Factors Affecting Rate of Reaction

Click to enlarge

How different factors can affect how quickly a reaction happens is a common topic in the chemistry curriculum. This graphic serves as a convenient summary for teachers and students alike of what these different factors are, and how and why they affect the rate of a reaction. However, it’s not only of interest to those teaching or learning about chemistry; as we’ll see, knowledge of these factors can have some everyday applications too!

Before considering what can affect how quickly a reaction happens, we need to think about how a reaction happens in the first place. ‘Collision theory’ offers a simplified way of looking at this, based on ideas of particles. For a reaction to happen, particles of the reactants need to collide, and with enough energy to break bonds and trigger the reaction. They also have to collide in the correct orientation for the reaction to occur. If their orientation relative to one another is incorrect, or they don’t collide with enough energy, then no reaction occurs and the particles simply bounce off each other.

Within a given time period, a certain number of collisions between reactant particles will occur. A certain proportion of these collisions will have enough energy to exceed the reaction’s ‘activation energy’, the energy required for the reaction to take place, and a reaction will occur. By changing the conditions of the reaction, we can affect both the frequency of collisions between particles, and the proportion of collisions that are successful and result in a reaction.

The first thing we can change, for solutions (substances dissolved in a solvent, commonly water), is concentration. Concentration is a measure of how much of something is dissolved in a solution; a higher concentration means there is more of the particular substance in the solution. Having more of the substance means there are a greater number of particles, which increases the chance of a collision between reactant particles occurring. This increases the frequency of collisions, leading to an increased rate of reaction. Note that the proportion of successful collisions remains the same, as changing the concentration doesn’t affect the energy of the particles.

We can also increase the temperature of the reaction to speed it up. This works because increasing the temperature increases the kinetic energy of the reactant particles. Because they are moving around faster, the collisions between them are more frequent; additionally, as on average the particles have more energy at a higher temperature, a greater number of collisions will have the energy required to react, so the proportion of successful collisions increases. Both of these factors combined lead to an increased rate of reaction.

For solid reagents, we can increase their surface area. Using a reagent in a big, solid lump means that only the particles at the surface are exposed and available for reaction initially. By dividing the lump into smaller, finer pieces, we increase the number of particles available for reaction. This increases the frequency of collisions, and hence the rate. Again, the energy of the collisions is unaffected, so the proportion of successful collisions remains the same.

For gaseous reagents we can increase the pressure at which the reaction is being carried out. Put simply, squeezing the same amount of a gas into a smaller space is a method of increasing pressure. This forces the gas molecules closer together, so the collisions between their particles become more frequent, increasing the rate of the reaction. Increasing the pressure doesn’t affect the energy of the particles, so the proportion of successful collisions remains constant.

The final way we can affect the rate of a reaction is to use a catalyst. A catalyst is a chemical or biological agent that speeds up the rate of a reaction, without itself being used up in the process. It provides an alternative way for the reaction to occur, which has a lower activation energy. This means that less energy is required for two particles to collide and successfully react, so the proportion of successful collisions increases, increasing the speed with which the reaction occurs.

After all this you might be wondering why this would be of any relevance to anyone outside of learning chemistry. However, it’s much more relevant than it might initially seem. For one, the enzymes on which many of our bodily processes depends are catalysts, allowing reactions in our bodies to happen faster. Concentration and surface area, meanwhile, are both important in the context of medicine; drug doses are designed to lead to a specific concentration of the drug in our blood streams, a concentration at which the medicine is most effective, and medicines will often be powdered rather than in solid lumps to allow them to exert their effects more quickly.

Elsewhere, the kitchen also provides ample examples of how you can leverage knowledge of the factors affecting the rate of a reaction to your benefit. Always having problems with crying whilst chopping onions? Putting the onion in the fridge before cutting it lowers its temperature, slowing the rate of the reaction that produces the chemical that induces the tears. In fact, the reason we keep food in the fridge in the first place is of course to slow the rate of the spoilage of food.

There are occasionally other factors, in addition to those detailed here, that can affect the rate of a reaction. These include light intensity, the physical state of the reagents, and for some reactions the solvent in which the reaction is occurring. Because most pre-university chemistry courses here in the UK don’t discuss these in any great detail they aren’t included in this graphic, but they’re also worth being aware of.

Enjoyed this post & graphic? Consider supporting Compound Interest on Patreon, and get previews of upcoming posts & more!

DOWNLOAD

SUBSCRIBE

The graphic in this article is licensed under a  Creative Commons Attribution-NonCommercial-NoDerivatives 4.0 International License. See the site’s content usage guidelines.

 

References & Further Reading