Kinetic particle model of matter - Physics (0625) - Cambridge IGCSE

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The Kinetic Particle Model of Matter: Study Notes (IGCSE Physics 0625)

Hello future physicist! This chapter is all about understanding what everything around us—solids, liquids, and gases—is made of, and how they behave. This is the foundation of the "Thermal Physics" section, explaining why things heat up, cool down, and change state. It’s simple, intuitive, and extremely important!

1. The Core Idea: The Kinetic Model

The Kinetic Particle Model is a scientific theory that explains the properties of solids, liquids, and gases based on how their particles (atoms or molecules) move.
The main idea is straightforward: All matter is made up of tiny particles that are constantly moving.

Key Relationship: Motion and Temperature (Core 2.1.2/2)

The speed of the particles is directly linked to the temperature of the substance:

When a substance is heated, the particles gain energy.
They move faster (their average kinetic energy increases).
We measure this increased motion as a rise in temperature.

Analogy: Think of people in a room. If the temperature is low, they stand still or walk slowly. If the temperature rises, they start running, jumping, and moving chaotically!

Absolute Zero

If heat makes particles move, what happens when it gets as cold as possible?

There is a lowest possible temperature where the particles have the least possible kinetic energy. They are still vibrating, but their movement cannot decrease further.
This temperature is called Absolute Zero.
Absolute Zero is $0 \text{ K}$ (Kelvin) or $-273 \text{ }^{\circ}\text{C}$ (Celsius).

Quick Review: Temperature is a measure of the average kinetic energy of the particles.

2. The Three States of Matter (Core 2.1.1/1 & 2.1.2/1)

The model describes solids, liquids, and gases based on three things: the arrangement, the separation, and the motion of their particles.

2.1 Solids

Arrangement: Regular, fixed pattern (a lattice).
Separation: Very close together.
Motion: They only vibrate about fixed positions. They cannot move past each other.
Forces (Supplement 2.1.2/6): Very strong attractive forces hold them firmly in place.

Properties: Solids have a fixed shape and a fixed volume. They are hard to compress.

2.2 Liquids

Arrangement: Random arrangement. No fixed pattern.
Separation: Still very close together (similar density to solids), but slightly further apart than in a solid.
Motion: They move randomly and quickly, sliding past one another.
Forces (Supplement 2.1.2/6): Strong attractive forces still exist, but they are weak enough to allow particles to change places.

Properties: Liquids have a fixed volume but take the shape of their container. They are hard to compress.

2.3 Gases

Arrangement: Completely random.
Separation: Far apart—the average separation is much greater than the size of the particles themselves.
Motion: They move rapidly and randomly in all directions.
Forces (Supplement 2.1.2/6): Attractive forces are negligible (almost zero) because the particles are so far apart.

Properties: Gases have no fixed shape and no fixed volume. They fill their container and are easy to compress.

Memory Aid (Comparison Table):

State	Arrangement	Separation	Motion
Solid	Regular (Fixed Lattice)	Very close	Vibration only
Liquid	Random	Close	Slide past each other
Gas	Random	Far apart	Rapid and random

3. Changes of State (Phase Changes) (Core 2.1.1/2 & 2.2.3/3)

Changes of state occur when a substance absorbs or releases energy, causing the particle arrangement or motion to change.

Melting: Solid to Liquid. (Energy is absorbed, overcoming fixed lattice forces.)
Boiling/Evaporation: Liquid to Gas. (Energy is absorbed, overcoming all attractive forces.)
Condensation: Gas to Liquid. (Energy is released, particles lose speed and attractive forces pull them closer.)
Solidification (Freezing): Liquid to Solid. (Energy is released, particles slow down and settle into fixed positions.)

Note: You do not need to know the terms for gas to solid (deposition) or solid to gas (sublimation) for this syllabus section.

4. Pressure in Gases (Core 2.1.2/3 & Supplement 2.1.2/7)

Gases exert pressure on the walls of their container. The kinetic model explains exactly why this happens:

Gas particles are moving very fast and randomly.
They frequently collide with the interior walls of the container.
Each collision involves a change in the particle's momentum, which means a tiny force is exerted on the wall.
Since there are trillions of particles constantly colliding, the combined effect of all these tiny forces spread over the wall’s area creates a constant measurable pressure.

Recall the definition of pressure: $p = \frac{F}{A}$. In gases, the collisions provide the force (F) acting on the area (A) of the container wall.

What affects Gas Pressure? (Core 2.1.3/1)

We need to understand how two things affect the pressure of a fixed mass of gas:

1. Changing Temperature (at constant volume)

If you increase the temperature, the particles move faster.
Faster particles hit the walls more frequently and harder.
Result: The pressure increases.

Example: Leaving a deodorant can in a hot car. The gas inside heats up, the pressure rises, and the can might explode!

2. Changing Volume (at constant temperature)

If you decrease the volume (make the container smaller), the particles have less space to move.
The particles hit the walls more frequently because the walls are closer together.
Result: The pressure increases.

Supplement Focus: Boyle’s Law (2.1.3/3)

For a fixed mass of gas held at constant temperature, the pressure is inversely proportional to the volume.
Mathematically, this is expressed as: $$p V = \text{constant}$$

This means if you double the volume, the pressure halves, provided the temperature doesn't change. When plotting pressure (p) against volume (V), you get a curve (inversely proportional). If you plot pressure (p) against $1/V$, you get a straight line passing through the origin.

5. Evidence for the Kinetic Model: Brownian Motion (Core 2.1.2/4 & 5)

How do we know the particles are *really* moving randomly? The proof comes from observing Brownian Motion.

Did you know? Brownian motion was first observed by botanist Robert Brown in 1827 while looking at pollen grains in water, but he didn't know *why* they moved. Einstein later confirmed the kinetic explanation.

The Observation:

If you look at tiny smoke particles suspended in air, or tiny pollen particles in water, under a microscope, you will see them moving in a jerky, random, zig-zag path.

The Explanation: (Crucial Distinction!)

The key to understanding Brownian motion is distinguishing between the two types of particles involved (Supplement 2.1.2/8):

Microscopic Particles (e.g., smoke or pollen, which are visible under the microscope).
Atoms or Molecules (the invisible particles of the surrounding gas or liquid).

Step-by-Step Explanation:

The atoms/molecules of the gas or liquid are moving extremely fast and randomly.
These fast-moving, invisible particles constantly collide with the much larger microscopic particle (like a smoke particle).
Because the collisions happen randomly and unevenly on all sides, the microscopic particle is knocked first one way, then another, causing the observable jerky, random motion.

Summary: The random movement of visible microscopic particles (Brownian motion) is caused by the random, frequent collisions with the invisible, fast-moving atoms or molecules of the surrounding fluid, providing strong evidence for the kinetic particle model of matter.

6. The Absolute Scale of Temperature (Core 2.1.3/2)

In Physics, especially when dealing with gases, we must use the Kelvin (K) scale, often called the Absolute Scale.

Why? Because the Kelvin scale starts at Absolute Zero ($0 \text{ K}$), meaning the temperature reading is directly proportional to the average kinetic energy of the particles. You can't say a gas at $50 \text{ }^{\circ}\text{C}$ has twice the kinetic energy of a gas at $25 \text{ }^{\circ}\text{C}$, but you *can* say a gas at $300 \text{ K}$ has twice the kinetic energy of a gas at $150 \text{ K}$.

Converting between Kelvin and Celsius

The Kelvin scale uses the same size divisions as the Celsius scale, but it is shifted by 273 degrees.

To convert Celsius ($\theta$) to Kelvin ($T$): $$T \text{ (in K)} = \theta \text{ (in }^{\circ}\text{C)} + 273$$

Example:

Room temperature ($20 \text{ }^{\circ}\text{C}$) in Kelvin is: $20 + 273 = 293 \text{ K}$.
The boiling point of water ($100 \text{ }^{\circ}\text{C}$) in Kelvin is: $100 + 273 = 373 \text{ K}$.

Common Mistake Alert! Always remember to convert Celsius temperatures to Kelvin before using them in gas law calculations (like Boyle's Law, even though Boyle's Law specifies constant temperature, other gas laws you might encounter later rely on Kelvin).

Key Takeaway Summary

The kinetic particle model links particle motion directly to temperature. Solids vibrate in fixed positions, liquids slide past each other, and gases move randomly and far apart. Gas pressure is caused by collisions, and this model is confirmed by Brownian Motion, the random movement caused by tiny, invisible particle impacts. Finally, remember to use the Kelvin scale, which starts at absolute zero ($-273 \text{ }^{\circ}\text{C}$).