Reaction Rate: The Speed at Which a Reaction Happens
What Exactly is Reaction Rate?
Imagine two races: a snail moving across a garden and a cheetah sprinting across the savannah. Both are moving, but their speeds are vastly different. Chemical reactions are similar. Some, like an explosion, are incredibly fast. Others, like the formation of rust on a car or the ripening of a banana, can take much longer. The reaction rate is simply a measure of this "chemical speed."
Scientifically, we can define reaction rate in a few ways. It can be the speed at which the reactants (the starting substances) are used up. Alternatively, it can be the speed at which the products (the new substances formed) are created. For example, in the reaction between hydrochloric acid and magnesium metal:
We could measure the reaction rate by observing how quickly the magnesium metal disappears (the reactant is used up) or how quickly hydrogen gas bubbles are produced (the product is formed).
The Collision Theory: Why Reactions Happen
To understand what affects reaction rate, we first need to understand why reactions happen at all. The Collision Theory provides the answer. This theory states that for a reaction to occur, the particles of the reactants (atoms, molecules, or ions) must:
- Collide with each other.
- Collide with sufficient energy (known as the activation energy).
- Collide with the correct orientation.
Think of it like trying to open a door with a key. Simply touching the key to the door (a collision) isn't enough. You need to push the key in with enough force (sufficient energy) and the key must be the right way up (correct orientation). If any of these three conditions are not met, the door won't open. Similarly, if the collision between reactant particles isn't just right, no reaction will take place. The reaction rate depends on how often successful collisions happen.
Factors That Control the Speed of a Reaction
Several factors can change the number of successful collisions, thereby speeding up or slowing down a reaction. These are the "dials" scientists and engineers can turn to control chemical processes.
1. Concentration and Pressure
For reactions involving liquids or gases, increasing the concentration of the reactants makes the particles more crowded. In a gas, increasing the pressure has a similar effect, squeezing the particles closer together. When particles are closer together, they collide more frequently. More collisions per second mean a higher chance of successful collisions, leading to a faster reaction rate.
Example: A glowing splint will glow brighter when placed in pure oxygen than in air (which is only about 21% oxygen). The higher concentration of oxygen particles leads to more frequent and energetic collisions with the splint, increasing the combustion reaction rate.
2. Surface Area
If a reaction involves a solid, only the particles on the surface of the solid are available for collision. Breaking the solid into smaller pieces dramatically increases its total surface area. This exposes more particles, making them available for collision. More available particles lead to more frequent collisions and a faster reaction.
Example: A large log burns slowly in a fireplace, but the same mass of wood in the form of small twigs or sawdust can burn very quickly, even explosively. The sawdust has a much larger surface area exposed to oxygen, allowing for a much faster reaction rate.
3. Temperature
This is often the most important factor. Increasing the temperature does two things. First, it makes particles move faster, causing them to collide more frequently. Second, and more importantly, it provides the particles with more kinetic energy. A higher proportion of the colliding particles will then have energy equal to or greater than the activation energy ($E_a$). This significant increase in successful collisions causes the reaction rate to increase dramatically.
Example: Food spoils much faster on a warm day than in a refrigerator. The lower temperature inside the fridge slows down the chemical reactions conducted by bacteria and molds, preserving the food for longer.
4. Catalysts
A catalyst is a substance that speeds up a chemical reaction without being permanently changed or used up itself. It works by providing an alternative pathway for the reaction that has a lower activation energy ($E_a$). With a lower energy barrier, a much larger fraction of collisions become successful, greatly increasing the reaction rate.
| Factor | Effect on Reaction Rate | Reason (Collision Theory) |
|---|---|---|
| Increased Concentration/Pressure | Increases | More particles in a given space leads to more frequent collisions. |
| Increased Surface Area | Increases | More particles are exposed and available for collision. |
| Increased Temperature | Increases | Particles have more energy; more collisions have energy ≥ $E_a$. |
| Addition of a Catalyst | Increases | Provides an alternative reaction pathway with a lower $E_a$. |
Reaction Rates in Action: From Biology to Industry
The principles of reaction rate are not confined to a chemistry lab; they are at work all around us and inside us.
In Our Bodies (Enzymes): The chemical reactions in living organisms must happen quickly at relatively low temperatures. This is made possible by enzymes, which are biological catalysts. For instance, the enzyme amylase in your saliva catalyzes the breakdown of starch into sugars. Without amylase, this reaction would be far too slow to provide your body with energy from food.
In Food Preservation: We use low temperatures (refrigeration and freezing) to slow down the reaction rates of decay and bacterial growth. Canning food involves heating it to destroy microorganisms and then sealing it to prevent new oxygen (a reactant) from getting in, thus slowing down spoilage reactions.
In Industrial Processes: The Haber Process[1] for producing ammonia ($N_2 + 3H_2 \rightarrow 2NH_3$) is a classic example. This reaction is naturally very slow. To make it economically viable, chemists use high pressure (to increase concentration), a high temperature (to increase energy), and an iron catalyst (to lower the activation energy). Controlling the reaction rate is essential for producing fertilizer on a global scale.
Common Mistakes and Important Questions
A: No. A catalyst only speeds up a reaction that is already possible (thermodynamically favorable). It cannot force a reaction that is impossible to occur. It simply makes the possible happen much faster.
A: Not exactly. The relationship is not linear. A common rule of thumb is that for many reactions, the rate approximately doubles for every 10°C rise in temperature. So, increasing the temperature from 20°C to 30°C might double the rate, but going from 30°C to 40°C would double it again (a four-fold increase overall).
A: For the vast majority of chemical reactions, yes. However, some physical processes and a few rare chemical reactions can slow down with increasing temperature. For the scope of school chemistry, you can assume that increasing temperature increases the reaction rate.
Footnote
[1] Haber Process: An industrial method for synthesizing ammonia from nitrogen and hydrogen gases, using an iron catalyst, high temperature, and high pressure. It is critically important for producing nitrogen-based fertilizers.
[2] Activation Energy ($E_a$): The minimum amount of energy that colliding particles must have for a reaction to occur. It is the energy barrier that must be overcome.
[3] Catalyst: A substance that increases the rate of a chemical reaction without being consumed in the process. It works by providing an alternative reaction pathway with a lower activation energy.
[4] Enzyme: A protein that acts as a biological catalyst, speeding up specific biochemical reactions within living organisms.