ch 5 - Chemical Kinetics Flashcards

1
Q

Change in Gibbs Free Energy (delta G)

A

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

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2
Q

collision theory of chemical kinetics

A

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)

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3
Q

activation energy (E sub a)

A

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

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4
Q

rate of reaction equation

A

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

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5
Q

Arrhenius equation

A

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

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6
Q

frequency factor

A

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

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7
Q

transition state theory

A

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

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8
Q

reaction coordinate

A

traces the reaction from reactants to products

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9
Q

transition state

A

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

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10
Q

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

A

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)

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11
Q

factors that can alter experimental rates

A

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

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12
Q

homogeneous catalysis

A

catalyst is in same physical phase as reactants

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13
Q

heterogeneous catalysis

A

catalyst is in a distinct phase

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14
Q

notation of reaction rate

A

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

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15
Q

rate units

A

mol/L x s or M/s

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16
Q

rate law expression

A

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
Q

law of mass action

A

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

18
Q

equilibrium for reversible reaction

A

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
Q

zero-order reaction

A

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
Q

first-order reaction

A

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
Q

radioactive decay rate law

A

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
Q

concentration of radioactive sample A at any time t expression

A

[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
Q

second-order reaction

A

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
Q

mixed-order reactions

A

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
Q

what I am responsible to know about mixed-order reactions

A

large value for [A] at beginning results in k sub 3&raquo_space; 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