Chemistry (0620) | Solids, liquids and gases

Hello IGCSE Chemists! Understanding the World Around You

Welcome to the fascinating world of States of Matter! This chapter is all about understanding the physical world—how materials exist as solids, liquids, or gases, and how they transition between these forms. Everything we see, touch, and breathe is made of tiny particles, and chemistry explains how their behavior changes with energy.

Don't worry if this seems tricky at first. We will break down these concepts using the idea that everything is made of moving particles: the Kinetic Particle Theory (KPT). Master KPT, and this chapter becomes simple!

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1.1 Solids, Liquids, and Gases: Distinguishing Properties and Structure

The Big Idea: Kinetic Particle Theory (KPT)

The Kinetic Particle Theory (KPT) is the foundation of this topic. It states that all matter is made up of tiny particles (atoms, molecules, or ions) that are in constant, random motion. The state of matter (solid, liquid, or gas) depends entirely on how much energy these particles have and how strongly they are attracted to each other.

Comparing the Three States (Core Content)

You need to know the basic differences between the three states in terms of volume, shape, and compressibility.

• Solids: Fixed shape and fixed volume. Cannot be compressed easily.
• Liquids: No fixed shape (takes the shape of the container) but fixed volume. Cannot be compressed easily.
• Gases: No fixed shape and no fixed volume (fills the entire container). Highly compressible.

Particle Structure, Arrangement, and Motion

We use KPT to describe the internal structure of each state:

State	Particle Separation	Particle Arrangement	Particle Movement (Motion)
Solid	Very close together, touching.	Regular, fixed pattern (lattice).	Vibrate about fixed positions. Low kinetic energy.
Liquid	Close together, touching.	Random arrangement (no fixed pattern).	Slide over each other. Moderate kinetic energy.
Gas	Far apart. Large spaces between particles.	Completely random.	Move rapidly and randomly in all directions. High kinetic energy.

Analogy Tip:
Think of a movie theatre:
Solid: Everyone is sitting neatly in their assigned seats (fixed arrangement, vibrating).
Liquid: Everyone is mingling in the foyer, close together but free to move past each other (random, sliding).
Gas: Everyone rushes outside into the street, completely spread out and moving fast (random, far apart).

2. Changes of State (Phase Changes)

Changes of state are physical changes—the substance itself doesn't change chemically (H₂O is still H₂O, whether ice, water, or steam), only the energy and organization of its particles change.

Definitions (Core Content)

• Melting: Solid to liquid (e.g., ice turning to water).
• Boiling: Liquid to gas throughout the liquid (occurs at a specific boiling point).
• Evaporating: Liquid to gas only at the surface, occurring below the boiling point.
• Freezing: Liquid to solid (e.g., water turning to ice).
• Condensing: Gas to liquid (e.g., steam turning back to water droplets).

Explaining Changes of State using KPT (Supplement Content)

These changes are explained by the transfer of thermal energy (heat).

Heating Processes (Energy Absorbed, Endothermic):

1. Melting: When a solid is heated, the particles gain kinetic energy and vibrate more intensely. At the melting point, they have enough energy to partially overcome the strong forces holding them in fixed positions. The regular structure breaks down, and the particles can now slide over each other (liquid state).
2. Boiling: When a liquid is heated to its boiling point, the particles gain enough energy to completely overcome the forces of attraction. They escape rapidly from the liquid to become free gas particles, moving far apart.

Cooling Processes (Energy Released, Exothermic):

1. Condensing: When a gas is cooled, the particles lose kinetic energy and slow down. The forces of attraction between particles become strong enough to hold them close together, forming a liquid.
2. Freezing: When a liquid is cooled, particles lose more kinetic energy and move slower. At the freezing point, the attractive forces lock the particles into fixed, regular positions, forming a solid.

Heating and Cooling Curves (Supplement Content)

When you heat a pure substance, the temperature increases steadily until it hits a melting or boiling point. At these points, the temperature stops rising, even though you are still supplying heat.

Why does the temperature stay constant during melting or boiling?

The added thermal energy (sometimes called latent heat) is not increasing the particles' kinetic energy (which determines temperature). Instead, this energy is being used to break the strong forces of attraction between the particles, allowing the change of state to occur.

Similarly, during freezing or condensation, energy is being released (exothermic), but the temperature remains constant as the new forces of attraction are being formed.

3. Gases and External Factors

Because gas particles are far apart and move rapidly, they are greatly affected by changes in temperature and pressure (1.1 Core & Supplement).

Effect of Temperature on Gas Volume (If Pressure is Constant)

Core Description: Increasing the temperature increases the volume of a gas.

KPT Explanation (Supplement):

1. When you heat a gas, the particles absorb thermal energy, converting it into kinetic energy.
2. The particles move faster and collide with the container walls more frequently and more forcefully.
3. To keep the pressure constant (as in a balloon or movable piston), the volume must increase, allowing the particles more space to spread out and hit the walls less often per unit area.

Effect of Pressure on Gas Volume (If Temperature is Constant)

Core Description: Increasing the pressure decreases the volume of a gas.

KPT Explanation (Supplement):

1. Pressure is caused by gas particles colliding with the container walls.
2. When you increase the external pressure on a gas (e.g., pushing down on a piston), the space available to the particles (the volume) decreases.
3. Because the same number of particles is now forced into a smaller space, they collide with the walls more often, maintaining the higher pressure.

Common Mistake to Avoid: When discussing gas pressure, remember it's the collisions with the container walls that cause pressure, not collisions between the gas particles themselves!

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4. Diffusion (Spreading Out)

What is Diffusion? (Core Content)

Diffusion is the net movement of particles from a region of higher concentration to a region of lower concentration, down a concentration gradient, due to their random motion.

• Example: If someone sprays perfume in one corner of a room, the scent eventually spreads everywhere. The perfume molecules move from where they are concentrated (the corner) to where they are not concentrated (the rest of the room).

Diffusion in terms of KPT (Core Content)

Diffusion happens because particles are constantly moving randomly. In liquids and gases, there are enough spaces and enough kinetic energy for the particles to spread out and mix completely.

Did you know? Diffusion is much faster in gases than in liquids because gas particles are much further apart, meaning they travel longer distances before colliding with another particle and changing direction.

Factors Affecting the Rate of Diffusion in Gases (Supplement Content)

The speed at which a gas diffuses is related to the mass of its particles.

Rule: Lighter gas particles diffuse faster than heavier gas particles.

Why?

At the same temperature, all gas particles have the same average kinetic energy. The formula for kinetic energy (KE) is:

$$KE = \frac{1}{2} m v^2$$

Where $m$ is mass and $v$ is velocity (speed).

If KE is constant, then if the mass ($m$) is smaller, the velocity ($v$) must be greater to keep the energy equal. Therefore, lighter particles move faster and spread out quicker.

• Practical Example: Ammonia gas (NH₃, \(M_r=17\)) diffuses much faster than hydrogen chloride gas (HCl, \(M_r=36.5\)). If you put them at opposite ends of a tube, the white solid product (ammonium chloride) will form closer to the heavier HCl end, showing the lighter NH₃ travelled further in the same time.

Key Concept: Relative Molecular Mass (M_r)

You must be able to calculate the $\mathbf{M_r}$ of a gas to predict its rate of diffusion. The smaller the $\mathbf{M_r}$, the faster the rate of diffusion.