Video Transcript
Considering the diagram, which of
the following statements is not true of a catalyzed reaction? (A) Fewer particles have sufficient
energy to react. (B) The catalyst provides an
alternative reaction pathway. (C) The catalyst is not consumed
during the reaction. (D) The activation energy is
smaller. Or (E) the overall energy change is
the same.
Here we are looking to discern
which statement is not a correct statement for catalysis reactions. Scientists use collision theory to
understand how reactants turn to products and at what rate. Collision theory states that
reactant particles will turn into products when they collide with enough energy to
overcome the energy barrier called the activation energy. The activation energy is defined as
the minimum amount of energy required for a reaction to occur.
When the activation energy of a
given reaction is high, such as the pink line in the diagram, then conversion to
product is low. Conversely, when the activation
energy is low, such as in the dotted red line, conversion to product is high because
it requires less energy to overcome the barrier.
When activation energies become
very high, it’s common to use catalysts. A catalyst is a substance that
increases the rate of a reaction without undergoing a permanent chemical change. The result is that a catalyst
lowers the activation energy, allowing the reaction to proceed more easily via a new
reaction pathway. It’s important to note that in both
the catalyzed and noncatalyzed reaction pathways, the overall energy change, or Δ𝐻
of the reaction, is the same.
And with this information, we can
revisit and answer the question. Considering the diagram, which of
the following statements is not true of a catalyzed reaction? And the correct answer is (A):
fewer particles have sufficient energy to react.
