Collision Theory — Definition & Overview

(Video) Collision Theory & Reactions Part 1 | Reactions | Chemistry | FuseSchool

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Collision theory explains part of how chemical reactions occur. Chemicals must collide with each other to react, but very collision does not create a reaction.

An example is used with two diatomic molecules represented by $A_2$ and $B_2$. When these react they form two molecules of $AB$. The orientation of the molecules matters to cause a reaction.

When two atoms collide they can hit each other and then move away or combine. When they combine this is called a fruitful collision. For a fruitful collision, the particles must collide with enough energy that the reaction can occur. This amount of energy needed is called the activation energy.

Part of the reason I love chemistry is that chemical reactions are so unique. They create everything we use in modern life. Some examples are fertilizers, batteries, and clothing fibers. And sometimes, chemistry is just fun. It's cool to mix two chemicals and see them change colors. What if we want to see a violent reaction? You can always put an alkali metal in water!

But, what's happening at the molecular scale? What are the atoms and molecules doing? What causes chemical reactions to occur? To answer these questions, we need collision theory.

Components of Chemical Reactions

There are three components to chemical reactions.

• Collisions. For a reaction to happen, the reactants must collide.

• Anything that increases the number of collisions will increase the reaction rate.

• For example, consider the states of matter. We know the movement of particles is fastest in gases. But slower in liquids and solids. So, we would expect the most collisions between gas molecules. Gas reactions tend to be faster.

• Increasing the temperature will raise the average kinetic energy. The result is more collisions.

• Orientation. Many colliding molecules do not react. Why? The molecules need to have the proper orientation.

• Consider the reaction $2CH_{3}CH_{2}Br + 2 OH^{-} \rightarrow 2CH_{3}CH_{2}OH + Br_{2}$. This reaction is from organic chemistry. We are forming ethanol. (And yes, that's the alcohol in drinks.)

• Let's consider where the hydroxide $(OH^{-})$ needs to hit.

• When the hydroxide hits the methyl ($CH_{3}$) group, nothing happens. Methyl groups are very stable.

• What happens if the hydroxide hits the bromine? The carbon-bromine bond is a polar covalent bond. Bromine is electronegative. So, it has a partial negative dipole. Hydroxide has a full negative charge. So, there will be an electrostatic repulsion between them. Nothing happens!

• So, the hydroxide has to collide with the methylene ($CH_{2}$) carbon for the reaction to happen.

• Activation Energy. Sometimes, molecules will collide with the correct orientation and still not react. What's happening? They don't have enough energy.

• Molecules store potential energy in their bonds. For reactions to arise, the bonds need to break. So, the collision needs enough kinetic energy to break the bond. Anything that raises the kinetic energy increases the reaction's rate.

• On an energy diagram, activation energy is the difference between the reactants' energy and the transition state's energy.

• Catalysis lower the activation energy by providing an alternative reaction mechanism.

Connecting Collision Theory and Reaction Rate

We can connect reaction rates and collision theory with the Arrhenius Equation.

• Remember, the equation is $k= Ae^{\frac{-E_{a}}{RT}}$.

• k is the rate constant.

• A is the frequency factor.

• $A=z\times\rho$

• z is the number of collisions per second.

• $\rho$ is the fraction of collisions with proper orientation. $\rho$ is always greater than zero and less than 1.

• $E_{a}$ is the activation energy.

• $R$ is the universal gas constant. For the Arrhenius equation, we use $8.314 \frac{J}{mol\times K}$.

• $T$ is the kelvin temperature.

• Let's simplify the Arrhenius Equation by taking the natural log of both sides. $ln{(k)}=ln{(z\rho)}+ \frac{-E_{a}}{R}\times\frac{1}{T}$.

• The equation is in the y=mx+b form.

• If we plot $ln{(k)}$ versus $\frac{1}{T}$ we get a linear graph. The graph's slope is $\frac{-E_{a}}{R}$. The intercept is $ln{(z\rho)}$.

• The activation energy term is negative. So, for larger $E_{a}$ values, $ln{(k)}$ is smaller. The larger the activation energy, the slower the reaction rate.

• Higher temperatures make the activation energy term smaller. So, $ln{(k)}$ is larger. The reaction is faster.