Chemical Kinetics

1
Chemical Kinetics

• Chapter 14
• Brown, LeMay, and Bursten

2
Factors that Affect Reaction Rate

• Physical state of reactants - area of contact
  (demo)
• Concentration of reactants (cigarette)
• Temperature - what do you think? (refrigerating
  food)
• Catalyst - enzymes in cells

3
Collision Theory

• Effective collisions of reactants lead to
  products
• Force must be adequate
• Orientation must be correct

4
Rate

• Expressed as a an event per unit time
• Can follow the appearance of a product or the
  disappearance of a reactant
• Units would be Molarity/time or
• mol/liter-s
• Rates are always positive
• Average can be calculated for a long period of
  time
• Graph - slope is instantaneous rate

5
Graph

• Calculus tells us that the average rate
  approaches the instantaneous rate as ?t
  approaches 0.

6
Rate Comparison of Reactants and Products

• For A ? B, rate of disappearance of A equals the
  rate of appearance of B
• For A ? 2B, the rate of disappearance of A 1/2
  rate of appearance of B

7
Concentration and Rate

• Need to determine this experimentally!
• General form
• Rate kAmBn etc.
• From a table of rates at various concentrations,
  exponents can be calculated
• From any single set of data, k can be
  calculated.
• Now the general expression can be determined.

8
Practice table
9
More

• Units of k vary - whatever is necessary to get
  mol/liter-sec
• Language - first order etc. overall order
• What does second order mean?
• What does 0 order mean?
• Usually 0, 1, or 2 but may be something else.
• NOT THE COEFFICIENTS IN THE EQUATION!

10
k

• Changes with temperature and presence of a
  catalyst
• Not affected by concentration

11
Calculations with 0, First, and Second Order

• Table from Masterton
• Can be used to calculate
• how much is gone in x seconds
• how much is left
• how long until half (or any other amount) is
  gone or left

12
First Order

• Exponential Decay (coin example)
• A ? products (decomposition)
• Rate kA1
• Integrate this lnAt/lnAi -kt
• lnAt lnAi kt (describe this line)

13
Half-life of First Order

• Does the half-life of a radioactive element
  depend upon concentration?
• No it is fixed
• t1/2 ln2/k or .693/k

14
Zero Order

• A ? products
• Rate kA0 k
• Rate is independent of concentration! (wire
  example.)
• At Ai kt
• Describe line.

15
Half-life of Zero Order

• Simple t1/2 A0/2k

16
Second Order

• Hard to find a model
• A ? products rate kA2 (collision between 2
  As to break a bond)
• Integrate to 1/At 1/Ai kt (describe this
  line)

17
Half-life of Second Order

• t1/2 1/kA0
• Therefore, half-life depends upon initial
  concentration of reactant.

18
Temperature and Rate

• Must have correct orientation
• Must have a minimum amount of kinetic energy -
  called activation energy (EA)
• Like a hill - diagram
• Arrangement at top is the activated complex or
  transition state

19
Arrhenius Equation

• KE of particles is a distribution at any
  temperature.
• Fraction of particles with KE equal or greater
  than Ea f e-Ea/RT
• Found that rate did not increase linearly with
  temperature
• K Ae-Ea/RT
• A is a factor that considers number of collisions
  and probability of a favorable collision

20
Determining Ea

• Take natural log of previous equation
• lnk -Ea/RT lnA
• Graph lnk vs 1/T, determine slope
• Nongraphically, convert above equation for two
  temperatures
• Ln(k1/k2) Ea/R(1/T2 - 1/T1)

21
Reaction Mechanisms

• Describes how a reaction actually occurs
• Each step is called an elementary step
• Molecularity - number of molecules involved in a
  collision.
• Single molecule - unimolecular
• Two molecules - bimolecular
• Three - termolecular
• Other - unlikely

22
Multistep Mechanisms

• Several steps must add to give original reaction
• Intermediate - neither a reactant nor product in
  the final reaction
• Multistep mechanisms involve one or more
  intermediates

23
Rate Law and Mechanisms

• Important to come up with mechanisms that explain
  the rate law seen experimentally
• Rate Laws follow directly from molecularity.
• Table

24
Multistep Mechanisms

• Most likely
• The slowest step is the rate-determining step
• Example
• NO2 CO ? NO CO2
• It is second order for NO2,0 order for CO
• Propose a mechanism.

25
Catalysis

• Definition - substance that increases the rate of
  a chemical reaction without being used up in the
  process.
• Very common and important - body, industry
• Lowers the overall activation energy of a
  chemical reaction (Arrhenius equation)
• Usually provides a completely different mechanism

26
Energy diagram
27
Kinds of Catalysts

• Homogeneous - same phase
• Examples are gases, substances in solution
• Heterogeneous - different phase
• Often a solid in contact with a liquid or gas
• Reactants are usually adsorbed on surface of
  catalyst at active sites