BLB 11 th Chapter 14

1. Will the reaction occur? Ch. 5, How fast will the reaction occur? Ch How far will the reaction proceed? Ch. 15

 Review: concentration units, graphing (line, slope)  Work with reaction rates.  Determine a rate law for a given reaction from experimental data.  Work with rate laws.  Understand the factors that affect the rate of a chemical reaction.  Calculate the activation energy.

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

 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. (sect. 6)

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 (section 5) 4. Catalysts: rate ↑ by changing the mechanism and reaction energy (section 7) 5. Other physical things like stirring and grinding solid reactants.

 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:

In this reaction, the concentration of butyl chloride, C 4 H 9 Cl, was measured at various times.

 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.

 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)

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

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 (section 4) Shows how concentration changes with time Graphical determination of the order

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.

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

Units of the rate constant, k

1. What is the order with respect to NO? 2. What is the order with respect to H 2 ? 3. What is the overall order? 4. If [NO] is doubled, what is the effect on the reaction rate? 5. If [H 2 ] is halved, what is the effect on the reaction rate? 6. What are the units of k?

1. Calculate the rate of reaction when the concentration of PtCl 2 (NH 3 ) 2 is 0.020M. 2. What is the rate of Cl¯ production under these conditions?

(b) Calculate rate when [NO] = M and [H 2 ] = M.

Initial Rates Method 1.Find two experiments in which all but one reactant’s concentration is constant. 2.Observe the relationship between concentration change and rate change to determine the order for that reactant. 3.Repeat for other reactant(s).

Exp. #[NH 4 + ] o [NO 2 ¯] o Initial Rate (M·s -1 ) x x x Determine the rate law and calculate k.

Exp. #[NO] o [Cl 2 ] o Initial Rate (M·min -1 ) Determine the rate law and calculate k.

Exp. #[BrO 3 ¯] o [Br¯] o [H + ] o Initial Rate (M·s -1 ) x x x x Determine the rate law and calculate k.

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

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

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

Time (s) Pressure CH 3 NC (torr) , , , , ,

Time (min) [C 12 H 22 O 11 ]

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

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

1 st order reaction

 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

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

BLB 39

To Nuclear Chemistry, Ch. 21

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

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.

Molecular Collisions

 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

 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

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

 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.

y = m x + b

 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

Molecularity is the number of molecules reacting.

 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  Examples: ◦ Catalytic converter (p. 608) ◦ Enzymes in the body (pp ) ◦ Ozone depletion (pp )

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