Module 2 Flashcards
Rate of Reaction
Change in amount of a reactant / product
Rate of Reaction Equation
Δconcentration / Δtime
How do you calculate the rate of appearance/disappearance of reactants/products? (aA+bB -> cC+dD)
-1/a * Δ[A]/Δt = -1/b * Δ[B]/Δt = 1/c * Δ[C]/Δt = 1/d * Δ[D]/Δt
or
-1/a * rate[A] = -1/b * rate[B] = 1/c * rate[C] = 1/d * rate[D]
Ex: Rate of disappearance of A = a/b * rate[B] = a/c * rate[C] = a/d * rate[D] **Always Positive!
Factors that affect reaction rates
- Effective Collisions (Correct orientation + KE > Activation E)
- Concentration
- Temperature
- Surface Area
- Presence of a Catalyst
NOT Pressure
Do heterogeneous or homogenous solutions react faster?
Homogenous solutions react faster
Rate Law
Rate = k[reactant] ^ rxn order
Ex: aA+bB->cC+dD
R = k[A]^x * [B]^y
Reaction Order
The sum of exponents of concentrations,
Ex: rate = [A]^2 * [B]^3 Overall Rxn Order = 5
Rxn Order in respect to A = 2
Reaction Rate Units
Molarity / second
How do you experimentally determine the rate law
rate2/rate1 = ([A2]/[A1])^n
Rate Laws for different orders
Order: 0 »_space; rate = k
1»_space; rate = k[A]
2»_space; rate = k[A]^2
Integrated Rate Laws
0 order – [A]f = -kt + [A]i
1 order – ln[A]f = -kt + ln[A]i
2 order – 1/[A]f = kt + 1/[A]i
Half-Life Equations
0 order – t=[A]i/2k
1 order – t=(ln2)/k
2 order – t=1/k[A]i
Integrated Rate Law Units
0 – M/s
1 – 1/s
2 – 1/M*s
Activation Energy
Required energy for a reaction to occur
Arrhenius Equation
k = Ae^(-Ea/RT)
ln(k1/k2) = Ea/R * (1/T2 - 1/T1)
A = Frequency Factor
Ea = Activation Energy
T = Temp in Kelvin
R = 0.008314
