[maccha]

My chemistry textbook explains that catalysts provide an alternate route to the reaction with a lower activation energy barrier, but my biochemistry text says that catalysts lower the activation energy by stabilizing the transition state. I'm confused by these two definitions- do they contradict each other?

[Wald_ron]

No they don't.

[SirRoderick]

Nope.

[DevaDevil]

My chemistry textbook explains that catalysts provide an alternate route to the reaction with a lower activation energy barrier, but my biochemistry text says that catalysts lower the activation energy by stabilizing the transition state. I'm confused by these two definitions- do they contradict each other?

What your teacher means by alternative route may also be the exact same intermediates, but coordinating to the catalyst.

For example, ammonia oxidation: in order to get nitrogen, you first need to break a N-H bond in the ammonia, this can be done by high temperature (very difficult as ignition temperature is usually higher than a flame's temperature), or this can be done quite efficiently by a Pt- catalyst at a lot lower temperature. In other words, the catalyst lowers the activation barrier.

In both cases the intermediates will be compounds such as NH₂, NH etc, but in case of the Pt-catalyst, these compounds are more stable as they can coordinate to the Pt-surface and share in the electron density of the metal, offsetting the loss of paired electrons. This lowers the energy required to break a N-H bond.

So the catalyst lowers the activation energy, because the intermediates are more stabilized.

(think of the energy you require as a mountain; the higher the energy you need, the higher the mountain, with the most unstable intermediates at the peak. The more unstable they are, the harder they are to make out of a stable compound, so you need a long uphill climb to get there. If those intermediates are stabilized, they require less energy to make, and the mountain will be smaller, requiring less effort to get over it.)

[ajkoer]

My observations is that not all catalysts act in the same manner. For example, I recalling seeing a copper compound (the so-called catalyst) that is involved in a series of chemical reactions, which at the end, has the compound regenerated. This would be an example of introducing different reaction sequences/paths to lower the activation threshold to achieve a reaction. Note, although the compound is chemically involved in the reaction, as it is NOT consumed, it meets the definition of a catalyst.

There is also the nano-particle example. Here, the surface area seems to lower the activation energy required.

Others may provide more examples.

[g-bones]

so typically, when discussing catalyst in terms of strictly organic or general chemistry, they change the mechanism by which a reaction occurs, typically by producing different intermediates along the reaction pathway.  These intermediates are lower in energy than would be for a different reaction pathway without the catalyst that delivers the same product.  In biochemistry, things like enzymes do not change the mechanism for the most part, the hold the two reactive parts of any substrates in close proximity and/or activate them through hydrogen bonding etc (thereby stabilizing any developing charges in the transition state).  So for the most part the molecules are closer in energy to the transition state by being held so closely to each other.  this is not to say all enzymes proceed this way but, I think this is the point your text was getting at.  Hope this helps.