The collision model of kinetics states that, in order for two molecules to react, the molecules must first collide with each other with a certain amount of energy called the activation energy. If the molecules do not have enough energy, they will simply bounce off each other.
The activation energy component of the collision model is integral to the understanding of chemical behavior. The air that we breath is composed heavily of `N_2` and `O_2` gases which are constantly colliding with each other but, fortunately, don't react to form `NO_2, NO_3, "or" N_2O_4` at room temperature. The reason for this is that the energy from collision is below the activation energy.
As temperature increase, the average energy of molecules increases and thus the activation energy is reached by more molecules. This examples why heating reactions up often increases the reaction rate. By heating up the reaction, the molecules have more energy and more of the molecules will have sufficient energy to react.
If we plot the reaction progress vs. energy, we can see a chart like this where the little "hill" represents the activation energy. Every reaction has an activation energy, though some are negligible at room temperature whereas others are practically insurmountable. The takeaway from this section is that for every reaction, there is an energy cost in order for the reaction to occur.
The Arrhenius Equation gives the relationship between reaction rates and temperature:
`"ln"(K_2/(K_1))=-E_a/R(1/(T_2)-1/(T_1))`
`K_n` is the rate constant of temperature `n` .
`T_n` is the temperature corresponding to `n` .
`E_a` is the activation energy of the reaction.
`R` is the gas constant.
With this equation, we can calculate the activation energy of a reaction given the reaction rate at two different temperatures.
Catalysts serve the function of reducing the activation energy of a reaction. By reducing the activation energy, the catalyst is effectively speeding up the reaction.
An important property of catalysts is that catalysts are not consumed during the reaction. If you use a catalyst at the beginning of a reaction, it must be present after the reaction is complete in order for it to be considered a catalyst. This does not mean that the catalyst stays the same throughout the reaction; most of the time, catalysts are broken at some point during the reaction and re-formed by the end.
One way of visualizing catalysts is to imagine yourself driving down a highway during rush hour; chances are, you're going to be driving pretty slowly. Now imagine that a new lane opens up just for you to drive on. Suddenly, you're able to drive much faster and get to your destination more quickly. This is the essence of catalysts; the start and end are the same, but catalysts open up an alternative, easier pathway.
We can sum up this statement by saying that catalysts speed up reactions by lowering the activation energy, without being consumed during the reaction.
1. The collision model states that in order for molecules to react, they must first collide into each other with sufficient energy.
2. The Arrhenius equation gives a relationship between rate and temperature.
3. Catalysts reduce the rate of a reaction by lowering the activation energy. An important property of catalysts is that they are not consumed during the reaction.