ch 5 - Chemical Kinetics Flashcards Preview

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Flashcards in ch 5 - Chemical Kinetics Deck (25)
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1

Change in Gibbs Free Energy (delta G)

determines whether or not a reaction will occur by itself without outside assistance

2

collision theory of chemical kinetics

states that the rate of a reaction is proportional to the number of collisions per second between the reacting molecules (not all collisions result in a chemical reaction though)

3

activation energy (E sub a)

also called energy barrier; the minimum energy of collision necessary for a reaction to take place

4

rate of reaction equation

rate = Z x f; where Z = total number of collisions occurring per second and f = fraction of collisions that are effective

5

Arrhenius equation

k = Ae^(-E sub a/(RT)); where k = rate constant of reaction; A = frequency factor; E sub a = activation energy of reaction; R = ideal gas constant; T = temp in kelvin; as the frequency factor of the reaction increases, the rate constant increases in a direct relationship

6

frequency factor

also called attempt frequency - a measure of how often molecules in a certain reaction collide, with the unit s^-1

7

transition state theory

when molecules collide with energy equal to or greater than the activation energy, they form a transition state in which the old bonds are weakened and the new bonds begin to form; the transition state then dissociates into products, fully forming new bonds

8

reaction coordinate

traces the reaction from reactants to products

9

transition state

also called activated complex: has greater energy than both the reactants and products. Activation energy is the energy required to reach this state; once this is formed it can either dissociate into products or revert back to reactants with no energy input

10

free energy change of the reaction (delta G sub rxn)

the difference between the free energy of the products and the free energy of the reactants: negative indicates an exergonic reaction (energy given off), positive indicates an endergonic reaction (energy absorbed)

11

factors that can alter experimental rates

reaction concentrations (increased reactants increase rate), temperature (generally increase increases rate until denaturation of catalysts), medium, catalysts

12

homogeneous catalysis

catalyst is in same physical phase as reactants

13

heterogeneous catalysis

catalyst is in a distinct phase

14

notation of reaction rate

A + B -> C; rate of reaction with respect to A = -delta[A]/delta t, B = -delta[B]/delta t; C = + delat[C]/delta t. Reactants have negative signs because they are being consumed

15

rate units

mol/L x s or M/s

16

rate law expression

rate = k[A]^x[B]^y ; k = reaction rate coefficient or rate constant; x and y = order of the reaction; x = order with respect to reactant A; y = order with respect to reactant B; overall order of reaction is sum of x and y

17

law of mass action

the equilibrium constant expression used to determine the intermediate molecule's concentration

18

equilibrium for reversible reaction

K sub eq = the ratio of the rate constant for the forward reaction, k, divided by the rate constant for the reverse reaction k sub -1

19

zero-order reaction

one in which the rate of formation of product C is independent of changes in concentrations of any of the reactants; rate = k[A]^0[B]^0 = k and k has units M/s

20

first-order reaction

rate that is directly proportional to only one reactant, such that doubling the concentration of that reactant results in doubling of the rate of formation of the product: rate = k[A]^1 or =k[B]^1 and k has units s^-1

21

radioactive decay rate law

from the rate law, in which the rate of decrease of the amount of radioactive isotope A is proportional to amount of A: rate = -(delta [A])/delta t = k[A]

22

concentration of radioactive sample A at any time t expression

[A]sub t = [A] sub 0 x e^-kt; [A] sub t = concentration of A at time t, [A] sub 0 = initial concentration of A, k = rate constant; t = time; e = Euler's number

23

second-order reaction

has a rate that is proportional to either the concentrations of two reactants or to the square of the concentration of a single reactant: rate = k[A]^1[B]^1 or rate = k[A]^2 or rate = k[B]^1; k has units of M^-1 s^-1

24

mixed-order reactions

sometimes refer to non-integer orders (fractions) and in other cases to reactions with rate orders that vary over the course of the reaction; fractions are more specifically described as broken-order: rate = (k sub 1 [C][A]^2)/(k sub 2 + k sub 3[A]); A represents single reactant and C is a catalyst

25

what I am responsible to know about mixed-order reactions

large value for [A] at beginning results in k sub 3 >> k sub 2. reaction will appear to be first-order with respect to A. at end when [A] is low k sub 2>> k sub 3[A], making the reaction appear second-order with respect to A