Why do some reactions happen and others don

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Why do some reactions happen and others dont?
Are the products more stable than the reactants?
Thermodynamics
Does the reaction go at a reasonable rate?
Kinetics
Chapter 14 Chemical KineticsRates of Reactions
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Control of Reactivity
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Collision Theory
For a reaction to take place - Molecules must
collide - They must do so in the correct
orientation - They must collide with an energy
greater than the activation energy

• Consider NO O3 ? NO2 O2

product molecules separate
Molecules collide
Bonds are formed and break
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What would control how fast a reaction happens?
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So, what controls the rate of a reaction?

• Number of collisions
• How often they collide in a shape that allows new
  bonds to form
• The energy of the colliding reactant molecules

• Well consider dependence on
• Concentration
• a. Rate laws
• b. Concentration vs. time relationships
• Temperature and activation energy
• Mechanism

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Concentration Dependence

• It makes sense that as concentration increases,
  the number of collisions per second will increase
• Therefore, in general, as concentration
  increases, rate increases
• But, it depends on which collisions control the
  rate
• So, you cant predict concentration dependence
  it must be measured experimentally

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An analogy of love
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business
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The CO2 in my DietCoke is

1. Saturated
2. Unsaturated
3. Supersaturated

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The CO2 in my DietCoke is now

1. Saturated
2. Unsaturated
3. Supersaturated

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Rates
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Concentration-Time Curvessimulation
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The reaction

1. Speeds up as it goes
2. Slows down as it goes
3. Keeps the same rate

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Types of measured rates

• Rate over time
• Instantaneous rate
• Initial rate

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Example of rate measurement
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Rate Laws (also called Rate Equations)

• For the reaction 2 N2O5 ? 4 NO O2
• Rate kN2O5
• For the reaction NO2 ? NO ½ O2
• Rate kNO22
• For the reaction CO NO2 ? CO2 NO
• Rate kCONO2

first order reaction
second order reaction
first order in CO and in NO2 second order
overall
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What is the overall order for a reaction with
rate kCO22H

1. 0 order
2. 1st order
3. 2nd order
4. 3rd order

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Determining a Rate Law

• Determining the rate law must be done by
  experiment the reaction equation does not tell
  you the rate law
• Two methods Initial Rates and the Graphical
  Method

• Method of Initial Rates
• Measure the rate of the reaction right at the
  start.
• Vary the starting concentrations
• Compare initial rates to initial concentrations

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Determining a Rate Law Initial Rate Method

• Isolation of variables Vary only one
  concentration at a time and keep temperature
  constant
• If concentration doubles and
• Rate does not change, then zero order
• Rate doubles, then first order
• Rate quadruples, then second order
• General Rule

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Initial Rate Method Example 1
What is the rate law?
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Simulation A? B

1. 0 order
2. 1st order
3. 2nd order

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Simulation C? D

1. 0 order
2. 1st order
3. 2nd order

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Simulation E? F
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Initial Rate Method Example 2
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Initial Rate Method Example 2
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Concentration-Time Relationships
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Graphical Method for Determining Rate Laws
A plot of 1/R vs. Time will be linear.
A plot of concentration vs. Time will be linear.
A plot of lnR vs. Time will be linear.
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Graphical Method for Determining Rate Laws
How it works 1. Collect R over an interval of
times. 2. Make plots of R vs. time lnR
vs. time 1/R vs. time Only one will be linear.
That tells you the reaction order. The slope of
the linear plot is the rate constant.
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Graphical Method for Determining Rate Laws

• Example 2 H2O2 ? 2 H2O O2
• Time(min) H2O2(mol/L)
• 0 0.0200
• 200 0.0160
• 400 0.0131
• 600 0.0106
• 800 0.0086
• 1000 0.0069

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Graphical Method for Determining Rate Laws Order

• Example 2 H2O2 ? 2 H2O O2
• Time(min) H2O2(mol/L)
• 0 0.0200
• 200 0.0160
• 400 0.0131
• 600 0.0106
• 800 0.0086
• 1000 0.0069

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Graphical Method for Determining Rate Laws k
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Kinetics Lab
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Example 1
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Example 2
The decomposition of nitrous oxide at 565 oC, 2
N2O ? 2 N2 O2 is second order in N2O. If the
reaction is initiated with N2O equal to 0.108
M, and drops to 0.940 M after 1250 s have
elapsed, what is the rate constant?
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Half-Life t1/2

• the time it takes for half the reactant
  concentration to drop to half of its original
  value

First Order Reaction 2 H2O2 ? 2 H2O O2Rate
kH2O2 k 1.05 x 10-3/min
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Cool things about half-life
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Calculations involving Half-Life

• For a first order reaction

What is the relationship between t1/2 and k?
What is the relationship between t1/2 and k for a
second order reaction?
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Radioactive Decay

• All radioisotopes decay via first order
  reactions. Instead of concentrations, amounts are
  used.

Measured as radioactive activity, in counts per
minute (cpm) using a detector.
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Radioactive Decay Example 1

• Radioactive gold-198 is used in the diagnosis of
  liver problems. The half-life of this isotope is
  2.7 days. If you begin with a 5.6-mg sample of
  the isotope, how much of this sample remains
  after 1.0 day?

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Radioactive Decay Carbon Dating
C-14 In living thing
Sunlight Nitrogen
Atmospheric C-14
C-14 Dead thing
Sunlight Nitrogen
Atmospheric C-14
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Radioactive Decay Example 2

• The Carbon-14 activity of an artifact in a burial
  site is found to be 8.6 counts per minute per
  gram. Living material has an activity of 12.3
  counts per minute per gram. How long ago did the
  artifact die? t1/2 5730 years

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What is the relationship between half-life and k
for a second order reaction?
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Reacting the Fuel Fission Reactions
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Chain Reactions
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Controlling the Reactions Control Rods
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