Kinetics
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Rate of conversion (J)
J = -(1/a)(dnA/dt) = -(1/b)/(dnb/dt) A & B = -, C & D = +
Reaction rate (r)
r = J/V = -(1/aV)(dnA/dt)
Order of the reaction
The power to which the concentration of a reactant is raised
Rate Law
r = d[P]/dt = k[A]^a[B]^b, Expressed as a function of the reactant concentrations, set in square brackets, at constant temp
Rate constant
k
Partial orders
a & b
Total order (n)
Sum of the partial orders
First order reactions aA -> products
r = k[A] = -(1/a)(d[A]/dt) [A] = [A]o•e^(-ka•t) ln([A]/[A]o) = -ka•t
First order half life
t(1/2) = ln2/ka
Second order reactions
2 types: aA -> products aA + bB -> products
Second order reactions: type 1
d[A]/dt = -ka[A]^2 ka = a • k
Second order half life:type 1
t(1/2) = 1/([A]o•ka)
Second order reaction type: 2
r = k[A][B] {1/(a[B]o-b[A]o}ln{([B]/[B]o)/([A]/[A]o)} = kt
Nth order reaction
([A]/[A]o)^n-1 = 1+[A]^n-1•(n-1)•ka•t t(1/2) = (2^n-1 - 1)/(n-1)•[A]o^n-1•ka
Zeroth order reaction
r = d[A]/dt = -ka [A] = -Ka•t + [A]o t(1/2) = [A]o/2ka
Reaction mechanism
The process by which the reaction occurs
Reaction intermediates
Species that are formed in one step and consumed in another
Stoichiometric number
The number of occurrences of a step
Elementary reaction
Each step
Initial rate method
Most common way for determining rate laws
Collision theory
Reactants must collide to react. If temperature increases, so does the speed of the molecules and the frequency of collisions
Activation energy
Energy that must be overcome for the reaction to occur
Activated complex
A transition state where kinetic energy is stored as potential energy
Arrhenius’s Equation
k = Ae^(-Ea/RT)