7. Kinetics Flashcards
What is a reaction intermediate?
a chemical that is produced in a reaction mechanisms elementary step and then subsequently consumed in another elementary step.
The intermediate does not appear in the overall reaction.
What is the rate-determining step?
In a sequence of elementary steps, the slowest step is the rate-determining step which governs the kinetics of the overall reaction.
Collision theory: what three elements determine reaction rate?
- frequency of reactant chemical collisions
- the orientation of the colliding molecules
- their energy
What is the activated complex?
The activated complex is the transition state of a chemical reaction. On the graph, the activated complex is a local (or global) maxima while the intermediates formed are local minima.
Activation energy is the energy needed to get to the highest transition state.
Collision theory states reactions proceed faster based on collision frequency, orientation, and their energy. What three ways can we speed up a reaction?
- Lowering the activation energy (i.e. less energy is needed for the reaction to proceed)
- increased reactant concentration (increased collision frequency)
- increased temperature (increased energy to overcome the Ea)
note that you cannot really manipulate chemical orientation
What is a catalyst?
Catalysts speed up chemical reactions by lowering the activation energy through stabilizing the transition state (of the RDS). Catalysts are NOT consumed by the reaction and are regenerated at the end of the reaction.
t or f, catalysts effect the direction of a reaction.
FALSE, catalysts only speed up the reaction in whatever way the reaction is occurring (determined by thermodynamics and equilibrium).
Thus, catalysts decrease the Ea of both the forward and reverse reaction.
How does a catalyst change the ΔH of a reaction?
A catalyst would not influence ΔH as this is a thermodynamic property. A catalyst will lower the Ea.
What is the rate law of a reaction? What is a part of it?
Reactions are almost always measured as the rate of reactants disappearing (and thus, you do not see products in a rate law expression). A rate law expression typically only involves reactants of the rate-determining step.
aA + bB –> cC + dD
rate = k[A]^x[B]^y
Explain this.
This is the rate expression for the reaction aA + bB –> cC + dD (assuming that this is the RDS of a reaction).
k = the rate constant
x = reaction order with respect to A
y = reaction order with respect to B
x + y = overall reaction order
t or f, the rate law of a reaction can only be determiend experimentally?
true, rate orders and rate constants are found experimentally.
However, there are exceptions: For elementary steps of a reaction, the rate law is first order for uni-molecular steps and second order for bimolecular steps.
Explain rate orders of 0, 1, and 2.
We can hold the concentrations of given reactants constant and then change one reactant and see how reaction rate changes.
If increasing [] gives no change = 0
If [] and rate are proportional = 1
if doubling [] leads to quadrupling rate = 2
reactants with a rate order of 0 are excluded from the rate law expression (b^0 = 1).
How do you find the rate constant (k) with experimental data?
You first find the individual orders of the reaction. Then you may use any of the experimental data and plug that into the rate law expression, isolating for k.
How does temperature affect the rate constant?
Arrhenius equation: lnK = lnA - (Ea / RT)
Increasing temperature increases the rate constant (larger denominator means we are subtracting something smaller)
Similarly, catalysts decrease Ea which increases the rate constant.
Reaction rate is measured as M / s. Explain how the order of the reaction changes the units of the rate constant?
first order: R = k[A], thus s^-1
second order: R = k[A][B], thus M^-1 s^-1
third order: R = k[A][B]^2, thus M^-2 s^-1
the units become whatever is needed to make the reaction rate M/s
