1. AP Chapter 14

2.  Chemical kinetics is the area of chemistry that involves the rates or speeds of chemical reactions.  The more collisions there are between molecules, the faster the reaction occurs.

3. 1. The physical state of the reactants – gases and liquids tend to react faster than solids - think surface area. 2. The concentrations of reactants. 3. The temperature at which the reaction occurs. 4. The presence of a catalyst.

4.  The speed of a chemical reaction is its reaction rate.  Reaction rates are usually expressed as changes in concentration per unit of time. ( M/s )

6.  The preceding slide shows that the rate of the reaction can be expressed as either the rate of disappearance of reactant A, or the appearance of product B.  The overall concentration does not change! Avg rate of disappearance of A = ∆ [A] ∆ t

7.  C4H9Cl( aq ) + H2O( l ) → C4H9OH ( aq ) + HCl ( aq )  It is typical for rates to decrease as a reaction proceeds, because the concentration of the reactants decreases.

9.  An instantaneous rate is the rate at a particular moment in the reaction.  It is determined from the slope of the curve at the point of interest.

11. Instantaneous rate = ∆ [C4H9Cl] ∆ t

12.  In the preceding reaction, 1 mol of C4H9OH is produced for every mol of C4H9Cl consumed. (this is a 1:1 ratio)  Not all equations have this ratio:  2HI (g) → H 2 (g) + I 2 (g)

13.  For a general equation:  a A + b B → c C + d D Rate = 1 ∆[A] 1 ∆[B] 1 ∆[C] 1 ∆[D] a ∆t b ∆t c ∆t d ∆t ===

14.  The absorption of electromagnetic radiation by a substance at a particular wavelength is directly proportional to its concentration.

15.  The quantitative relationship between rate and concentration is expressed by rate law:  Rate = k [reactant 1] m [reactant 2] n...  The constant k is the rate constant  The exponents m, n, etc are the reaction orders for the reactants.  The sum of the reaction orders gives the overall reaction order.

16.  The exponents in a rate law determine how the rate is affected by the concentration of each reactant.  The values of these exponents must be determined experimentally  In most rate laws, reaction orders are 0, 1, and 2.

17.  Rate laws can be used to determine the concentration of reactants or products at any time during a reaction.  A first-order reaction is a reaction whose rate depends on the concentration of a single reactant raised to the first power.  This form of a rate law is called a differential rate law.

18.  This uses the following calculation:  ln[A] t - ln[A] 0 = -ktOR  ln = - kt [A] t [A] 0

19. m is the slope and b is the y-intercept. A reaction that is not first order will not yield a straight line.

20. Kinetic data for the conversion of methyl isonitrile. In (b) above, the plot of the natural log as a function of time yields a straight line, which confirms that it is a first order reaction.

21.  A second order reaction is one whose rate depends on the reactant concentration raised to the second power or on the concentrations of 2 different reactants, each raised to the first power.

22.  Rate = = k[A] 2  = kt + OR = kt  One way to tell the difference between first- and second-order rate laws is to graph both ln[A] t and 1/[A] t. If the ln[A] t plot is linear, the reaction is first order. If the 1/[A] t plot is linear, the reaction is second order. ∆[A] ∆t 1 [A] t 1 [A] 0 1 [A] t 1 [A] 0

23.  Half-life of a reaction, t 1/2 is the time required for the concentration of a reactant to reach ½ of its initial value, [A] t1/2 = ½[A] 0.  In a first-order reaction, the concentration of the reactant decreases by ½ in each of a series of regularly spaced time intervals, t ½.

24. Half Life of a First Order Reaction

25.  The rate of most chemical reactions increase as the temperature increases.

26.  Reactions occur as a result of collisions between molecules.  This is why the magnitude of rate constants increase with increasing temperature.  The greater the kinetic energy of the colliding particles, the greater the energy of the collision.

28.  Activation energy is the minimum amount of energy needed for a reaction to occur. EaEa  A collision with energy E a or greater can cause the atoms of the colliding molecules to reach the activated complex, or transition state.  The activated complex is a particular arrangement of atoms at the top of the barrier (or hill.)

29. Reaction Coordinate Potential Energy A.E. Products Reactants

30. Reaction Coordinate Potential Energy Product Reactant Activated Complex

31. Reaction Coordinate Potential Energy Product Reactant ΔH

34.  Because the kinetic energy of molecules depends on the temperature, the rate constant of a reaction is very dependent on temperature.  This relationship is given by the Arrhenius equation: ln k = -+ ln A E a RT k is the rate constant, E a R = gas constant, 8.314 J/mol-K T is the absolute temperature

35.  The term A is called the frequency factor.  It relates to the number of collisions that are favorably oriented for a reaction.  It remains nearly constant as the temperature varies.  *Reaction rates decrease as E a increases.*

36.  Reaction mechanisms detail the individual steps that occur in a chemical reaction.  Each step, called an elementary reaction, has a well-defined rate law that depends on the number of molecules (molecularity) in the step.  These are either unimolecular, (1 reactant molecule), bimolecular (2) or termolecular, 3.  The follow rate laws: unimolecular follows first order overall, etc.

38.  Multistep reactions involve two or more elementary reactions, or steps.  An intermediate that is produced in one elementary step is consumed in another later step and doesn’t show up in the overall equation.

39.  The slowest elementary step is the one that limits the speed of the overall reaction and is called the rate-limiting step, or rate determining step.  The rate determining step governs the rate law for the overall reaction.

40.  A catalyst is a substance that changes the speed of a chemical reaction without taking part in the chemical process itself.  It provides an alternate pathway for the reaction to occur.

41.  A homogeneous catalyst is a catalyst that is present in the same phase as the reacting molecules.

42.  A heterogeneous catalyst is a catalyst that exists in a different phase from the reactant molecules, like a solid catalyst involved with gaseous reactant molecules.

43.  Biological catalysts that are necessary for many large inter-related chemical reactions that occur in body systems are called enzymes.  These body system reactions must occur at suitable/specific rates in order to sustain life.

44.  Active site – the specific location on the enzyme where the reaction occurs.  Substrate - the substance that undergoes a reaction at the active site.  Lock and key model – a model that shows how enzyme actions occur.