AP Chemistry

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AP Chemistry

• Unit 13
• Kinetics

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Questions

• What is the collision theory?
• What are four ways to increase the rate of
  reaction?

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Chemical Kinetics

• Chemical Kinetics a study of the rates of
  chemical reactions, factors affecting these
  rates, and the sequence of molecular events that
  result in chemical reactions

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Chemical Kinetics Cont.

• Reactions occur at different rates
• Kinetics allow for the measurement of a chemical
  reaction as well as a prediction of the products

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Chemical Kinetics Cont.

• Kinetics allow for the mechanism of a chemical
  reaction to be studied

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Factors Affecting Rates of Reaction

• 1. Concentration the larger the amounts of
  reactants, the faster a reactions occurs
• 2. Temperature At higher temperatures,
  particles move faster and allow for particles to
  interact more with other particles

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Factors Affecting Rates of Reaction

• 3. Surface Area the larger the surface area,
  the more likely molecules will interact with each
  other
• 4. Catalysts increase the rate of reaction
  while remaining unchanged

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Rates of Reaction

• The unit used for rates of reaction is mole per
  liter (M) per unit time
• Rate ? substance / ? t
• Rates are always written as positive values

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Rates of Reaction

• Rates can be viewed as
• Rate - rate of disappearance of reactants
• OR
• Rate rate of appearance of products

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Example

• H2O2 decomposes to form H2O and O2
• 2 H2O2 (aq) ? 2 H2O (l) O2 (g)
• Rate - ? H2O2 / ? t OR
• Rate ? H2O / ? t OR
• Rate ? O2 / ? t

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Discussion of Example

• 2 H2O2 (aq) ? 2 H2O (l) O2 (g)
• Rate - ? H2O2 / ? t OR
• Rate ? O2 / ? t
• O2 forms 1/2 as fast as H2O2 decomposes

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General Rule for Rate of Reaction

• The general rate of reaction is found by dividing
  the rate of disappearance of a reactant by the
  rate of formation of a product by the
  stoichiometric coefficient of the reactant or
  product in the balance equation

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General Equation for Rate of Reaction

• a A b B ? c C d D
• a, b, c, d are coefficients in balanced
  equations
• A, B, C, D represent reactants and products

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General Equation for Rate of Reaction
Rate
OR
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General Equation for Rate of Reaction
Rate
OR
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Average Rate of Reaction

• Average Rate of Reaction measures the change in
  concentration per unit between two points

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Example

• A 2 B ? 3 C 2 D
• Initially A 0.4658 M. After 125 seconds, A
  0.4282 M.
• A) Calculate the average rate of reaction for
  A
• B) What is the rate of formation for C?
• A) 0.0003 B) 0.0009

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The Rate Law

• The concentrations of certain substances must be
  determined by experiment
• The Rate Law relates the rate of reaction to the
  concentration of reactants

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The Rate Law

• The Rate Law is
• Rate k Am Bn
• A molar concentration of A
• B molar concentration of B

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The Rate Law

• m, n determined based on experimental data
• m n order of reaction
• k rate constant

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More About k

• k depends on
• a. Particular reaction
• b. Temperature
• c. Catalysts
• The units for k are M(1-Overall order) s-1

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Method of Initial Rates

• The method of initial rates is a procedure to
  determine the order of a reaction by varying
  initial amounts of reactants and analyzing rates

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Method of Initial Rates Continued

• The amounts of reactants vary by multiples
• Use of this method also minimizes the effects of
  reversible reactions
• Determine the rate of reaction for
• 2 NO Cl2 ? 2 NOCl

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Examples of Method of Initial Rates

• Exp Initial Initial Initial Rate NO
  Cl2
• 1 0.0125 0.0255 2.27 x 10-5
• 2 0.0125 0.0510 4.55 x 10-5
• 3 0.0250 0.0255 9.08 x 10-5
• Calculate the rate of reaction

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Solution

• Focus on each reactant one at a time
• Consider the changes which occur
• Compare the change in concentration to the rate
• Ans) Rate k NO2Cl2

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• Ans) Rate k NO2Cl2
• Questions
• 1. What is the order with respect to NO?
• 2. What is the order with respect to Cl2?
• 3. What is the overall reaction order?

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Review WS 13.1

• See Handout

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Determine Rate of Reaction

• Exp Initial Initial Initial Rate NO
  Cl2
• 1 0.12 0.18 7.91 x 10-2
• 2 0.060 0.18 3.95 x 10-2
• 3 0.060 0.090 9.88 x 10-3

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Determine Rate of Reaction

• Exp Initial Initial Initial Rate O2
  H2
• 1 0.025 0.15 0 2.813 x 10-8
• 2 0.050 0.189 1.425 x 10-7
• 3 0.025 0.260 4.875 x 10-8

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Collision Theory

• 1. The reacting molecules must collide with each
  other
• 2. The reacting molecules must collide with
  sufficient energy

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Collision Theory

• 3. The molecules must collide in an orientation
  that can lead to the rearrangement of atoms

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Factors Increasing Collisions

• 1. Increases in concentration usually result in
  increased rate of reaction
• 2. Increase in temperature means particles move
  faster and therefore collisions occur more
  frequently

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Activation Energy

• An energy barrier must be surmounted by reactants
  in order for a reaction to occur
• This energy barrier is called activation energy
  (Ea)

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Activation Energy

• Reactions with low activation energies occur
  quickly
• Reactions with high activation energies occur
  more slowly
• The energy of the reactant molecules serve as the
  basis for Ea

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Transition State Theory

• As molecules collide, bonds break gradually while
  new bonds formed
• The transition state occurs when atoms have some
  characteristics of the reactants while forming
  the products

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Transition State Theory

• Example
• I- CH3-Br ? Initial State
• I CH3 Br ? Transition State

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Transition State Theory

• Example
• I- CH3-Br ? Initial State
• I CH3 Br ? Transition State

I- CH3-Br ? CH3-I Br - Final State
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Transition State Theory Continued

• The transition state occurs at the moment when
  the Ea has been reached
• The resulting species has some characteristics of
  both products and reactants

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Classwork / Homework

• Pg. 589 12, 30, 32, 38 a b, 44, 60

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Writing Reactions

• 1. The heating of sodium carbonate
• 2. Sulfur trioxide is bubbled through water
• 3. Aqueous sodium iodide is mixed with lead (II)
  nitrate
• 4. Solid sodium oxide is added to water

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Reaction Profile

• A reaction profile shows how the potential energy
  of a reaction changes with time

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Reaction Profile
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Activity

• Construct a reaction profile that is endothermic
• ?H Ea (forward) - Ea (reverse)

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Review

• 1. What is activation energy?
• 2. What are the three components of the
  collision theories?
• 3. How does a reaction profile of an exothermic
  reaction differ from that of an endothermic
  profile?

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Reaction Mechanisms

• The process by which a reaction occurs is called
  a reaction mechanism
• The mechanism essentially describes the bonds
  that are broken and formed as well as energy
  changes

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Reaction Mechanisms

• In a mechanism, the proposed steps are referred
  to as elementary reactions

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Reaction Mechanisms

• Species that form and are used up in the
  mechanism (not appearing in the end) are called
  intermediates
• Catalysts-are used up in an early step but
  reappear at the end

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Exa-mple Of A Me-chan-ism

• Rxn NO2 CO ? NO CO2
• Proposed Mechanism
• NO2 NO2 ? NO3 NO
• NO3 CO ? NO2 CO2
• When the mechanism is added together, the overall
  reaction should be derived.
• Identify the intermediates.

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Write The Overall Equation for

• The following mechanism is proposed for the
  conversion of ozone (O3) into O2
• O3 ? O2 O
• O3 O ? 2 O2
• Overall Rxn 2 O3 ? 3 O2
• Identify Any Intermediates.

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Elementary Steps

• Elementary steps are the independent steps which
  must occur in the reaction
• ??The rate law for elementary steps equals the
  coefficients as exponents

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Elementary Steps

• The rate law for the rate-determining step is
  directly related to the rate law for the overall
  reaction

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Write Rate Laws For Each Elem. Step

• Step I NO2 F2 ? NO2F F
• Step II F NO2 ? NO2F
• Answers
• Rateequation 1 k NO2 F2
• Rateequation 2 k F NO2
• Which rate law is preferred? (Consider any
  intermediates)

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Write Rate Laws for

• Step I H2O2 ? 2 OH
• Step II H2O2 OH ? H2O HO2
• Step III HO2 OH ? H2O O2
• Determine the overall mechanism.
• Identify any intermediates.
• Which step is the slow step?

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Multistep Mechanisms

• Since many reactions occur in multiple steps,
  there is usually one step that is slower than all
  others
• This step is known as the rate determining step
  for the overall reaction
• The rate determining step governs the rate for
  the overall reaction

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Example

• Step I NO2 NO2 ? NO3 NO (slow)
• Step II NO3 CO ? NO2 CO2 (fast)
• Write the rate law for this mechanism
• Ans) Rate k NO22

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Classwork

• WS 13.2

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First Order Reactions

• First Order Reaction involve only one reactant
• Rate k A

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First Order Reactions

• Integrated Rate Law will describe the
  concentration of reactants per unit time

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• Equation
• ln At / A0 -k t
• ln natural log
• At concentration of A at a certain time
• A0 initial concentration of A
• k rate law constant
• t certain time

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Variation of Equation

• ln At / A0 -k t
• ln At - ln A0 - k t
• ln At - k t ln A0

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Graphs of First Order Reactions

• The graphs of first order reactions are
• ln At vs t (time), produces a straight line
  See Handout
• ln At - k t ln A0
• y m x b
• m - k OR
• k - m

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Example

• The first order rate constant for a particular
  reaction is 1.45 yr-1. An initial concentration
  of 5.0 x 10-7 g/cm3 is present. What is the
  concentration after 1 year?
• Ans) 1.17 x 10-7 g/cm3

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Follow Up Question

• How long will it take until the concentration of
  the sample is 3 x 10-2 g/cm3?
• Ans) .3522 years

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Example 2

• The decomposition of dimethyl ether is first
  order with a rate constant of 6.8 x 10-4 s-1. If
  the initial pressure is 135 torr, what is the new
  pressure after 1420 seconds?
• Hint concentration is proportional to pressure
  from ideal gas law, and Eulers number (?x)
  undoes ln
• Ans) 51 torr

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Zero-Order Rate Law

• Rate k A0 OR
• Rate k
• The integrated rate law for 0th Order
• A - k t A0
• See Handout

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Second-Order Rate Law

• Rate k A2
• The integrated rate law
• 1 / A k t 1 / A0
• See Handout

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Summary of Integrated Rate Laws

• Zero Order
• A - k t A0
• 1st Order
• ln At - k t ln A0
• 2nd Order
• 1 / A k t 1 / A0

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Review - Integrated Rate Law

• 1. Which order of reaction
• a. Graphs ln A
• b. Graphs 1/A
• c. Graphs A

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Beers Law

• Beers Law can be used to determine the
  concentration of solutions
• In order to utilize Beers Law, colorimeters or
  spectrophotometers are used

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Basics of Colorimeters/Spectrophotometers

• A beam of light (P0) is sent through a sample in
  a cuvette
• As the light goes through the sample, some of the
  light is absorbed by particles in the sample
• The amount of light that comes out on the other
  side (P) is analyzed

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Cuvette
Incident Light
Resulting Light
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Absorbance Transmittance

• Transmittance measures the amount of light that
  is passed through the sample
• T LightIN / LightOUT x 100

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Beers Law

• A a b c
• A absorbance
• a molar absorptivity constant
• b cell path length (diameter of cuvette)
• c concentration of sample

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Lab - Beers Law

• Proper Use of Cuvettes Colorimeter
• Lab Procedure

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Determine Rate of Reaction

• Exp Initial Initial Initial Rate Br2
  Cl2
• 1 1.0 2.5 0.045
• 2 2.0 2.5 0.72
• 3 2.0 7.5 19.44