Diffusion: Why Smells Spread and Particles Move

Welcome to the exciting world of Diffusion! This is a fundamental concept in Chemistry, essential for understanding the properties of matter (Section 1.0 of your syllabus).
Diffusion explains everyday occurrences, like why a tiny bit of perfume sprayed in one corner of a room soon reaches the other side, or why food colouring mixes without you having to stir it.

Don't worry if the terminology seems complex; diffusion is simply about particles moving around randomly and spreading out!

1. Defining Diffusion (Core Content)

What is Diffusion?

Diffusion is the net movement of particles (atoms, ions, or molecules) from a region of high concentration to a region of low concentration, resulting in the uniform distribution of particles.

Think of it like a crowded cinema emptying out. People (particles) naturally move from the very crowded entrance (high concentration) to the empty foyer (low concentration) until they are spread out.

Key Point: Diffusion is a result of the random motion of particles, which is always occurring due to their kinetic energy (as explained by the Kinetic Particle Theory).

Quick Review: Key Terms
  • Diffusion: Spreading out of particles.
  • Concentration: How crowded the particles are in a given volume.
  • Kinetic Particle Theory: The idea that all matter is made of tiny particles that are constantly moving.

2. Explaining Diffusion using Kinetic Particle Theory (Core Content)

To explain diffusion, we must rely on the concepts introduced in Section 1.1 (States of Matter).

The syllabus requires you to describe and explain diffusion in terms of kinetic particle theory:

  1. The particles in gases and liquids are in constant, random motion. They possess kinetic energy.

  2. In a region of high concentration, particles are packed closely together. Due to their random motion, they frequently collide with each other.

  3. Although particles move in all directions, because there are statistically more particles moving out of the crowded area than moving into it, there is a net movement.

  4. The particles spread out until they are uniformly distributed, meaning the concentration is equal everywhere. At this point, movement is still happening, but there is no longer any net (overall) movement from one region to another.

Analogy: Imagine a busy school corridor and an empty classroom connected by a door. In the first minute, most students will move from the corridor (high concentration) into the classroom (low concentration). Eventually, the students will be spread out equally between both spaces, even though individuals are still running around!


3. Diffusion in Different States of Matter (Core Context)

Diffusion occurs in both liquids and gases, but the speed varies significantly because of the different arrangements of particles:

Diffusion in Gases

  • Speed: Very fast.

  • Reason: Gas particles are far apart and move very rapidly. They have large empty spaces between them, allowing them to spread out quickly.

Real-world example: Smelling cooking or chemicals from across the room.

Diffusion in Liquids

  • Speed: Much slower than gases.

  • Reason: Liquid particles are closely packed (though randomly arranged). They have to move past many neighbours, resulting in more frequent collisions, which slows down the net rate of spreading.

Practical Example: Dropping potassium manganate(VII) (purple crystals) into a beaker of water. The purple colour takes a long time (hours or days) to spread throughout the entire volume.

Diffusion in Solids

In solids, particles are held in fixed positions in a regular lattice structure. Their movement is limited to vibration. Therefore, diffusion in solids is generally negligible (does not happen in practical terms).

Key Takeaway (Core): Diffusion is the random spreading of particles from high to low concentration, driven by the constant, random movement of those particles. It is fastest in gases and slowest (or non-existent) in solids.


4. Factors Affecting the Rate of Diffusion (Supplement Content)

For students taking the Extended syllabus, you need to understand what makes some particles diffuse faster than others.

The speed of diffusion is mainly affected by two factors: temperature and particle mass.

4.1 Effect of Temperature (Core/Extension Link)

When you increase the temperature, the particles gain more kinetic energy.

  • Higher kinetic energy means the particles move faster.

  • Moving faster means they can travel from the high concentration area to the low concentration area more quickly.

Result: An increase in temperature leads to an increase in the rate of diffusion. (This is why hot tea cools quicker than cold tea—the water molecules move faster, allowing the heat energy to spread out more rapidly.)

4.2 Effect of Relative Molecular Mass (Mᵣ) on Gases (Supplement Content)

This is the most important concept for Extended candidates in this section.

You need to describe and explain the effect of relative molecular mass on the rate of diffusion of gases.

The Rule: Lighter is Faster

For gases at the same temperature, particles with a lower relative molecular mass (\( M_r \)) diffuse faster.

The Scientific Explanation:

  1. Temperature is a measure of the average kinetic energy (KE) of the particles.

  2. If two gases are at the same temperature, their particles must have the same average KE.

  3. Kinetic energy is calculated using the formula: \( \text{KE} = \frac{1}{2} \text{mass} \times \text{velocity}^2 \) or \( \text{KE} = \frac{1}{2}mv^2 \).

  4. If the KE is constant, a particle with a smaller mass (\( m \)) must have a higher velocity (\( v \)) to keep the equation balanced.

Therefore, lighter particles move at a greater speed, leading to a faster rate of diffusion.

The Classic Diffusion Experiment (\( \text{NH}_3 \) vs \( \text{HCl} \))

A common experiment to prove this relationship involves the diffusion of two gases: ammonia (\( \text{NH}_3 \)) and hydrogen chloride (\( \text{HCl} \)).

Step 1: The Setup
Two cotton wool swabs are soaked, one in concentrated aqueous ammonia (which releases \( \text{NH}_3 \) gas) and one in concentrated hydrochloric acid (which releases \( \text{HCl} \) gas). These are placed at opposite ends of a long glass tube simultaneously.

Step 2: The Reaction
The two gases diffuse towards the centre of the tube. When they meet, they react to form a visible white ring of solid ammonium chloride:
\( \text{NH}_3(g) + \text{HCl}(g) \rightarrow \text{NH}_4\text{Cl}(s) \)

Step 3: Calculating Relative Molecular Mass (\( M_r \))

  • Ammonia (\( \text{NH}_3 \)): \( 14 + (3 \times 1) = 17 \)

  • Hydrogen Chloride (\( \text{HCl} \)): \( 1 + 35.5 = 36.5 \)

Step 4: The Result and Explanation
The white ring of ammonium chloride forms closer to the \( \text{HCl} \) end of the tube.

  • Observation: The ammonia gas travelled a longer distance than the hydrogen chloride gas in the same amount of time.

  • Conclusion: Since \( \text{NH}_3 \) (Mᵣ = 17) is much lighter than \( \text{HCl} \) (Mᵣ = 36.5), the ammonia molecules move faster and diffuse more quickly.

Did you know? This relationship between diffusion rate and mass is formally known as Graham’s Law of Diffusion. For IGCSE, simply understanding the "lighter is faster" principle is sufficient.


5. Common Mistakes and Quick Review

Common Misconception Alert!

When asked to explain diffusion, many students mistakenly say that particles stop moving once they are spread out evenly.

Mistake to Avoid:
The particles never stop moving. Even when the concentration is uniform, the particles are still in constant, random motion. We say that the diffusion has stopped because the net movement (the overall flow from one side to the other) has become zero.

Quick Review Checklist

Use this checklist to ensure you have mastered the chapter:

  • Can I define diffusion clearly?

  • Can I explain diffusion using the idea of random particle motion and concentration? (Core)

  • Can I explain why diffusion is slower in liquids than in gases? (Core)

  • Can I state the relationship between a gas's mass and its diffusion rate? (Supplement: Lighter = Faster)

  • Can I explain *why* lighter particles diffuse faster in terms of kinetic energy (\( \text{KE} = \frac{1}{2}mv^2 \))? (Supplement)

Key Takeaway (Supplement): The rate of diffusion is inversely related to the mass of the particles. Lighter gas particles move faster at the same temperature, causing them to spread out more rapidly.