Activation energy

Imagine waking up on a day when you have lots of fun stuff planned. Does it ever happen that, despite the exciting day that lies ahead, you need to muster some extra energy to get yourself out of bed? Once you’re up, you can coast through the rest of the day, but there’s a little hump you have to get over to reach that point.

The activation energy of a chemical reaction is kind of like that “hump” you have to get over to get yourself out of bed. Even energy-releasing (exergonic) reactions require some amount of energy input to get going, before they can proceed with their energy-releasing steps. This initial energy input, which is later paid back as the reaction proceeds, is called the activation energy and is abbreviated $\text E_{\text A}$start text, E, end text, start subscript, start text, A, end text, end subscript.

Why would an energy-releasing reaction with a negative ∆G need energy to proceed? To understand this, we need to look at what actually happens to reactant molecules during a chemical reaction. In order for the reaction to take place, some or all of the chemical bonds in the reactants must be broken so that new bonds, those of the products, can form. To get the bonds into a state that allows them to break, the molecule must be contorted (deformed, or bent) into an unstable state called the transition state. The transition state is a high-energy state, and some amount of energy – the activation energy – must be added in order for the molecule to reach it. Because the transition state is unstable, reactant molecules don’t stay there long, but quickly proceed to the next step of the chemical reaction.

In general, the transition state of a reaction is always at a higher energy level than the reactants or products, such that $\text E_{\text A}$start text, E, end text, start subscript, start text, A, end text, end subscript always has a positive value – independent of whether the reaction is endergonic or exergonic overall. The activation energy shown in the diagram below is for the forward reaction (reactants $\rightarrow$right arrow products), which is exergonic. If the reaction were to proceed in the reverse direction (endergonic), the transition state would remain the same, but the activation energy would be larger. This is because the product molecules are lower-energy and would thus need more energy added to reach the transition state at the top of the reaction “hill.” (An activation energy arrow for the reverse reaction would extend from the products up to the transition state.)

The source of activation energy is typically heat, with reactant molecules absorbing thermal energy from their surroundings. This thermal energy speeds up the motion of the reactant molecules, increasing the frequency and force of their collisions, and also jostles the atoms and bonds within the individual molecules, making it more likely that bonds will break. Once a reactant molecule absorbs enough energy to reach the transition state, it can proceed through the remainder of the reaction.

The activation energy of a chemical reaction is closely related to its rate. Specifically, the higher the activation energy, the slower the chemical reaction will be. This is because molecules can only complete the reaction once they have reached the top of the activation energy barrier. The higher the barrier is, the fewer molecules that will have enough energy to make it over at any given moment.

Many reactions have such high activation energies that they basically don't proceed at all without an input of energy. For instance, the combustion of a fuel like propane releases energy, but the rate of reaction is effectively zero at room temperature. (To be clear, this is a good thing – it wouldn't be so great if propane canisters spontaneously combusted on the shelf!) Once a spark has provided enough energy to get some molecules over the activation energy barrier, those molecules complete the reaction, releasing energy. The released energy helps other fuel molecules get over the energy barrier as well, leading to a chain reaction.

Most chemical reactions that take place in cells are like the hydrocarbon combustion example: the activation energy is too high for the reactions to proceed significantly at ambient temperature. At first, this seems like a problem; after all, you can’t set off a spark inside of a cell without causing damage. Fortunately, it’s possible to lower the activation energy of a reaction, and to thereby increase reaction rate. The process of speeding up a reaction by reducing its activation energy is known as catalysis, and the factor that's added to lower the activation energy is called a catalyst. Biological catalysts are known as enzymes, and we’ll examine them in detail in the next section.