On this page:
Reaction mechanisms, molecularity Collision theory of chemical change
Anatomy of moleclar collisions Activation energy Catalysts
Temperature and kinetic energy
The Arrhenius law
Deterining the activation energy The pre-exponential factor
What you should be able to do Concept map
General Chemistry Virtual Textbook → kinetics/dynamics → collision /activaton
Collision and activation
the Arrhenius Law
index  | rate  laws  1 | rate  laws  2 | activation  | mechanisms  | solutions  | catalysis  | experimental
Note: this document will print  in an appropriately modified format (15 pages)
Why are some reactions so much faster than others, and why are reaction rates independent of the thermodynamic
tendency of the reaction to take place?These are the central questions we
address in this unit. In doing so, we open the door to the important topic of
reaction mechanisms : what happens at the microscopic level when chemical reactions take place? We can thank Prof. Svante Arrhenius  for unlocking this door!
To keep things as simple as possible, we will restrict ourself to reactions that take place in the gas phase.The same principles will apply to reactions in liquids and solids, but with added
complications that we will discuss in a later unit.
1  Reaction mechanisms
The mechanism  of a chemical reaction is the sequence of actual events that take place as reactant molecules are converted into products. Each of these events constitutes an elementary step  that can be represented as a coming-together of discrete particles ("collison") or as the breaking-up of a molecule ("dissociation") into simpler units. The molecular entity that emerges from each step may be a final product of the reaction, or it might be an intermediate  — a species that is created in one elementary step and destroyed in a subsequent step, and therefore does not appear in the net reactio
n equation.
Step
A reaction mechanism must ultimately be understood as a "blow-by-blow"description of the molecular-level events whose sequence leads from reactants to products. These elementary steps (also called elementary reactions ) are almost always very simple ones involving one, two, or [rarely] three chemical species which are classified, respectively, as
unimolecular    A →by far the most common bimolecular    A + B →
termolecular
A +
B +
C →
very rare
In a gas at room
temperature and normal atmospheric pressure,there will be about 1033collisions in each cubic centimetre every second.If every collision between two reactant molecules yielded products, all reactions would be
complete in a fraction of a second.
2  Collision theory of chemical change
Molecules must collide before they can react
This fundamental rule must guide any analysis of an ordinary chemical reaction mechanism.
This explains why termolecular processes are so uncommon. The kinetic
theory of gases tells us that for every 1000 binary collisions, there will be only one event in which three molecules simultaneously come together. Four-way collisions are so improbable that this process has never been demonstrated in an elementary reaction.
Consider a simple bimolecular step
A +
B → products
Clearly, if two molecules A and B are to react, they must approach closely enough to disrupt some of their existing bonds and to permit the creation of any new ones that are needed in the products. We call such an encounter a collision .
The frequency of collisions between A and B in a gas will be proportional to the concentration of each; if we double [A], the frequency of A-B
collisions will double, and doubling [B] will have the same effect. So if all collisions lead to products, than the rate of a bimolecular process will be first-order in A and B, or second-order overall:
rate = k [A][B]
<
Not all collisions are equal
When two billiard balls collide, they simply bounce off of each other. This is also the most likely outcome if the reaction between A and B requires a significant disruption or rearrangement of the bonds between their atoms.In order to effectively initiate a reaction, collisions must be sufficiently energetic  (kinetic energy) to bring about this bond disruption. More about this further on.
And there is often one additional requirement. In many reactions,especially those involving more complex molecules, the reacting species must be oriented in a
react to domanner that is appropriate for the particular process. For example, in the gas-phase reaction of
dinitrogen oxide with nitric oxide,the oxygen end of N 2O must hit the
nitrogen end of NO; reversing the
orientation of either molecule prevents the reaction.
Owing to the extensive randomization of molecular motions in a gas or liquid,there are always enough correctly-oriented molecules for some of the molecules to react. But of course, the more critical this orientational requirement is, the fewer collisions will be effective.
Anatomy of a collision
Energetic collisions between molecules cause interatomic bonds to stretch and bend farther, temporarily weakening them so that they become more
susceptible to cleavage. Distortion of the bonds can expose their associated electron clouds to interactions with other reactants that might lead to the formation of new bonds.
Chemical bonds have some of the properties of
mechanical springs, whose potential energy
depends on the extent to which they are
stretched or compressed. Each atom-to-atom
bond can be described by a potential energy
diagram that shows how its energy changes
with its length. When the bond absorbs energy
(either from heating or through a collision), it is
elevated to a higher quantized vibrational state
(indicated by the horizontal lines) that weakens
the bond as its length oscillates between the
extended limits corresponding to the curve.
A particular collision will typically excite a number of bonds in this way. Within about 10–13 second this excitation gets distributed among the other bonds in the molecule in rather complex and unpredictable ways that can concentrate the added energy at a particularly vulnerable point. The affected bond can stretch and bend farther, making it more susceptible to cleavage. Even if the bond does not break by pure stretching, it can become distorted or twisted so as to expose nearby electron clouds to interactions with other reactants that might encourage a reaction.
Consider, for example, the isomerization of cyclopropane to propene which takes place at fairly high temperatures in the gas phase.
We can imagine the collision-to-product sequence in the following [grossly oversimplified] way:
Note that
To keep things simple, we do not show the hydrogen atoms here. This is
reasonable because C–C bonds are weaker then C–H bonds and thus less likely to be affected.
The collision at  will usually be with another cyclopropane molecule, but
because no part of the colliding molecule gets incorporated into the product, it can in principle be a noble gas or some other non-reacting species;
Although the C–C bonds in cyclopropane are all identicial, the instantaneous
localization of the collisional energy can distort the molecule in various ways (
), leading to a configuration sufficiently unstable to initiate the rearrangement to the product.
Unimolecular processes also begin with a collision
The cyclopropane isomerization described above is typical of many decomposition reactions that are found to follow first-order kinetics, implying that the process is unimolecular. Until about 1921, chemists did not understand the role of collisions in unimolecular processes. It turns out that the mechanisms of such reactions are really rather complicated, and
The chemical reactions associated with most food spoilage are catalyzed by enzymes produced by the bacteria which mediate
these processes.
Here is a short YouTube video  on activation energy.
The "reaction coordinate " plotted along the abscissa represents the changes in atomic coordinates as the
system progresses from reactants to products. In the very simplest elementary reactions it might correspond to the stretching or twisting of a particular bond, and be shown to a scale. In general, however, the reaction coordinate is a rather abstract concept that cannot be tied to any single measurable and scaleable quantity.
The activated complex  (also known as the transition state ) represents the structure of the system as it exists at the peak of the activation energy curve. It does not correpond to an identifiable intermediate structure (which would more properly be considered the product of a separate elementary process), but rather to whatever configuration of atoms exists during the collision, which lasts for only about 0.1 picosecond.
Activation energy diagrams always incorporate the energetics (ΔU  or ΔH ) of the net reaction, but it is
important to understand that the latter quantities depend solely on the thermodynamics of the process which are always independent of the reaction pathway. This means that the same reaction can exhibit different activation energies if it can follow alternative pathways.
With a few exceptions for very simple processes,activation energy diagrams are largely conceptual constructs based on our standard collision model  for
chemical reactions. It would be unwise to read too much into them.
that at very low pressures they do follow second-order kinetics. Such
reactions are more properly described as pseudounimolecular . The details are beyond the scope of this course, but a good introduction can be found on this U. Arizona page .
Activation energy
Higher temperatures, faster reactions
It is common knowledge that chemical reactions occur more rapidly at higher temperatures. Everyone knows that milk turns sour much more rapidly if stored at room temperature rather than in a refrigerator, butter goes rancid more quickly in the summer than in the winter, and eggs hard-boil more quickly at sea level than in the mountains. For the same reason, cold-blooded animals such as reptiles and insects tend to be noticeably more lethargic on cold days.
It is not hard to understand why this should be. Thermal energy relates direction to motion at the molecular level. As the temperature rises,
molecules move faster and collide more vigorously, greatly increasing the likelyhood of bond cleavages and rearrangemens as described above.Activation energy diagrams
Most reactions involving neutral molecules cannot take place at all until they have acquired the energy needed to stretch, bend, or otherwise
distort one or more bonds. This critical energy is known as the activation energy  of the reaction. Activation energy diagrams  of the kind shown below plot the total energy input to a reaction system as it proceeds from reactants to products.
In examining such diagrams, take special note of the following:
Gallery of activation energy plots
Activation energy diagrams can describe both exothermic and endothermic reactions:
... and the activation energies of the forward reaction can be large, small, or zero (independently, of course, of the value of ΔH):
Processes with zero activation energy most
commonly involve the combination of
oppositely-charged ions or the pairing up of
electrons in free radicals, as in the
dimerization of nitric oxide (which is an odd-
electron molecule).
In this plot for the dissociation of
bromine, the E a is just the enthalpy
of atomization
Br2(g)→ 2 Br· (g)
and the reaction coordinate
corresponds roughly to the
stretching of the vibrationally-
excited bond. The "activated
complex", if it is considered to
exist, is just the last, longest
"stretch". The reverse reaction,
being the recombination of two
radicals, occurs immediately on
contact.
Where does the activation energy come from?
In most cases, the activation energy is
supplied by thermal energy, either
through intermoleculr collisions or (in
the case of thermal dissocation) by
thermal excitation of a bond-stretching

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