Can activation energy be zero?

[mjassim]

I've been trying to find a definitive answer to this question since long, and I still haven't found one!

My teacher at school says that it can't be zero, according to the Arrhenius equation. She says if the activation energy were to be zero, every collision would result in a reaction, which does not take place, and hence activation energy is not zero.

But doesn't a reaction take place when there is an effective collision, ie, with sufficient energy(>E_a) and with proper orientation (according to the collision theory)? Or is there something I have misunderstood about activation energy all along?


It'd also be great if someone could tell me what the meaning of activation energy actually is. I've searched online, and I've come up with various answers, and now I'm confused.

 

[Mapes]

Hi mjassim, welcome to PF.

Between the beginning and the end of a process (e.g., a reaction), sometimes there is an intermediate state with a higher energy than either of the other two states. For the process to proceed, the system must reach this energy, even if only instantaneously. The activation energy is the difference between the maximum intermediate energy and the starting energy.

Not every process has an activation energy. For a ball at the top of a hill, for example, there's no barrier to overcome before rolling initiates. There's not even a requirement that a chemical reaction has to have an activation energy, but I believe it would imply that both the starting and ending species are barely chemically bound. This state could be difficult to enforce experimentally.

 

[physics girl phd]

There is a nice simulation http://phet.colorado.edu/simulations/sims.php?sim=Reversible_Reactions" that might help you understand reversible reactions better (you set the energy of the products and reactants, as well as temperature, etc.). You'll need to have Java installed on your computer.

 

[mjassim]

@Mapes:

"The activation energy is the difference between the maximum intermediate energy and the starting energy."

So that basically means if I were to measure (somehow) the energy of a reaction system at, say, 293K, and the energy at this temperature is that "maximum intermediate energy", then the activation energy is zero?

Or is it that, activation energy is defined for a reaction at a certain temperature and with certain conditions (such as STP) and all?


@physics girl phd:

Thanks for the link! I've tried it out, and it seems to hold well with what I think. But I'm not sure if considering a reversible reaction would apply to irreversible reactions too. In any case, I'll show this to my teacher, and see what she has to say.

Thanks a lot, both of you!

 

[chemisttree]

I think your teacher has this one right! Think about it this way... the activation energy prevents every single molecule from undergoing the reaction in the same instant (and there is a word for reactions like that). Kind of like the drain of a tub filled with water restricts the flow of the water from the tub. To complete the analogy, the same water in the tub would have zero restriction... if there were no tub in the first place! A tub-shaped mass of water would have no restriction to flow but is difficult to achieve in the laboratory.

 

[mjassim]

Hmm... okay, I understand that part well. So does that mean that activation energy can theoretically be zero, but not in practice?


I've checked the Wikipedia page for activation energy and it says that it is "the energy that must be overcome in order for a chemical reaction to occur." In the same paragraph, it states it "may otherwise be denoted as the minimum energy necessary for a specific chemical reaction to occur."

Now, I'm confused as to whether activation energy is the excess energy supplied to the initial energy of the reactants, or it is the energy already present in the system. Can you please clarify?

Thanks a lot for helping out!

 

[Mapes]

Typically the necessary energy is already in the system. In a collection of molecules at a given temperature, there will be a distribution of energies. Some molecules will have enough energy to get over the "hump," some won't (but may later). The higher the temperature, the sooner a given molecule will randomly attain enough energy for the reaction to proceed. The reaction rate for the entire system is usually modeled as being exponentially dependent on temperature, $r\propto \exp(-E/kT)$ (this is the Arrhenius model).