Activation Energy


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Sarah Sharma 2J
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Joined: Fri Sep 29, 2017 7:05 am

Activation Energy

Postby Sarah Sharma 2J » Sun Mar 11, 2018 3:44 pm

When activation energy is at its optimum, what does that mean for collisions? Are they taking place at the fastest at this point?

Zane Mills 1E
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Re: Activation Energy

Postby Zane Mills 1E » Sun Mar 11, 2018 8:20 pm

Not really sure what you mean by optimum but activation energy is essentially an energy barrier between reactants and products. It's not so much collisions as it is stability, as catalysts help to stabilize intermediates and bring down the 'barrier' between products and intermediates. I guess optimum would be a good word for the enzyme b/c the better the enzyme the lower the Ea and thus the faster the rxn will proceed.

Nathan Tu 2C
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Joined: Fri Sep 29, 2017 7:07 am

Re: Activation Energy

Postby Nathan Tu 2C » Sun Mar 11, 2018 8:53 pm

If activation energy is reached in the reactants, then when the reactants collide to form a product or intermediate, they will form the product/intermediate. Without reaching the activation energy required, reactant collisions will temporarily create the product but then reverse back into reactants because of insufficient energy.

Varsha Sivaganesh 1A
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Joined: Thu Jul 13, 2017 3:00 am

Re: Activation Energy

Postby Varsha Sivaganesh 1A » Mon Mar 12, 2018 12:14 am

I wouldn't necessarily describe this situation as having an "optimum" activation energy. Think of it as, if there is enough energy supplied, then the reactants can reach the activation energy which means they have enough energy to break bonds so that the reaction can occur. If activation energy is lowered (i.e. through use of a catalyst), then the reaction can occur faster.

Guangyu Li 2J
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Re: Activation Energy

Postby Guangyu Li 2J » Mon Mar 12, 2018 2:01 am

If the kinetic energy of the molecules reaches the activation energy, they will be able to collide with each other effectively and allow the reactions to proceed.

The reactions can be speeded up by lowering the energy barrier of the reactions. One of the most significant examples of this point is the catalyst. If the activation energy required for a reaction decreases, it will proceed faster than before.


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