Chemistry (0620) | Rate of reaction

Welcome to the World of Reaction Rates!

Have you ever noticed that some chemical changes happen almost instantly, like lighting a match, while others take days or years, like iron rusting? This chapter, Rate of reaction, is all about understanding how fast chemical processes occur and, more importantly, how we can speed them up or slow them down.

This topic is vital, not just for exams, but for real-life applications, from manufacturing plastics quickly to safely storing medicines that must react slowly! Don't worry if the concepts seem a little abstract at first; we will use plenty of analogies to make sense of the chemical world.

1. What is the Rate of Reaction?

1.1 Defining Rate and Effective Collisions

The rate of reaction simply means how quickly reactants are used up, or how quickly products are formed. It is measured as the change in concentration or quantity of a substance over time.

$Rate = \frac{Change \ in \ amount \ of \ substance}{Time \ taken}$

The Requirement for Reaction (Collision Theory Basics)

For any chemical reaction to occur, the reacting particles (atoms, ions, or molecules) must follow two simple rules. This is called the Collision Theory (Supplement 5):

1. Collision: The particles must first collide with each other.
2. Effective Collision: The collision must happen with the correct orientation AND with enough energy.

A collision that meets both conditions (correct angle and sufficient energy) is called an effective collision. Only effective collisions lead to the formation of products.

Key Takeaway: The faster the rate of reaction, the greater the frequency of effective collisions.

1.2 Measuring the Rate of Reaction (Core 3, 4)

To measure the rate in the lab, we usually track an easily observable change.

Practical Methods for Investigating Rate

We investigate the rate of reaction by monitoring the change in a physical property that occurs as the reaction proceeds:

• If a gas is produced:
  We can measure the volume of gas produced over time using a gas syringe.
  Example: Reaction between magnesium and acid.
• If mass is lost (due to gas escaping):
  We can measure the decrease in mass over time by carrying out the reaction on a balance. (The cotton wool plug allows gas to escape but prevents acid spray from leaving).
  Example: Reaction between calcium carbonate (marble chips) and dilute acid.
• If the solution becomes cloudy (forms a precipitate):
  We can measure the time taken for a cross drawn underneath the reaction vessel to disappear (become obscured). This is called the time taken to form a specified amount of precipitate.
  Example: Sodium thiosulfate reacting with dilute acid.

Interpreting Graphs from Rate Experiments

When you plot the amount of product formed (y-axis) against time (x-axis), the graph always has a characteristic shape:

• The graph starts steep (high gradient) $\rightarrow$ The reaction is fastest at the start (when reactant concentration is highest).
• The slope gradually becomes less steep $\rightarrow$ The reaction slows down as reactants are used up.
• The slope becomes horizontal (zero gradient) $\rightarrow$ The reaction has stopped (one reactant is fully used up).

Analogy: Racing a Car
Imagine a speed graph. When the line goes up steeply, you are accelerating quickly (fast rate). When the line flattens out, you have reached your cruising speed (reaction ending).

2. The Energy Barrier: Activation Energy

2.1 Defining Activation Energy (\(E_a\)) (Supplement 5)

In order for particles to collide effectively and break old bonds to form new ones, they need a certain minimum amount of energy.

The Activation Energy (\(E_a\)) is defined as the minimum energy that colliding particles must possess in order to react (Supplement 5d).

2.2 Analogy for Activation Energy

Imagine you need to push a rock over a hill for it to roll down and complete the reaction.

• The height of the hill is the Activation Energy, \(E_a\).
• If you only push the rock halfway up the hill (low energy), it rolls back down (no reaction).
• You must give the rock enough energy (a strong push) to get it over the top (the \(E_a\)) before the reaction can proceed.

This energy barrier explains why mixing fuel and oxygen isn't enough; you need a spark (which provides the \(E_a\)) to start the combustion reaction.

3. Factors Affecting the Rate of Reaction

We can change the rate of reaction by adjusting conditions. Remember, all these explanations rely on increasing the frequency of effective collisions (Core 1, Supplement 6).

3.1 Concentration (Solutions) and Pressure (Gases)

Increasing the concentration of solutions or the pressure of gases increases the rate of reaction.

Explanation using Collision Theory:

• Increase Concentration/Pressure: This means you have a greater number of particles per unit volume (they are packed closer together).
• Result: The particles are more likely to bump into each other. This increases the frequency of collisions.
• Outcome: A higher frequency of effective collisions, thus a faster rate.

Did you know? In industry, reactions often run under high pressure to increase the rate, even though building high-pressure equipment is expensive. It saves time, making the process more economical!

3.2 Surface Area (Solids)

For reactions involving a solid reactant, breaking the solid into smaller pieces increases the rate.

Explanation using Collision Theory:

• Increase Surface Area: When a solid lump is crushed into a powder, the total area exposed to the other reactant (the liquid or gas) increases dramatically.
• Result: More reactant particles (in the liquid/gas) can collide with the solid particles simultaneously.
• Outcome: This increases the frequency of collisions between the phases, leading to a faster rate.

Analogy: Digestion. We chew food (increasing surface area) so digestive enzymes (reactants) can work quickly!

3.3 Temperature

Increasing the temperature almost always increases the rate of reaction significantly.

Explanation using Collision Theory:

When you heat up reactants, you give the particles more energy (increased kinetic energy). This speeds up the rate for two reasons:

1. Higher Collision Frequency (Minor Effect): The particles move faster, so they collide more frequently.
2. More Effective Collisions (Major Effect): Because the particles have more energy, a much greater proportion of particles now have energy equal to or greater than the Activation Energy (\(E_a\)).

Crucial Point: The second factor (getting over the \(E_a\) barrier) is far more important than the first. A small temperature rise can double the reaction rate because so many more particles now have the necessary energy for an effective collision.

Memory Aid: Think of a stadium crowd. If the crowd is warm (high T), they are running around faster (more collisions) and they have the energy to jump the fence (\(E_a\))!

3.4 Catalysts and Enzymes

A catalyst is a substance that increases the rate of a chemical reaction but is itself chemically unchanged at the end of the reaction (Core 2).

Enzymes are biological catalysts (found in living organisms).

Explanation using Collision Theory:

• Role of the Catalyst: A catalyst works by providing an alternative reaction pathway that requires a lower Activation Energy, \(E_a\) (Supplement 7).
• Result: Since the energy barrier is now lower, a much larger proportion of the reactant particles possess the minimum energy needed to react, even at the original temperature.
• Outcome: This drastically increases the frequency of effective collisions, leading to a much faster rate.

Analogy: The Shortcut. If the Activation Energy is a large mountain, the catalyst digs a tunnel through it, making it easier for everyone (all the particles) to get to the other side quickly!

4. Summarising the Factors (Extended Level Synthesis)

When explaining why a factor affects the rate, always use these key terms:

1. Describe the change (e.g., "Increasing concentration means more particles per unit volume.")
2. Explain the effect on collisions (e.g., "This increases the frequency of collisions.")
3. Explain the effect on \(E_a\) (if applicable, usually temperature or catalyst).
4. State the final outcome (e.g., "...therefore, the frequency of effective collisions increases, and the rate is faster.")

Common Mistakes to Avoid

• Mistake: Saying "Temperature increases activation energy."
  Correction: Temperature increases the *kinetic energy* of particles. Only a catalyst can change the Activation Energy (\(E_a\)).

• Mistake: Saying "More collisions happen."
  Correction: Always use the term "frequency of collisions" (collisions per unit time) to show you understand that rate is dependent on time.

Congratulations! You now understand the essential principles behind how fast chemical reactions occur, governed by the crucial concept of effective collisions and the energy barrier, \(E_a\).