Rates of Reaction (Speed of Reaction)
What Does "Speed of Reaction" Actually Mean?
Imagine a race between two people. The one who reaches the finish line first has the greater speed. In a chemical reaction, the 'race' is the transformation of starting materials (reactants) into new substances (products). The rate of reaction is a measure of how quickly this transformation happens.
We can define it more precisely as the change in the amount (concentration, mass, volume) of a reactant or product per unit of time. For example, if we are producing a gas, we could measure the rate by how much gas is collected every second. The key idea is change over time. A fast reaction has a large change in a short time, while a slow reaction has a small change over a long period.
Consider the classic reaction between an acid and a metal, like hydrochloric acid and magnesium ribbon: $ Mg_{(s)} + 2HCl_{(aq)} \rightarrow MgCl_{2(aq)} + H_{2(g)} $ . The fizzing you see is hydrogen gas ($ H_2 $) being produced rapidly. We could measure the rate by timing how long it takes for the magnesium to completely disappear, or by measuring the volume of gas produced every 10 seconds.
The Collision Theory: Why Reactions Happen
To understand why reactions have different speeds, we need the Collision Theory. This theory states that for a reaction to occur:
- Particles must collide. Atoms, molecules, or ions must bump into each other.
- The collisions must have sufficient energy. This minimum energy required is called the activation energy ($ E_a $).
- The collisions must have the correct orientation. The particles must hit each other in the right way for bonds to break and form.
Think of trying to open a door with a key. You need to collide the key with the lock (collision), you need to push with enough force to turn it (energy > $ E_a $), and you need the key to be the right way up (correct orientation). If any of these conditions aren't met, the door won't open (the reaction won't happen).
Factors That Control the Speed of a Reaction
Several factors influence the rate of a reaction by changing the frequency or effectiveness of collisions between particles. Let's explore the main ones.
1. Concentration (or Pressure for Gases)
Increasing the concentration of reactants means there are more reactant particles in the same volume. This leads to more frequent collisions, and therefore, more effective collisions per second. The reaction rate increases.
Example: A glowing splint will burst into flames more quickly in pure oxygen (high concentration) than in air (lower oxygen concentration).
2. Temperature
This is often the most dramatic factor. Increasing temperature does two things: it makes particles move faster, increasing collision frequency, and, more importantly, it gives a much larger proportion of the particles energy equal to or greater than the activation energy ($ E_a $). Even a small rise in temperature can cause a large increase in reaction rate.
Example: Food spoils slowly in the fridge (low temperature) but very quickly if left on a hot kitchen counter (high temperature).
3. Surface Area (for Solids)
If a reactant is a solid, only the particles on the surface are available for collision. Breaking a solid into smaller pieces (or powdering it) vastly increases its surface area. This exposes more particles, leading to more frequent collisions and a faster rate.
Example: A sugar cube dissolves slowly in tea, but the same mass of granulated sugar dissolves almost instantly because of its much larger surface area.
4. Catalysts
A catalyst is a substance that speeds up a chemical reaction without being permanently changed or used up. It works by providing an alternative reaction pathway with a lower activation energy ($ E_a $). With a lower energy barrier, a greater proportion of collisions become effective at any given temperature.
Example: In a car's catalytic converter, platinum and rhodium metals catalyze the conversion of harmful exhaust gases (like carbon monoxide) into less harmful substances (like carbon dioxide).
5. Nature of the Reactants
The type of substances involved affects the rate. Reactions involving simple ions in solution are often very fast, while those involving breaking strong covalent bonds within molecules are typically slower.
Example: The reaction between silver ions and chloride ions to form a white precipitate of silver chloride is instantaneous. The rusting of iron (a reaction with oxygen) is much slower.
| Factor | Effect on Rate | Reason (Collision Theory) | Practical Example |
|---|---|---|---|
| Concentration / Pressure | Increases | More particles per volume = more frequent collisions. | Higher concentration of acid dissolves metal faster. |
| Temperature | Increases greatly | Particles have more kinetic energy; a greater proportion have energy ≥ $ E_a $. | Cooking food uses heat to speed up chemical changes. |
| Surface Area | Increases | More exposed particles = more collision sites. | Powdered calcium carbonate fizzes more violently in acid than a single lump. |
| Catalyst | Increases | Lowers the activation energy ($ E_a $), so more collisions are effective. | Enzymes in your body catalyze the digestion of food. |
| Nature of Reactants | Varies | Strong bonds require more energy to break; weak bonds react more easily. | Sodium metal reacts violently with water; gold does not react at all. |
How to Measure the Rate of a Reaction
Scientists measure reaction rates by monitoring the change in quantity of a reactant or product over time. The method chosen depends on the reaction and what property changes conveniently.
| Method | What is Measured | Example Reaction | Calculation of Rate |
|---|---|---|---|
| Loss of Mass | Mass decrease (if a gas is produced) | Acid + Metal Carbonate → Salt + Water + CO2(g) | Rate = (Mass loss in grams) / (Time in seconds) = g/s |
| Volume of Gas | Volume of gas produced | Hydrogen Peroxide → Water + Oxygen (with a catalyst) | Rate = (Volume in cm3) / (Time in seconds) = cm3/s |
| Color Change | Light absorption using a colorimeter | Sodium thiosulfate + acid → cloudy yellow precipitate | Rate ∝ 1 / (Time for cross to disappear) |
| pH Change | Change in acidity using a pH probe | Acid + Alkali → Salt + Water (neutralization) | Rate = Δ[H+] / Δt (Change in H+ ion concentration over time) |
Real-World Applications: From Kitchens to Factories
Understanding and controlling reaction rates is vital in countless everyday and industrial processes. Here are a few detailed applications.
Food Preservation: We use low temperatures (refrigeration, freezing) to slow down the chemical reactions and microbial growth that cause food to spoil. Canning removes oxygen and uses heat to deactivate enzymes[1], creating a sterile environment where decay reactions are extremely slow.
Combustion Engines: In a car engine, fuel and air must mix and burn extremely rapidly to produce the power to move the car. The rate is controlled by the fuel injectors, air intake, and spark timing. Catalytic converters then use catalysts to speed up the reactions that convert polluting exhaust gases into safer ones.
Medicines and Drugs: Pharmaceutical companies carefully study the rate at which a drug breaks down in the body to determine the correct dosage and how often it should be taken. A slow-release medication is designed to react with body fluids at a controlled, slow rate to provide a steady effect.
Manufacturing Chemicals: The Haber Process for making ammonia ($ N_2 + 3H_2 \rightleftharpoons 2NH_3 $) is a classic example. To achieve an economically viable rate, factories use high pressure (to increase concentration of gases), a high temperature (to speed up particle collisions), and an iron catalyst (to lower the activation energy).
Important Questions
A: Increasing concentration only increases the frequency of collisions. Increasing temperature does two things: it increases collision frequency and, more importantly, it dramatically increases the fraction of particles with kinetic energy greater than or equal to the activation energy ($ E_a $). This means a much higher percentage of collisions are successful. A common rule is that a $ 10^{\\circ}C $ rise often doubles the reaction rate.
A: A catalyst remains chemically unchanged at the end of the reaction. It is often a solid that provides a surface for the reaction to occur on, or a substance that temporarily combines with reactants to form an intermediate compound. Once the products are formed, the catalyst is released back into its original form, ready to be used again. For example, the manganese(IV) oxide ($ MnO_2 $) catalyst used to decompose hydrogen peroxide will be left as a black solid at the bottom of the test tube after all the gas has been produced.
A: No, a reaction cannot go on forever. As reactants are used up and turned into products, their concentration decreases. According to collision theory, fewer reactant particles remain, so collisions become less frequent. Therefore, the rate of reaction is fastest at the beginning and gradually slows down over time until it effectively stops when one of the reactants is completely used up. A graph of amount of product vs. time starts steep (fast rate) and gradually levels off (slow rate).
Footnote
[1] Enzymes: Biological catalysts, usually proteins, that speed up biochemical reactions in living organisms. For example, amylase in saliva catalyzes the breakdown of starch into sugars.
[2] Activation Energy ($ E_a $): The minimum amount of energy that colliding particles must possess for a reaction to occur.
[3] Haber Process: An industrial chemical process that uses an iron catalyst to synthesize ammonia from nitrogen and hydrogen gases under high temperature and pressure.
[4] Catalyst: A substance that increases the rate of a chemical reaction without being consumed in the overall process.