Chapter 6 Chemical Kinetics - PowerPoint PPT Presentation

Loading...

PPT – Chapter 6 Chemical Kinetics PowerPoint presentation | free to download - id: 53e07b-YTZjY



Loading


The Adobe Flash plugin is needed to view this content

Get the plugin now

View by Category
About This Presentation
Title:

Chapter 6 Chemical Kinetics

Description:

Title: PowerPoint Presentation - Chapter 14 Kinetics Author: Swiftney, Ty SHS-Staff Last modified by: Windows User Created Date: 12/30/2004 11:01:46 PM – PowerPoint PPT presentation

Number of Views:45
Avg rating:3.0/5.0
Slides: 30
Provided by: Swiftne
Learn more at: http://www.swiftchem.org
Category:

less

Write a Comment
User Comments (0)
Transcript and Presenter's Notes

Title: Chapter 6 Chemical Kinetics


1
Chapter 6Chemical Kinetics
1
2
Kinetics
  • Why?
  • Studies the rate at which a chemical process
    occurs.
  • Besides information about the speed at which
    reactions occur, kinetics also sheds light on the
    reaction mechanism (exactly how the reaction
    occurs).

2
3
Factors That Affect Reaction Rates
  • Concentration of Reactants
  • As the concentration of reactants increases, so
    does the likelihood that reactant molecules will
    collide.
  • Temperature
  • At higher temperatures, reactant molecules have
    more kinetic energy, move faster, and collide
    more often and with greater energy.
  • Catalysts
  • Speed reaction by changing mechanism
  • Nature of Reactants
  • Speed of reaction may depend
  • on the complexity of the
  • molecules reacting

3
4
Reaction Rates
  • Rates of reactions can be determined by
    monitoring the change in concentration of either
    reactants or products (or both) as a function of
    time. ?A vs ?t

4
5
Reaction Rates
C4H9Cl(aq) H2O(l) ??? C4H9OH(aq) HCl(aq)
C4H9Cl M
  • In this reaction, the concentration of butyl
    chloride, C4H9Cl, was measured at various times,
    t.

5
6
Reaction Rates
C4H9Cl(aq) H2O(l) ??? C4H9OH(aq) HCl(aq)
Average Rate, M/s
  • The average rate of the reaction over each
    interval is the change in concentration divided
    by the change in time

6
7
Reaction Rates
C4H9Cl(aq) H2O(l) ??? C4H9OH(aq) HCl(aq)
  • Note that the average rate decreases as the
    reaction proceeds.
  • This is because as the reaction goes forward,
    there are fewer collisions between reactant
    molecules.

7
8
Reaction Rates
C4H9Cl(aq) H2O(l) ??? C4H9OH(aq) HCl(aq)
  • A plot of concentration vs. time for this
    reaction yields a curve like this.
  • The slope of a line tangent to the curve at any
    point is the instantaneous rate at that time.

8
9
Reaction Rates
C4H9Cl(aq) H2O(l) ??? C4H9OH(aq) HCl(aq)
  • The reaction slows down with time because the
    concentration of the reactants decreases.

9
10
Reaction Rates and Stoichiometry
C4H9Cl(aq) H2O(l) ??? C4H9OH(aq) HCl(aq)
  • In this reaction, the ratio of C4H9Cl to C4H9OH
    is 11.
  • Thus, the rate of disappearance of C4H9Cl is the
    same as the rate of appearance of C4H9OH.

10
11
Reaction Rates and Stoichiometry
  • What if the ratio is not 11?
  • H2(g) I2(g) ??? 2 HI(g)
  • Only 1/2 H2 is used for each HI made.

11
12
Reaction Rates and Stoichiometry
  • To generalize, for the reaction

Reactants (decrease)
Products (increase)
12
13
Reaction Rates Determination
  • There are several methods for determination of
    Rates
  • Change in Volume of a gas
  • Gas is given off can be collected and measured
  • Change in Mass
  • If some products are gases (they can escape),
    remains have less mass.
  • Change in transmission of light (Colorimetry)
  • Color(s) can change as reactant is used up or as
    product(s) are formed.

13
14
Reaction Rates Determination
  • There are several methods for determination of
    Rates
  • Change in concentration by Titration
  • Taking a sample, stopping the reaction and
    measure.
  • Change of concentration by Conductivity
  • Increase/decrease of ions changes conductivity
  • Clock Reactions
  • Letting a reaction run to a completion point.
    Measure the time.

14
15
Concentration and Rate
  • Each reaction has its own equation that gives its
    rate as a function of reactant concentrations.
  • (?this is called its Rate Law) (HL)
  • To determine the rate law we measure the rate at
    different starting concentrations.

15
16
Rate Laws (HL)
  • A rate law shows the relationship between the
    reaction rate and the concentrations of
    reactants.
  • For gas-phase reactants use PA instead of A.
  • k is a constant that has a specific value for
    each reaction.
  • The value of k is determined experimentally.
  • The Rate Constant is relative
  • k is unique for each reaction
  • k changes with T

16
17
Temperature and Rate
  • Generally, as temperature increases, so does the
    reaction rate.
  • This is because the rate constant, k, depends on
    the temperature.

17
18
The Collision Model (Kinetic Molecular Theory
(KMT))
  • In a chemical reaction, bonds are broken and new
    bonds are formed.
  • Molecules can only react if
  • they collide with each other correctly
  • They have enough activation energy

18
19
The Collision Model
  • Molecules must collide with the correct
    orientation and with enough energy to cause bond
    to break and reform again.

19
20
Activation Energy
  • The minimum amount of energy required for
    reaction to happen is called the activation
    energy, Ea.
  • Just as a ball cannot get over a hill if it does
    not roll up the hill with enough energy, a
    reaction cannot occur unless the molecules
    possess sufficient energy to get over the
    activation energy barrier.

20
21
Reaction Diagrams
  • It is helpful to visualize energy changes
    throughout a process on a reaction diagram like
    this one for the rearrangement of methyl
    isonitrile.

21
22
Reaction Diagrams
  • It shows the energy of the reactants and products
    (and, therefore, ?E (also known as ?H).
  • The high point on the diagram is the transition
    state.
  • The energy gap between the reactants and the
    activated complex is the activation energy
    barrier (Ea).

22
23
MaxwellBoltzmann Distributions
  • Temperature is defined as a measure of the
    average kinetic energy of the molecules in a
    sample.
  • At any temperature there is a wide distribution
    of kinetic energies.

23
24
MaxwellBoltzmann Distributions
  • As the temperature increases, the curve flattens
    and broadens.
  • Thus at higher temperatures, a larger population
    of molecules has higher energy.

24
25
MaxwellBoltzmann Distributions
  • If the dotted line represents the activation
    energy, as the temperature increases, so does the
    fraction of molecules that can overcome the
    activation energy barrier.
  • As a result, the reaction rate increases
  • more molecules with enough energy
  • Molecules moving faster (higher T), more likely
    to have more collisions

25
26
Factors affecting Rate of Reaction
  • Temperature increasing of molecules with
    correct Ea
  • Concentration - increasing of molecules
    available for collisions
  • Particle Size decreasing size while maintaining
    mass, more surface area to react
  • Pressure (gases) increasing pressure can result
    in more collisions
  • Catalyst lowers the Ea of a reaction
  • can increase the number of molecules with correct
    Ea

26
27
Catalysts
  • Catalysts increase the rate of a reaction by
    decreasing the activation energy of the reaction.
  • Catalysts change the mechanism by which the
    process occurs.

27
28
Catalysts
  • One way a catalyst can speed up a reaction is by
    holding the reactants together and helping bonds
    to break.

28
29
Enzymes
  • Enzymes are catalysts in biological systems.
  • The substrate fits into the active site of the
    enzyme much like a key fits into a lock.

29
About PowerShow.com