Activation energy, transition state, and reaction rate.
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 .
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 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 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 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.)
Reaction coordinate diagram for an exergonic reaction. Although the products are at a lower energy level than the reactants (free energy is released in going from reactants to products), there is still a "hump" in the energetic path of the reaction, reflecting the formation of the high-energy transition state. The activation energy for the forward reaction is the amount of free energy that must be added to go from the energy level of the reactants to the energy level of 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.
Activation energy and reaction rate
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.
Want to join the conversation?
- what is the defination of activation energy?(26 votes)
- The official definition of activation energy is a bit complicated and involves some calculus. But to simplify it:
Activation energy is the minimum energy required to cause a process (such as a chemical reaction) to occur.(106 votes)
- I thought an energy-releasing reaction was called an exothermic reaction and a reaction that takes in energy is endothermic. In the article, it defines them as exergonic and endergonic. Are they the same?(2 votes)
- Exothermic and endothermic refer to specifically heat. Exergonic and endergonic refer to energy in general.(21 votes)
- can a product go back to a reactant after going through activation energy hump? (sorry if my question makes no sense; I don't know a lot of chemistry)(1 vote)
- Theoretically yes, but practically no.
So this concept can be visualized with combustion or fire. While wood does not spontaneously burst into flame, if you add additional energy, with a match for an example, to the pile of wood, it starts a fire. What happens is that the energy in the match pushes the wood over the activation energy hump and starts the fire. Afterwards, the fire is self-sustaining because the fire creates enough heat to activate the rest of the wood.
Chemically, wood is composed of mostly carbon, which reacts with the oxygen in the air when 'activated' to create carbon dioxide.
So, for this reaction, carbon is the reactant and carbon dioxide is the product, which can be converted back into carbon (like photosynthesis) but requires more energy to do so.
The bottom line is that while it is possible, it will (in general) require additional energy to go back from a product to a reactant(8 votes)
- When mentioning activation energy: energy must be an input in order to start the reaction, but is more energy released during the bonding of the atoms compared to the required activation energy? Can the energy be harnessed in an industrial setting?(3 votes)
- In an exothermic reaction, the energy is released in the form of heat, and in an industrial setting, this may save on heating bills, though the effect for most reactions does not provide the right amount energy to heat the mixture to exactly the right temperature. Often the mixture will need to be either cooled or heated continuously to maintain the optimum temperature for that particular reaction. For endothermic reactions heat is absorbed from the environment and so the mixture will need heating to be maintained at the right temperature. By right temperature, I mean that which optimises both equilibrium position and resultant yield, which can sometimes be a compromise, in the case of endothermic reactions.(3 votes)
- I thought an energy-releasing reaction was called an exothermic reaction and a reaction that takes in energy is endothermic. In the article, it defines them as exergonic and endergonic. Are they the same?(4 votes)
- I read that the higher activation energy, the slower the reaction will be. This makes sense because, probability-wise, there would be less molecules with the energy to reach the transition state. Is there a limit to how high the activation energy can be before the reaction is not only slow but an input of energy needs to be inputted to reach the the products? In other words with like the combustion of paper, could this reaction theoretically happen without an input (just a long, long, long, time) because there's just a 1/1000000000000..... chance (according to the Boltzmann distribution) that molecules have the required energy to reach the products. Looking at the Boltzmann dsitribution, it looks like the probability distribution is asymptotic to 0 and never actually crosses the x-axis.(3 votes)
- Even if a reactant reaches a transition state, is it possible that the reactant isn't converted to a product? So even if the orientation is correct, and the activation energy is met, the reaction does not proceed?(3 votes)
- When a rise in temperature is not enough to start a chemical reaction, what role do enzymes play in the chemical reaction?(2 votes)
- I think you may have misunderstood the graph — the y-axis is not temperature it is the amount of "free energy" (energy that theoretically could be used) associated with the reactants, intermediates, and products of the reaction.
Temperature is related to the average amount of kinetic energy for a group of molecules. Some of those molecules will have much more than the average amount, some will have much less, and many will have an amount of energy close to the average — look up "Maxwell-Boltzmann distribution" for more information§.
This means that some reactant molecules will have enough energy to reconfigure themselves into the transition state. They can then release energy by converting into the products (or by going back to the reactants).
Enzymes provide a lower energy pathway for the reactants to become products — since less energy is required, more molecules at a give temperature will have enough to proceed through the reaction.
Does that help?
§Note: Examples of KhanAcademy information on this subject:
- What is the
diffrenece between transition state and activation complex?(2 votes)
- They are different because the activation complex refers to ALL of the possible molecules in a chain reaction, but the transition state is the highest point of potential energy(0 votes)
- It's saying that if there is more activation energy (the small amount of energy needed to break the molecules apart), the chemical reaction would be slower... I don't understand why. I would think that if there is more energy, the molecules could break up faster and the reaction would be quicker?(2 votes)
- No, if there is more activation energy needed only means more energy would be wasted on that reaction.
Well, if we speak of one enzyme. And we have reaction a and reaction B:
Imagine reaction A has activation energy 15kJ and reaction B has 20kJ.
Which one would be slower? Enzyme can accelerate it for 15kj in let's say 10min. If it has to accelerate for 20kj and it can only for 15kj, it means it would require doing it longer than 10min - 15 + 5kj. in two sessions.(1 vote)