Most chemical reactions have an activation barrier

A chemical reaction is a significant rearrangement of otherwise stable electron configurations in molecules or atoms. In order for any rearrangement of bonding electrons to take place, one or more of these things have to happen first:

close approach of atoms/molecules, which requires overcoming electron-electron repulsion (Coulomb) forces

bending or stretching of one or more chemical bonds

partial breakage or formation of one or more chemical bonds

Here's a quick example that we'll expand upon below. Consider the reaction in which a chloride ion (Cl-) attacks a methyl-bromide molecule (CH3Br) to displace the bromide and form methyl chloride (CH3Cl)

In the diagram, Cl- has to approach CH3Br, and to approach closely, there has to be enough relative translational energy (they have to be approaching fast enough) to overcome electron-electron repulsion. As a Cl-C bond begins to form, VSEPR theory suggests that the C-H bonds will bend due to repulsion of electron clouds, and finally, one bond (Cl-C) is formed while another (C-Br) is broken. All of these contribute to what we refer to as an energetic "barrier," or an "activation barrier" to the reaction.

Svante Arrhenius (1889), who made several foundational discoveries in chemistry, called this barrier the activation energy. In order for a reaction to occur, the relative energy (kinetic + potential) of the reactants has to at least equal the activation energy.

We often draw a hypothetical reaction coordinate to illustrate what the activation barrier looks like. In the case of an exothermic reaction, in which the energy of the reactants is greater than the energy of the products, that graph looks like this:

In a diagram like this, the reaction proceeds from left to right, from reactants to products. The activation energy, Ea, forms a barrier to the start of the reaction. The reaction must have some input of energy to get it going. If it's spontaneous, no extra input of energy beyond that generated by the reaction itself will be needed to keep it going.

The reaction coordinate of an endothermic reaction might look like this:

It's worth comparing these reaction coordinate graphs. Notice that they are approximately mirror images of each other. In fact, the reverse of an exothermic reaction is an endothermic reaction with a higher activation energy.

Activation energy ↔ rate constant

The activation energy is related to the rate of a chemical reaction. That just makes sense. In one extreme – no activation barrier at all – a reaction occurs very rapidly. And in the limit in which the activation energy is very high, the reaction should be very slow.

Arrhenius related the rate constant to the activation energy through the Arrhenius equation:

k is the rate constant (not the rate itself, but always linearly proportional to it), A is a "frequency factor," related to the frequency of collisions which lead to reactions, R is the molar gas constant (R = 8.314 J·mol-1·K-1) and T is the Kelvin temperature.

Because the exponential part of the function is in the denominator (negative exponent), we expect the rate constant to shrink exponentially with activation energy, but for that to be mitigated by raising the temperature. In fact. all reaction rates increase as heat is added.

The Arrhenius equation will be explored in more detail in another section.

The transition state

Let's take another look (see diagram below) at a single-displacement reaction. This time a hydroxyl ion (OH-) will attack bromomethane (CH3CH2Br) to displace the bromine (forming a Br- ion) and ethanol (CH3CH2OH). It's very similar to the methyl bromide ↔ methyl chloride reaction in the previous section. The reaction is

OH- + C2H5Br ⇌ C2H5OH + Br-

Arrhenius noted that there must be some transient structure which the reacting compounds form that is unstable, and could, with equal probability, fall apart into either reactants or products.

That structure, the transition state, is shown below. It is commonly marked with the "double dagger" symbol, .

The transition state structure is elusive, and defies most experimental attempts to follow a reaction through from start to finish. Nevertheless, we can often infer what the transition state must look like from other experimental data, including the details of the reaction mechanism and its rate and thermodynamic profile.

The activation energy of a reaction is the energy required to create the transition-state structure or the transition complex, a relatively high-energy structure.



means lasting for only a very short time.

Bromoethane and ethanol are stable products of either the forward (shown) or reverse reaction. The transition state is less stable and of high energy. Once it is formed, there is roughly equal probability of it falling apart either into reactants or products.

The role of a catalyst

A catalyst can come in many forms, and we'll get to some of those in a bit, but the role of a catalyst is simple: it alters the transition state to reduce its energy, thus making the barrier to the reaction lower.

Here's an example of how a catalyst might alter the activation energy of an exothermic reaction:

The green curve shows how a catalyst might modify the reaction energy curve to reduce the activation energy, Ea.

The corresponding curve for an endothermic reaction looks like this:

It's just that simple. Catalysts reduce the activation energy by modifying the form of the transition state. Now let's look at some examples of how that might happen.

Heterogeneous catalysis

One of the simplest types of catalysis to understand is heterogeneous catalysis. Heterogeneous means "of different character," and here it refers to the fact that the reactants and products are of a different physical phase than the catalyst itself. A nice example is the catalytic converter in a gasoline-powered car.

The reason the catalytic converter is there is to convert toxic, smog-forming carbon monoxide (CO) to carbon dioxide (CO2) by combination with oxygen in the reaction

2 CO (g) + O2 (g)   ⇌   2 CO2 (g)

In the gas phase and at fairly low partial pressures, this reaction proceeds exceedingly slowly because the activation energy is high and there are pretty stringent requirements on the alignment of an already improbable three-body collision:

To speed this reaction along (substantially), we employ metals like platinum (Pt) and Rhodium (Rh). These metal surfaces consist of metal atoms packed in a tight crystal-like lattice, as in the cartoon below (only the surface atoms are shown).

What essentially happens on metal surfaces like this in the presence of CO and O2, is that the CO molecules adsorb onto the metal surface, and are free to migrate around on it afterward. When two CO molecules of the same end-to-end alignment adsorb next to one another, they are aligned perfectly for the next collision with an O2 molecule, and the reaction proceeds rapidly from there. The role of the metal is to organize and align the CO molecules so that improbable 3-body gas-phase collisions aren't required.

Because of the way they are constructed (see diagram below), catalytic converters can be close to 100% efficient in terms of the gas molecules coming out of the exhaust pipe of a car.

Schematic diagram of an automobile catalytic converter. The engine exhaust is passed through a chamber containing a Pt, Pd and/or Rh adsorbed onto a zeolite, an inert material with a great deal of surface area. At exhaust-gas temperatures, the converter approaches 100% conversion efficiency.

A heterogeneous catalyst is not of the same phase as the reaction it catalyzes. Most heterogeneous catalysts are solids which catalyze liquid or gas-phase reactions.



Adsorption describes the sticking of a gas or liquid onto a solid surface, usually weakly.

Absorption is migration of a gas or liquid into a solid, such as a sponge or a sieve.

Homogeneous catalysis

A homogeneous catalyst is one that is of the same phase (gas, liquid, solid) as the reaction it catalyzes.

The most common kind of homogeneous catalysis is acid catalysis. Here's an example, the hydrolysis of the ester methyl acetate (or methyl ethanoate).

This reaction relies on free protons from the autoionizaton of water. You can see the reaction mechanism in the wide figure below. the auto ionization of water is

H2O   ⇌   H+ + OH-

The equilibrium constant expression for it is

which means that the concentration of free protons is only 1×10-7 M, a very small concentration that causes this reaction to proceed only very slowly in pure water.

But the addition of acid causes causes the proton concentration to rise, hastening the reaction. For every proton used to form the acetic acid product, a water ionizes, liberating a hydroxyl ion to form the methanol product. Charge neutrality demands dissociation of the water molecule.

Because the proton supply is replenished in the reaction, the acid doesn't, strictly speaking, take part in it – its concentration is unchanged. That's one definition of a catalyst: something that speeds up the reaction without being changed by the reaction itself.

In the reaction below, an acidic proton attacks the oxygen that connects the two parts of the ester molecule to form the acid (acetic acid in this case). That proton is replenished by the ionization of a water molecule, which also liberates an OH- ion. That ion binds to the other half of the ester to form methanol.



An ester is an organic (carbon-containing) molecule that consists of two organic elements, R1 and R2 below, joined by an ester group (O=C-O):


Our last class of catalyst is the most important for living things, enzymes. Enzymes are protein molecules (or sometimes protein / RNA complexes) that catalyze biological reactions.

While they're really in a class by themselves, enzymes can be considered to be heterogeneous catalysts. They are large molecules that present shaped surfaces or pockets on surfaces, into which reactants fit, and are pre-aligned to facilitate a reaction.

A good example is the ribosome, a very large protein + RNA complex that catalyzes the building of proteins from the code in DNA. The ribosome itself is built of more than half a million atoms. The picture gives a broad overview.

There is a groove into which the messenger RNA (mRNA) fits and can move along its axis. Other pockets contain transfer RNAs (tRNAs), which are attached, via the genetic code, to one of 20 amino acids, the building blocks of proteins.

The mRNA is "ratcheted" through its groove. At each stop, the appropriate tRNA aligns with the mRNA code,

and is then in position to add its amino acid to the growing protein chain. It's an amazing reaction that would proceed in geological time (VERY slowly) without the enzyme. With it, proteins of 100s of amino acids can be built in seconds.

Nearly every chemical reaction that occurs in living systems involves some specific enzyme.

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