Multistep & Steady State Analysis Flashcards
How does an chemical reaction normally proceed?
It is rare for an overall reaction to go directly from reactants to final products
Proceed via many elementary reactions
- some act in parallel (competition)
- some act in sequence, intermediate is formed
What is the rate constant in a sequential reactions?
There is a rate constant for every step of the reaction (eg k1 and k2)
What does k1»k2 mean ?
Relate it to the reaction
A -> B -> C
This means the first step is much quicker than the second step
This means there is a considerable time lag between the loss of A and the formation of C
The first reaction takes place quickly and the intermediate B builds up to a concentration close to A before reacting to form C
What does k2» k1 mean?
Relate it to the reaction
A -> B -> C
This means that the second step is much quicker than the first step
B reacts rapidly once it is formed so it’s concentration is always small
This means the formation it actually closely follows the decay of A
How do you write a differential rate equation for multi step reactions?
- for each species identify which reactions it occurs in
- rate of change of its conc will have a term for reach reaction
Each term will depend on the rate constant for the reaction multiply by the conc of all reactants for that reaction
Sign depends on consumption or formation
For the reaction:
A-> (k1) B
B -> (k2) C
Give the differential rate equation with respect to B
d[B]/dt= k1[A] - k2[B]
B is involved in two reactions so both k terms need to be included
In the first reaction it is produced and A is the reactant so you get +ve k1[A]
In the second reaction B is consumed and B is the reactant so you get -k2[B]
What does the rate equation include?
The rate equation includes the rate constants and the concentration of the reactants in the reaction
What assumptions can you make when the second reaction is much faster than the first reaction for
A -> B -> C
The rate of formation of C approximately matches the rate of decay of A
This means the rate of consumption effectively matches the rate of formation so
d[B]/dt= k1[A] - k2[B]=0 -B is used up as soon as it is formed (the change in concentration of B is 0)
This is the steady state hypothesis
What is the steady state hypothesis?
This states that the net rate of change of a highly reactive intermediate is effectively zero
This is because the intermediate is used up as quickly as it is formed so it’s concentration stays constant
What can steady state hypothesis be used for?
It can be used to estimate the concentration of intermediates
d[B]/dt= k1[A] - k2[B]=0
Rearranging this gives:
[B]= k1/k2[A]
So from this you can work out the concentration of B
What does the ratio of products formed depend on?
The ratio of products formed from each reaction depends on the relative magnitude of rate constants
What is a competing first order reaction?
This is where two reactions both have the same reactant so they compete with each other
Eg A-> B
A-> C
What is the differential rate equation with respect to A and what is k total ?
A-> B(k1)
A-> C(k2)
d[A]/dt= -k1[A]-k2[A]
K total is: -(k1+k2)
- the two rate constants added together
How does [A] vary with time in a competing first order reaction?
d[A]/dt= -k1[A]-k2[A]
1) define a total rate constant k= k1+ k2
2) rewrite rate equation
d[A]/dt= -ktotal [A]
3) integrate this with limits to get [A]=[A]oe^-ktotalt
4) proportion going to [B] is k1/ktotal (same for c) work out k total
How do you express reverse reactions?
K1 for forward and k-1 for reverse
