Think of a reaction as a hike from one valley to another over a mountain pass. The valleys are the reactants and products, and the pass is the energy barrier that every successful collision must get over. How high the destination valley sits compared with the start is ; how high the pass is above the start is the activation energy. A potential energy diagram draws exactly this picture, and it explains in one glance why a strongly exothermic reaction (like a match burning) may still need a spark to get started.
Reading a Potential Energy Diagram
The vertical axis shows the potential energy stored in the chemical bonds, and the horizontal axis shows the progress of the reaction. As reactant particles collide, their kinetic energy is converted into potential energy while bonds stretch and begin to break. At the very top of the curve is the activated complex (also called the transition state): an unstable, short-lived arrangement in which old bonds are partly broken and new bonds are partly formed. It can fall forward to products or back to reactants.
- = energy of activated complex energy of reactants.
- = energy of activated complex energy of products.
- = energy of products energy of reactants (negative if the products are lower).
- These three are linked:
A catalyst speeds up a reaction by providing a different pathway with a lower-energy activated complex. On the diagram, the catalyzed curve starts and ends at exactly the same levels, so does not change, but the peak is lower. Because the peak drops for travel in both directions, a catalyst lowers the forward and reverse activation energies by the same amount. A catalyst takes part in the mechanism but is regenerated by the end, which is why only a small amount is needed. Enzymes are biological catalysts that let reactions in your cells run quickly at body temperature.
Reaction Mechanisms
Most balanced equations hide what really happens. The chance of several particles all hitting each other at the same instant, correctly oriented, is almost zero, so most reactions happen as a series of simple collisions called elementary steps. The full sequence of steps is the reaction mechanism. An elementary step involving one particle is unimolecular, two particles bimolecular, and three particles termolecular (rare, because a three-way collision is so unlikely). Adding the steps together must give the overall balanced equation.
- An intermediate is formed in one step and used up in a later step, so it does not appear in the overall equation ( above).
- A catalyst is used up in an early step and regenerated in a later step, so it also cancels from the overall equation — but it was present at the start.
- The rate-determining step is the slowest step, the one with the highest activation energy. The overall reaction cannot go faster than this bottleneck, just as a line at a store moves only as fast as its slowest cashier.
- Because step 1 is the slow step above, the rate depends on how often two molecules collide, and changing has essentially no effect on the rate.
On a potential energy diagram, a two-step mechanism has two peaks with a small valley between them. The valley is the intermediate, and the step with the higher barrier is the rate-determining step. Combustion, the everyday reaction of burning a fuel, is far more complex: even burning methane proceeds through many elementary steps involving highly reactive fragments called free radicals, and the energy released by those steps keeps the chain going, which is why a lit flame sustains itself as long as fuel and oxygen keep arriving.
Energy Changes, Catalysts and Technology
| Technology or process | Energy change or chemistry involved | Efficiency and environmental points |
|---|---|---|
| Fossil-fuel power plants and car engines | Exothermic combustion of hydrocarbons heats steam or expands gases | Much of the released heat is lost rather than converted to useful work; burning releases CO₂, a greenhouse gas, and can release pollutants such as nitrogen oxides |
| Catalytic converters | Platinum, palladium and rhodium catalyze the conversion of CO, nitrogen oxides and unburned hydrocarbons into CO₂, N₂ and H₂O | A small amount of catalyst works for years because it is regenerated; it cuts smog-forming pollutants but not CO₂ |
| Hydrogen fuel cells | Hydrogen and oxygen combine to form water, releasing chemical energy directly as electricity | Water is the only product at the cell, but the overall impact depends on how the hydrogen is made; much hydrogen today comes from natural gas, which releases CO₂ |
| Hydroelectric, wind and solar power | No combustion; energy comes from falling water, moving air or sunlight | No CO₂ from fuel during operation, but dams flood land, and wind and solar output varies with the weather |
| Haber process for ammonia | Nitrogen and hydrogen react over an iron catalyst | The catalyst lets the reaction run at a useful rate at lower temperatures than would otherwise be needed, saving energy |
| Enzymes in industry | Biological catalysts used, for example, in processing paper pulp and in detergents | Lower temperatures and smaller amounts of harsh chemicals such as bleaches are needed |