UNIT 2 Flashcards
rates of reaction can be determined using the two formulas
A->B
rate of production of B = Δ[B]/Δt
rate of consumption of A = -Δ[A]/Δt
rate =
change in concentration/change in time
The General Reaction Rate
for a balanced reaction:
aA -> bB + cC
rate = -1/a(Δ[A]/Δt) = 1/b(Δ[B]/Δt) = 1/c*(Δ[C]/Δt)
Initial instantaneous rate vs instantaneous rate
initial instantaneous rate is at t=0.
instantaneous rate can be at any point on the graph
describe the rate law in words
an equation that shows how the reaction rate depends on the concentration of each reactant
Rate Law
aA + bB -> cC + dD
rate = k[A]^m[B]^n
[A], [B] - concentration in M
k - rate constant
m - reaction order in A
n - reaction order in B
what si the overall order of a reaction?
the sum of the individual orders
units of rate
always M/s
units of k
depends on the rate law expression
units of k if:
rate = k
rate = k[A]
rate = k[A][B]
rate = k[A][B]^2
M/s
1/s
1/(Ms)
1/(M^2s)
how are the order of reactions (m, n) and the rate constant (k) determined?
experimentally
give 2 ways of determining order of reaction and rate constant
- method of initial rates:
- vary initial concentration of one reactant at a time
- measure the initial reaction rate for each reaction
- solve a system of rate law equations to determine the order - graphical method/integrated rate law
- monitor the course of a reaction over time
- plot data
- the shape of the curve reveals reaction order
0th order reaction
A –(k)–> product
[A]t = -kt + [A]0
[A]0 is the initial concentration
[A]t is the concentration at time t
- plot the experimental [A] vs time
- if the graph is linear, the reaction is zero-order
- rate constant k = (-slope)
half life
the time required for reactant concentration to reach half of its original value
half life for a 0th order reaction
t(1/2) = [A]0/2k
- dependent on the (initial) concentration
- gets shorter over the course of the reaction (each successive half-life is half as long)
how are 0th order reactions possible?
in reactions whether the kinetics are governed by the availability of a catalyst
1st order reaction
A –(k)–> products
[A]t = [A]0e^-kt
ln[A]t - ln[A]0 = -kt
- plot the experimental ln[A] vs time
- if the graph is linear, the reaction is first-order
- the rate constant k = (-slope)
half life of 1st order reaction
t(1/2) = ln(2)/k
- t(1/2) is independent of the concentration
- every half-life is an equal period of time
2nd order reaction
A –(k)–> products
1/[A]t = kt + 1/[A]0
- plot the experimental 1/[A] vs. time
- if the graph is linear, the reaction is second-order
- the rate constant k = slope
half life of 2nd order reaction
t(1/2) = 1/k[A]0
- t(1/2) is dependent on the (initial) concentration
- t(1/2) gets longer over the course of the reaction, so each successive t(1/2) doubles in length
draw graphs for zeroth order, first order, and second order. also draw straight-line plots used to determine rate constant
how can we find the overall reaction from the elementary steps?
the elementary steps in the true reaction mechanism must sum to give the overall (stoichiometric) reaction
reaction intermediate
a species formed in one step of a reaction mechanism and consumed in a later step
molecularity
a classification of an elementary reaction based on the number of molecules (or atoms) on the reactant side of the chemical equation
unimolecular reaction
an elementary reaction that involves a single reactant molecule
eg rate = k[O3]
bimolecular reaction
an elementary reaction that results from an energetic collision of two reactant molecules
eg rate = k[O3][O]
termolecular reaction
an elementary reaction that involves three atoms or molecules (rare)
eg rate = k[O]^2[M]
how can you find a plausible mechanism for a reaction?
- elementary steps must add up to the overall reaction
- mechanism must correlate with the experimental rate law
the rate law of the overall reaction must be the rate law of
the rate limiting step
how do you determine the rate law for a reaction where the rate limiting step is the second step?
factors affecting reaction rates
- concentration; increased concentration of reactants = increased rate
- moving molecules are closer together
- increased likelihood that reactants will collide - temperature: observed increase in rate wen there is an increase in temperature
- due to collision and transition state theory
- molecules collide with energy greater than Ea more often - catalyst: a catalyst provides an alternative reaction mechanism that proceeds faster than the original (uncatalysed) mechanism
collision theory
for a bimolecular reaction to take place, reactants A and B must collide with proper orientation, and an energy greater than the activation energy, Ea
k =
Zpf
Z = volumetric collision frequency
p = fraction with correct orientation
f = fraction with sufficient energy
k depends strongly on temperature - so which of Z, p, and f depend on temperature?
collision frequency (Z)
- increase in temperature -> reactants collide more often (but not that more often)
fraction with correct orientation (p)
- increase T -> no change
fraction with sufficient energy (f)
- increase T -> reactants collide with greater kinetic energy
transition state
the unstable group of atoms that are the highest-energy species along the pathway from reactants to products
activation energy
the minimum energy required for a successful reaction
f =
e^(-Ea/RT)
as T increases, how does f change?
increases exponentially
draw a graph of collision energy vs fraction of collisions for two temperatures
Arrhenius equation
k = Ae^(-Ea/RT)
A = zp : the frequency factor (aka pre-exponential factor)
ln(k) = -Ea/R (1/T) + lnA
ln(k) = y
-Ea/R = m
1/T = x
lnA = b
Arrhenius plot
y intercept = ln(A)
slope = -Ea/R
catalysts
a substance that provides an alternative reaction mechanism that has a lower activation energy than the uncatalyzed mechanism -> faster reaction
- usually involved in the rate-limiting step of the new pathway
- often appears in the rate law of the catalysed reaction
- react early in a multi-step mechanism; regenerated in a later step, so not consumed
- do not show up in balanced, overall reaction
- catalysed reaction has the same endo/exothermicity as the uncatalysed reaction
Michaelis-Menten kinetics
enzymes are catalysts of biological organisms, typically protein molecules with large molecular weights
at low substrate concentration
- first order behaviour: rate increases linearly with [S]
at high substrate concentration
- zeroth order behaviour: once all the enzyme is complexed, rate of reaction saturates
define chemical equilibrium
the state reached when the concentrations of reactants and products remain constant over time
- forward and reverse reactions occur at the same rate
- reactions are still occurring, but there is no net reaction
- chemical equilibria are dynamic and reversible
Kc =
aA + bB <–> cC + dD
Kc = [C]^c[D]^d/[A]^a[B]^b
concentrations at equilibrium
Kc when you reverse the reaction
1/Kc
Kc when you multiply the reaction by 2
Kc^2
Kc when you add two reactions
Kc1 x Kc2
units of the equilibrium constant
equilibrium constants are defined as unites quantities
homogeneous equilibria
all species are present in the same phase.
heterogeneous equilibria
system’s state of equilibrium contains components from multiple phases
why are pure solids and liquids excluded from equilibrium expressions?
they are always in their standard state, where a=1: their concentrations don’t change
equilibrium constant Kp
for gases, we can use pressures instead of concentrations
Kp = Kc(RT)^Δn
Kp = Pc^c x Pd^d/Pa^a x Pb^b
Δn is the number of moles of gaseous products minus the number of moles of gaseous reactants
Kc > 10^3
reaction proceeds nearly to completion, products favoured
Kc < 10^-3
reaction proceeds hardly at all, reactants favoured
10^-3<Kc<10^3
appreciable concentrations of both reactants and products
reaction quotient, Qc
defined similarly to an equilibrium constant Kc except that the []’s in Qc can have ant values, not necessarily equilibrium values