1. Prepared by PhD Halina Falfushynska

2. C(s, diamond) C(s, graphite) ΔH ° rxn = Is the reaction favorable?

3.  studies the rates at which chemical reactions occur.  gives information about how the reaction occur, that is, the reaction mechanism  Determining the reaction mechanism is the overall goal of kinetic studies.

4. 1. Physical state of the reactants – states that promote contact have faster rates; homogeneous vs. heterogeneous 2. Concentration of the reactants: conc. ↑, rate ↑ (or pressure for gases) 3. Temperature: temp. ↑, rate ↑ due to higher molecular energy and speed 4. Catalysts: rate ↑ by changing the mechanism and reaction energy 5. Other physical things like stirring and grinding solid reactants.

5.  Rate – change in some variable per unit time  Reaction rate – change in concentration per unit time; M/s or mol/L·s  Rates are determined by monitoring concentration as a function of time.  Rates are positive quantities; for reactant A:

6.  Rates change over time: ◦ reactant rates decrease ◦ product rates increase  Instantaneous rate – rate at a specific time  Average rate – Δ[A] over a specific time interval  Initial rate – instantaneous rate at t = 0  Note: Rates and rate laws are not based on stoichiometry!! They must be determined experimentally.

9.  The molar ratios between reactants and products correspond to the relative rates of the reaction.  Relative rates – relationship between rates of reactant disappearance and product appearance at a given time. 2 HI(g) → H 2 (g) + I 2 (g)

10. 2 NO 2 (g) → 2 NO(g) + O 2 (g) Rate = 2.4 x 10 -5 M/s Rate = 8.6 x 10 -5 M/s Rate = 4.3 x 10 -5 M/s

12. 1. Differential rate law or rate law (section 3) Shows how the reaction rate changes with concentration Must be specific in how defined 2. Integrated rate law 3. Shows how concentration changes with time Graphical determination of the order

13. aA + bB → cC + dD  General form of rate law: [A], [B] – conc. in M or P rate = k[A] m [B] n k – rate constant; units vary m, n – reaction orders  Reaction orders and, thus, rate laws must be determined EXPERIMENTALLY!!!  Note: m ≠ a and n ≠ b  Overall order = sum of individual orders  Rate constant is independent of concentration.

14. Order (m)Δ[A] by a factor of:Effect on rate Zero (0)2, 4, 15, ½, etc.None 1 st (1) 22X 33X 2 nd (2) 24X 39X ½¼X

15.  1 st order y = mx + b Plot: ln[A] vs. t slope = − k integrate

16.  2 nd order y = mx + b Plot: 1/[A] vs. t slope = k integrate

17.  zero order y = mx + b Plot: [A] vs. t slope = − k integrate

18. Time (s) Pressure CH 3 NC (torr) 0502 2,000335 5,000180 8,00095.5 12,00041.7 15,00022.4

19. Time (min) [C 12 H 22 O 11 ] 00.316 390.274 800.238 1400.190 2100.146

20.  t 1/2 – time required for the concentration of a reactant to decrease by half of its initial value

21.  1 st order half-life ◦ All half-lives same length of time ◦ Independent of initial concentration Note: Radioactive decay follows 1 st order kinetics.

22. 1 st order reaction

23.  2 nd order half-life ◦ All half-lives different length of time ◦ Dependent on initial concentration  Zero half-life ◦ All half-lives different length of time ◦ Dependent on initial concentration

24. Order zero1 st 2 nd Rate lawrate = krate = k[A]rate = k[A] 2 Integrated rate law [A] t =−kt+[A] 0 ln[A] t =−kt+ln[A] 0 1/[A] t =kt+1/[A] 0 Straight- line plot [A] vs. tln[A] vs. t1/[A] vs. t Slope−k−k−k−kk Half-life (t 1/2 ) [A] o /2k0.693/k1/k[A] 0

25.  Generally, as temperature increases, so does reaction rate.  This is because k is temperature dependent.

26. Collision theory  In a chemical reaction, bonds are broken and new bonds are formed. In order for molecules to react, they must collide.  Collisions are either effective or ineffective due to orientation of molecules.  Collisions must have enough energy to overcome the barrier to reaction, the activation energy.  Temperature affects the number of collisions.

27. Molecular Collisions

28.  Energy barrier (hump) that must be overcome for a chemical reaction to proceed  Activated complex or transition state – arrangement of atoms at the top of the barrier

29.  Energy difference between the reactant and the highest energy along the reaction pathway  Reaction specific  Rate of reaction is dependent upon the magnitude of E a ; E a ↓, rate ↑ (generally)  Temperature independent

31.  At higher temperatures, more molecules will have adequate energy to react.  This increases the reaction rate. Maxwell-Boltzman Distri bution

32.  Svante Arrhenius developed an equation for the mathematical relationship between k and E a.  A is the frequency factor, which represents the number of effective collisions.

33. y = m x + b

35.  Reactions occur in a series of elementary steps collectively called a mechanism.  Determining the reaction mechanism is the overall goal of kinetic studies.  One step, the rate-determining step, is much slower than the other.  Usually, an intermediate (isolable) or a transition state (non-isolable) is formed at some point during the reaction.  molecularity – the number of molecules that participate in a reaction

36. Molecularity is the number of molecules reacting.

37.  Catalysts – increase the rate of a reaction without being consumed or changing chemically  Accomplished by lowering the activation energy and changing the reaction mechanism.  Heterogeneous vs. homogeneous catalysis

40. Heterogeneous catalytic ethylene hydrogenation: C 2 H 4 + H 2 → C 2 H 6