

o _____________________________ = the rate over a given time interval

o _________________________________ = the slope of the tangent line at a given point

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A general reaction rate is calculated by dividing rate expressions by stoichiometric coefficients

Example: For the reaction aA + bB cC + dD

o Disappearance of A =

o Appearance of D = Collision Theory of Reactants

Reactions occur when molecules collide to exchange or rearrange atoms

Effective collisions occur when molecules have the correct __________________ and

_____________________________

Factors Affecting Rates:

o ______________________________ and physical state of reactants and products

A rate law relates the rate of the reaction to the concentration of the reactants

o

o

Catalysts are substances that speed up a reaction but are unchanged by the reaction

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Rate Laws

For aA + bB cC + dD, the rate law: Rate = k[A]m[B]n

o k =

o The exponents must be determined by performing an experiment

o They are NOT derived from the ________________________________________ in the overall chemical equation!

The exponents m, n, and p: Rate = k[A]m[B]n[C]p

o Are the __________________________________

o Can be 0, 1, 2, or fractions (can be other whole numbers in fictional examples)

o Must be determined by experimentation

o Overall Order =

Interpreting Rate Laws: Rate = k [A]m[B]n[C]p

• If m = 1, reaction is 1st order with respect to A Rate = k [A]1

If [A] doubles, then rate ____________________ (goes up by a factor of ________)

• If m = 2, reaction is 2nd order with respect to A Rate = k [A]2

If [A] doubles, then rate __________________________ (increases by a factor of ______)

• If m = 0, reaction is zero order with respect to A Rate = k [A]0

If [A] doubles, rate ______________________________ Rate Constant, k

Relates ________________and _____________________________at a given temperature

General Formula for units of k:

Overall Order Units of k

0

1

2

3

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Example #1: The initial rate of decomposition of acetaldehyde, CH3CHO, was measured at a series of different concentrations and at a constant temperature.

CH3CHO(g) CH4(g) + CO(g)

[CH3CHO] (mol/L)

0.162 0.195 0.273 0.410 0.513

Rate (mol/L·min)

3.15 4.56 8.94 20.2 35.2

a) Using the above data, determine the order of the reaction; that is, determine the value of m

in the equation: Rate = k[CH3CHO]m

Strategy You are looking at how the concentration affects the rate so compare the two in a

proportion! Pick any two points from the given data.

𝑅𝑎𝑡𝑒2𝑅𝑎𝑡𝑒1

=[𝐴]2

𝑚

[𝐴]1𝑚 = (

[𝐴]2[𝐴]1

)𝑚

b) Using the same set of data from part a, and knowing the order of the reaction, determine the value of the rate constant, k (with units!)

c) Determine the rate of the reaction when [CH3CHO] = 0.452 mol/L

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Example #2: The data below are for the reaction of nitrogen (II) oxide with hydrogen at 800°C.

2NO(g) + 2H2(g) N2(g) + 2H2O(g)

Initial Concentration

(mol/L) Initial Concentration

(mol/L) Rate of Formation of N2

(mol/L·min)

Experiment [NO] [H2]

1 0.0010 0.0040 0.12

2 0.0020 0.0040 0.48

3 0.0030 0.0040 1.08

4 0.0040 0.0010 0.48

5 0.0040 0.0020 0.96

6 0.0040 0.0030 1.44

a) Determine the order of the reaction with respect to both reactants and write the rate law. b) Calculate the value of the rate constant (with units). c) Determine the rate of formation of product when [NO] = 0.0024 M and [H2] = 0.0042 M.

Strategy

Choose two experiments where the concentration of one reactant is constant and the other is changing

Solve for m and n separately using one experiment at a time 𝑅𝑎𝑡𝑒2

𝑅𝑎𝑡𝑒1=

[𝐴]2𝑚

[𝐴]1𝑚 = (

[𝐴]2

[𝐴]1)𝑚

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Example #3: The initial rate of a reaction A + B C was measured with the results below. (a) Write the rate law, (b) the value of the rate constant, k (with units), and (c) the rate of reaction when [A] = 0.050 M and [B] = 0.100 M.

Experiment [A] (M) [B] (M) Initial Rate (M/s)

1 0.10 0.10 4.0x10-5

2 0.10 0.20 4.0x10-5

3 0.20 0.10 1.6x10-4

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Potential Energy Diagrams

Molecules need a minimum amount of energy for a reaction to take palce.

____________________________________ ( ) – the minimum amount of energy that the reacting species must possess to undergo a specific reaction

_____________________________________ – a short lived molecule formed when reactants

collide; it can return to reactants or form products

o Formation depends on the ________________________________________ and the

______________________________________ of the molecules

On the potential energy diagram, identify:

o Energy of Reactants, Energy of Products, Hrxn, EA – Energy of Activation, Energy of Activated Complex, Catalyzed pathway, Exothermic or Endothermic?

***Catalysts lower activation energy!

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Reaction Mechanisms

Mechanism =

o Most reaction do not occur in a single step. They occur in a series of elementary steps!

Elementary step =

Rate determining step =

Overall reaction = Adding elementary steps gives the net (overall) reaction.

Ex.) COCl2 (g) COCl (g) + Cl (g) fast Cl (g) + COCl2 (g) COCl (g) + Cl2 (g) slow 2 COCl (g) 2 CO (g) + 2 Cl (g) fast 2 Cl (g) Cl2 (g) fast

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Intermediate – _________________ in one elementary step but _______________ in another

Ex.) NO (g) + O3 (g) NO2 (g) + O2 (g) NO2 (g) + O (g) NO (g) + O2 (g) --------------------------------------------------------------

Catalyst – a reactant in an elementary step, but is unchanged at the end of the reaction

o A substance that speeds up a reaction but is not permanently changed by the reaction

o

Example #4 Cl2(g) ↔ 2Cl(g) Fast

Cl(g) + CHCl3(g) CCl3(g) + HCl(g) Slow CCl3(g) + Cl(g) CCl4(g) Fast

Overall (Net) Reaction: Rate Determining Step: Intermediate(s): Catalyst(s):

Example #5 H2O2(aq) + I1-(aq) H2O(l) + IO1-(aq) Slow

H2O2(aq) + IO1-(aq) H2O(l) + O2(g) + I1-(aq) Fast Overall (Net) Reaction: Rate Determining Step: Intermediate(s): Catalyst(s):

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Example #6 O3(g) + Cl(g) O2(g) + ClO(g) Slow

ClO(g) + O(g) Cl(g) + O2(g) Fast Overall (Net) Reaction: Rate Determining Step: Intermediate(s): Catalyst(s):

------------------------------------End of Kinetics----------------------------------

Honors Unit 11 Notes Part II: Chemical Equilibria

OBJECTIVES

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

Write an equilibrium expression.

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

Nature of Equilibrium

What is equilibrium?

(dictionary) Equilibrium = a state of rest or balance due to equal action of opposing forces

Chemical Equilibrium = Not all chemical reactions are _______________________________!!!

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General Characteristics of Equilibrium

o __________________________________________________________________

o _____________________________________

o Can be approached from either direction (reaction can run in the forward direction or the reverse direction)

o After a period of time, the concentrations of reactants and products are __________________

The forward and reverse reactions _____________________ after equilibrium is attained!

The Equilibrium Expression, K or Kc

For the reaction:

The brackets "[ ]" represent:

"a, b, c, and d" represent:

The "c" in Kc indicates:

dDcCbB aA ba

dc

c

BA

DCK

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There are two cases when a species is not shown in the equilibrium expression:

1.

2. Example #7: Write the equilibrium expression for the oxidation-reduction reaction occurring between iron (III) chloride and tin (II) chloride:

2 FeCl3 (aq) + SnCl2 (aq) 2 FeCl2 (aq) + SnCl4 (aq)

Example #8: Write the equilibrium expression for the replacement of silver ions by copper:

Cu (s) + 2 Ag+ (aq) Cu2+ (aq) + 2 Ag (s)

Le Chatelier's Principle

Le Chatelier’s Principle:

“If a system at equilibrium is stressed, the system tends to shift its equilibrium position to counter the effect of the stress.”

How does a “stress” influence equilibrium?

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The Seesaw Analogy

“Stressors” (factors) that Can Cause Changes at Equilibrium

1.

2.

3.

4.

5. 1. Changes in Amount of Species_

o Add reactant; system shifts to the ___________________

o Add product; system shifts to the ___________________

o Remove reactant; system shifts to the _______________

o Remove product; system shifts to the ________________

Example #9: Predict the direction of shift of the following concentration changes on the reaction:

CH4(g) + 2S2(g) ⇔ CS2(g) + 2H2S(g)

(A) Some S2 (g) is added.

(B) Some CS2 (g) is added.

(C) Some H2S (g) is removed.

(D) Some argon gas (an inert gas) is added.

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2. Changes in Pressure or Volume_

o If pressure goes down (same as volume going up), system shifts towards increased

_________________________________________________

o If pressure goes up (same as volume going down), system shifts towards decreased

_________________________________________________

Example #10_ Predict the effect of increasing pressure (decreasing volume) on each of the reactions.

(a) CH4(g) + 2S2(g) ⇔ CS2(g) + 2H2S(g) (c) CO2(g) + C(s) ⇔ 2CO(g)

(b) H2(g) + Br2(g) ⇔ 2HBr(g) (d) PCl5(g) ⇔ PCl3(g) + Cl2(g)

3. Changes in Temperature

o Write heat as a product (for exothermic reactions) or reactant (for endothermic reactions)

o System shifts to get rid of added heat:

o System shifts _____________ for exothermic reactions as temperature goes ________

o System shifts ______________ for endothermic reactions as temperature goes ______

Example #11_ Predict the effect of increasing temperature on each of the following reactions.

(a) CO(g) + 3H2(g) ⇔ CH4(g) + H2O(g) ΔH < 0

(b) CO2(g) + C(s) ⇔ 2CO(g) ΔH > 0

(c) 4NH3(g) + 5O2(g) ⇔ 4NO(g) + 6H2O(g) EXOTHERMIC

(d) 2H2O(g) ⇔ 2H2(g) + O2(g) ENDOTHERMIC

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4. Adding an Inert Substance

o If a substance is NOT in the reaction (or in the Keq expression) any changes will have

______________________________ on equilibrium!

Ex.) If CO2 is added to the system below, there would be no effect on the equilibrium.

H2 (g) + Br2 (g) 2 HBr (g) 5. Adding a Catalyst

o Catalysts – increase the _________________________________________________ by lowering the activation energy

o Adding a catalyst will NOT affect equilibrium. It only changes the rate at which you reach equilibrium!

Example #12_ How can the reaction below be shifted to the right? List all possibilities!

CO(g) + 2 H2 (g) CH3OH (g) ∆H = positive

Example #13_ How can the reaction below be shifted to the right? List all possibilities! This process

(the Haber process) is used in industry to produce ammonia.

N2(g) + 3H2(g) 2NH3(g) ∆H = negative