Collision Theory

In the previous section, we learned how a reaction coordinate diagram can be used to extract information about a reaction's enthalpy, i.e. how much heat it releases or absorbs.

We will now learn how to observe different features of a reaction coordinate diagram to extract information about a reaction's rate, i.e. how fast reactants are converted to products. We will then learn the collision theory of reaction rates, which will help us understand the features and conditions that cause some reactions to be fast and others slow.

As before, many things we will discuss here will be partially familiar to you. For example, you are probably aware that some reactions happen in an instant (like the color changes we saw in pH indicators in Lesson 5) while others take months or years (the slow rusting of nails is just one example). You also know that reactions can be sped up by their conditions like heat. In this section, we will learn the reasons for these and other phenomena.

Particle Collisions

The framework we are learning is called "collision theory" because it makes a simple but important assumption: in order for a reaction to occur between two (or more) substances, particles of those substances must collide with each other. This makes intuitive sense: if things are going to form a bond, they must get close enough to interact.

Required For Reaction

Particles must collide
...

Depending on the reaction, these particles may be atoms, molecules, or ions. This assumption has some shortcomings (for example, it ignores reactions where there is only one reactant, like decompositions) but it will suffice to let us explore the dynamics of reaction rates.

This requirement that particles must collide for a reaction to occur is the main reason that it is rare for reactions to take place between two solids. In a solid, the particles are locked up in lattices and do not move around, so it is all but impossible for collisions between particles to occur.

Let's return to our reaction coordinate diagrams from earlier in the lesson. As we learned, every reaction begins with the potential energy of the atoms increasing to a transition state. The size of this "climb" - the difference in energy between the reactants and the transition state - is called the activation energy (E_a).

In order to "climb" this energetic hill, the reactant molecules need to come into their collision with sufficient kinetic energy. Essentially, this means that they need to be moving fast enough that some bonds can be broken as they collide, allowing the reactants to rearrange themselves into products.

Thus, we need more than just for our particles to collide with each other. They must do so with enough kinetic energy (i.e. enough speed) that a reaction can take place.

Required For Reaction

Particles must collide
Particles must have enough energy
...

Finally, let's think about the shape and arrangement of our particles. In the reaction at right, between CO and O₂, there are two ways we can imagine these molecules colliding. If the O₂ molecule strikes the oxygen atom in CO, no CO₂ molecule can be formed, because CO₂ does not have an O-O bond. We only form CO₂ if the molecules collide with the O₂ molecule striking the carbon atom (and if the colliding molecules have enough kinetic energy, as mentioned above).

So we can add a third requirement to our list. In order for a reaction to occur under collision theory, the particles must collide, they must do so with sufficient kinetic energy, and they must do so in the correct orientation.

Required For Reaction

Particles must collide
Particles must have enough energy
Particles must have correct orientation

Factors Affecting Reaction Rates

Using this framework, we can now learn how different features of a reaction, and different reaction conditions, can affect how fast reactions proceed.

Reaction Features: Orientation Factor

Some reactions are intrinsically slower than others because of the orientation requirements imposed by the molecules involved. In some reactions, almost any orientation will work. For example, when O₂ reacts with H₂, both reactant molecules are symmetrical, so any collision can produce a reaction. In the example of CO and O₂ above, the odds of a correct orientation are roughly 50:50. And in the case of some geometrically complex molecules, the odds of a correct orientation can be so small that the reaction becomes very slow or even functionally non-existent.

Reaction Features: Activation Energy

The other factor that can make one reaction intrinsically faster or slower than another is its activation energy. As mentioned above, for a reaction to occur, particles need to collide with enough kinetic energy to get over the "energetic hill" that is the transition state. If that hill is shorter, a given particle collision is more likely to have sufficient energy, so the reaction is likely to go faster.

Reaction Conditions: Concentration

Going back to the first idea we touched on, we know that, for a reaction to take place, particles must collide. If collisions happen more frequently, then the reaction should happen faster. One of the simplest ways for this to occur is for the concentration of one or both reactants to be increased. If the reactant concentration is greater, then the reaction rate will tend to be faster.

Reaction Conditions: Temperature

Recall from our discussion of Kinetic Molecular Theory in Lesson 1 that temperature is a measure of the kinetic energy of particles. At higher temperatures, particles are moving faster, so they have greater kinetic energy. And, as we have just discussed, if a reaction is to occur, particles must collide with enough energy to overcome the activation energy of the reaction.

For this reason, raising the temperature of a reaction will almost always increase its rate. In the diagram at right, when the molecules collide with low energy, the will likely just bounce off each other. But in a high-energy collision, a reaction may be initiated, with bonds breaking and forming.

Reaction Conditions: Surface Area

When a solid reactant is broken into smaller and smaller pieces, the rate at which it reacts increases. Since only the molecules on the surface of the solid can collide with other molecules, the greater the surface area, the more collisions can take place per second, and the faster the reaction.

Occasionally this is dramatically (and tragically) illustrated in grain elevators, where very small particles of grain dust or flour offer such a large surface area that combustion, if allowed to occur, can do so at an explosive rate. Grain elevator explosions were unfortunately common before strict safety regulations were imposed.

This phenomenon is not limited to solids either. When gasoline is burned in car engines, a carburetor or fuel injector is used to render the liquid gasoline into a very fine aerosol made of extremely small liquid particles. This allows for much more rapid combustion than if the same amount of gasoline were dumped into the cylinder in a single glob.

Reaction Conditions: Catalysts

A final factor that can affect the rate of a reaction is a catalyst. A catalyst is a substance that increases the rate of a reaction without itself being consumed in the process.

One example of a catalyst is the catalytic converter used in automobile exhaust systems. Normally, some exhaust gases (like NO₂) break down very slowly, but when they come in contact with the catalyst in the converter, they react very rapidly to form much less toxic and harmful gases (like N₂ and O₂).

Catalytic converters contain metals like platinum and rhodium; these metals are not used up in the reaction, they just help it to go faster. Once the product molecules detach from the metal, it is free to assist the next reaction, and the next, and so on. This is a wonderfully efficient process with two pitfalls. First, while the metal is not consumed in the reaction, its surface can get dirty and worn out over time, eventually requiring replacement. Second, the metals required are extremely expensive, making catalytic converters attractive targets for theft.

Catalysts work by providing an alternate mechanism, or path, for the reaction that has a lower activation energy than the original path. Returning to our reaction coordinate diagrams, a catalyst lowers the energy of the transition state, as shown in the image at right.

Essentially, catalysts provide an easier way to get from reactants to products and the lower activation energy results in a faster reaction.

This week's lab experiment illustrates the effects of these various factors on the rate of the reaction between magnesium metal and hydrochloric acid. This is the first part of a two-part lab for this lesson.

Since there is no known catalyst for this reaction, we can’t illustrate the effect of catalysts. However, we do explore the effects of concentration, temperature, and the surface area of the solid on the rate of the reaction.

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