Single-Step Reaction Kinetics Using Collision Theory

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This Demonstration considers the kinetics of a single-step reaction; an example might be the formation of hydrogen iodide [1]. Two molecules in the gas state are shown colliding with one another; hydrogen molecules (black) move faster than iodine molecules (red) since they are less massive. The rotational states are not taken into account. Two conditions must be satisfied in order for the chemical reaction to take place: correct geometry () and correct force (). The transition state is the intermediate configuration in which old bonds break and new bonds form. Three collision possibilities are considered: an ideal collision with perfect alignment (a very rare occurrence), acceptable conditions that exist within certain limits, and two unaligned molecules that cannot cause a reaction to occur. These three cases are considered for both forward and reverse reactions for a total of six possible reactions [2]. Use the "kinetic energy" slider to increase the energy until sufficient force can be reached. Then the "time" slider will show the collision in progress.


The "theory" window shows the chemical reaction for four different types of schemes. The top-right plot shows the condition for the reaction to take place as the union of sets using a Venn diagram. This shows the possible combinations of geometry and energy. The top-left plot considers just the energy, independent of geometric factors. If the energy exceeds the activation barrier (indicated as Ea, activation energy), the reaction can always occur. If the energy of the products is lower than the energy of the reactants, there will be a heat release (, positive change in enthalpy); otherwise, there will be a heat absorption (, negative change in enthalpy). Use the "time" slider to move the reaction coordinates and check if the molecular energy is high enough to surmount the energy barrier. The bottom-left plot shows the state of the covalent bonds of both products and reactants; the energy barrier is determined by the transition state, when the old bonds are half broken and the new ones are half made. The bottom-right plot formalizes the collision type. Lines identify the molecule: a dashed red line for iodine, a black line for hydrogen and a dashed blue line for hydrogen iodide.

The last case considers the effects of a catalyst; lowering the transition state energy does not modify the thermodynamic equilibrium but it does enable the reaction to run faster. The top plots show the catalyst effect only on the activation energy (sometimes it could make easier having a correct geometry); without catalyst the activation barrier (Ea) is higher than the case with it ( ). The plots show that the enthalpy does not change with a catalyst, and the Venn diagram is appropriately modified; with the catalyst a lower temperature is required to exceed the activation energy so the correct force circle is shifted to the left.


Contributed by: D. Meliga and S. Z. Lavagnino (June 2018)
Additional contribution by: G. Follo
Open content licensed under CC BY-NC-SA



Snapshot 1: two molecules with the correct orientation that bounce elastically because the kinetic energy is below the activation energy

Snapshot 2: two molecules with the correct orientation and a kinetic energy level that is above the activation energy start the chemical reaction

Snapshot 3: molecular collision theory applied to a case with the correct energy but the wrong orientation

Snapshot 4: molecular collision theory applied to a reaction with catalyst


[1] G. C. Pimentel and R. D. Spratley, Understanding Chemistry, San Francisco: Holden-Day Inc., 1971.

[2] D. McQuarrie and J. Simon, Physical Chemistry: A Molecular Approach, Sausalito, CA: University Science Books, 1997.

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