Effective Collisions
Definition and meaning of Effective Collisions in chemistry.
Effective collisions are collisions between reactant particles that actually produce a chemical reaction, occurring only when the particles collide with kinetic energy equal to or greater than the activation energy and with the correct spatial orientation for the necessary bonds to break and form.
In more detail
According to collision theory, particles must physically collide to react, but the vast majority of collisions are "unproductive" because they fail one or both requirements. The Maxwell-Boltzmann distribution shows that only a small fraction of molecules possess energy at or above the activation energy at any given temperature, so raising temperature sharply increases the fraction of collisions that are effective. Even energetic collisions can fail if the molecules meet with the wrong geometry, which is captured by the steric (orientation) factor, p, in the modified Arrhenius equation.
Key facts
| Field | Physical Chemistry |
|---|---|
| Energy requirement | Kinetic energy ≥ activation energy (Ea) |
| Geometric requirement | Correct orientation, described by the steric factor, p |
| Typical fraction effective | Usually a very small percentage of total collisions |
In the gas-phase reaction NO(g) + O3(g) → NO2(g) + O2(g), a collision is effective only if the nitrogen atom of NO strikes a terminal oxygen atom of O3 with enough kinetic energy to overcome the activation energy; collisions with the wrong orientation, or with too little energy, simply bounce apart unreacted.
Frequently asked questions
Why do most molecular collisions not lead to a reaction?
Most collisions fail because the colliding particles either lack the minimum kinetic energy (activation energy) needed to break existing bonds, or they collide with an orientation that does not allow the reacting atoms to align properly.
How does increasing temperature increase the number of effective collisions?
Raising the temperature increases the average kinetic energy of the particles, and by the Maxwell-Boltzmann distribution a larger fraction of molecules then have energy at or above the activation energy, which increases both collision frequency and the proportion of collisions that are effective.