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 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
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.)
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.
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- what is the defination of activation energy?(30 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.(116 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?(23 votes)
- Exothermic and endothermic refer to specifically heat. Exergonic and endergonic refer to energy in general.(26 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)(3 votes)
- 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(18 votes)
- Is there a difference between the terms endothermic/exothermic reaction and endergonic/exergonic reaction?
I only learned endothermic and exothermic as a reaction that stores energy and a reaction that releases energy.(4 votes)
- Endothermic and exothermic refers to sign of the enthalpy of a reaction. Whether the net enthalpy change is positive or negative respectively. Where enthalpy is equivalent to heat.
Endergonic and exergonic refers to the sign of the Gibbs free energy of a reaction. Whether the net free energy change is positive or negative respectively. Free energy taking into consideration both the enthalpy and entropy change of a reaction.
Hope that helps.(10 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.(8 votes)
- I don't get this. If a molecule has more activation energy, shouldn't it be more likely to reach the high barrier required and complete the chemical reaction faster? If I have more energy when I wake up, it is easier to get out of bed and it takes me less time to do so. Shouldn't chemical reactions be the same?(5 votes)
- yeah, like amathakbari said-activation energy is the amount of energy needed to activate the complex that ocurrs at the transition state. it isn't energy you have(4 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?(4 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.(4 votes)
- What is the activation energy of the reaction?(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?(5 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)