Know that microscopic particles may be moved by collisions with light fast-moving molecules and correctly use the terms atoms or molecules as distinct from microscopic particles

2.1.2 Particle Model

Learning objectives

• Explain how microscopic particles are set into motion by collisions with fast‑moving molecules or atoms.
• Describe the three states of matter and the way particle arrangement changes between them.
• Link temperature to the average kinetic energy of particles and state the significance of absolute zero.
• Explain qualitatively why pressure is produced by particle collisions with the walls of a container.
• Distinguish clearly between atoms, molecules and the broader term microscopic particles.
• Recall and use the ideal‑gas relationship \(pV = \text{constant}\) (for a fixed mass of gas at constant temperature) and the temperature‑conversion formula.

Core concepts

• All matter consists of particles that are far too small to be seen with the naked eye.
• These particles are in constant random motion; the average speed increases as the temperature rises.
• Higher temperature → higher average kinetic energy → more energetic and more frequent collisions.
• At absolute zero (‑273 °C or 0 K) the kinetic energy of particles is at its minimum; in the idealised picture motion ceases.

States of matter – particle arrangement

The particle model distinguishes three familiar states of matter. The schematic below (suggested diagram) shows the relative spacing and freedom of movement.

• Solid: particles vibrate about fixed positions in a regular lattice; strong intermolecular forces keep them tightly packed.
• Liquid: particles are close together but can move past each other, giving a definite volume but no fixed shape.
• Gas: particles are widely spaced and move independently, filling any container.

Temperature, kinetic energy and absolute zero

The average kinetic energy of a particle is given (qualitatively) by

higher temperature ⇔ higher average kinetic energy ⇔ faster particles.

When the temperature reaches absolute zero the kinetic energy is at its lowest possible value; in the ideal gas model the particles would be at rest.

Pressure – a result of particle collisions

Pressure is the force exerted per unit area on the walls of a container. It arises because particles constantly strike the walls. The more frequent and the more energetic the impacts, the larger the pressure.

Real‑world links

• Hot‑air balloon: heating the air inside the envelope raises the kinetic energy of the gas particles, increasing the pressure and causing the balloon to expand and become less dense than the surrounding air.
• Tyre pressure: on a cold morning the air in a tyre is cooler, so the particles move more slowly and the pressure falls; warming the tyre (e.g., after a long drive) raises the pressure.
• Pressure cooker: heating the sealed steam raises the kinetic energy of the water vapour, increasing the pressure and allowing food to cook faster.

Atoms, molecules and microscopic particles

It is essential to keep the three terms distinct:

• Atoms are the smallest units of a chemical element that retain the element’s properties. They consist of a nucleus surrounded by electrons. Example: a hydrogen atom (H).
• Molecules are two or more atoms chemically bonded together. They can be elemental (O₂) or compound (H₂O, CO₂). Molecules are larger than individual atoms but still microscopic.
• Microscopic particles is the broad umbrella term that includes atoms, molecules, ions, electrons, neutrons, dust grains, pollen grains, etc.—any particle too small to be seen directly.

Collisions transfer motion

When a fast‑moving molecule (or atom) collides with a stationary microscopic particle, part of its momentum is transferred, causing the target to move. In an elastic collision the speed of the initially stationary particle becomes

\[

v{\text{target}} = \frac{2m{\text{fast}}}{m{\text{fast}}+m{\text{target}}}\;u_{\text{fast}}

\]

For the purposes of the IGCSE/A‑Level syllabus it is enough to understand the qualitative idea: the moving particle “kicks” the stationary one, and the energy transferred depends on their relative masses.

Brownian motion – evidence for the kinetic particle model

Brownian motion is the observable jittery movement of tiny suspended particles (e.g., pollen grains) in a liquid or gas. It provides direct macroscopic evidence that:

• Fluid molecules are in constant rapid motion.
• Countless random collisions with these molecules continually transfer tiny amounts of momentum to the suspended particle.
• Higher temperature → faster molecules → more energetic collisions → more vigorous Brownian motion.

Quick‑look (core quantitative tools)

Summary checklist

• All matter consists of microscopic particles that are always moving.
• Temperature ↔ average kinetic energy ↔ speed of particles.
• Absolute zero (‑273 °C or 0 K) is the theoretical limit where kinetic energy is minimal.
• Pressure results from particles striking the walls of a container; increasing temperature raises both the frequency and energy of those impacts.
• Collisions can transfer kinetic energy, setting previously stationary particles into motion (e.g., Brownian motion).
• Terminology:

  • Atoms – single‑element units.
  • Molecules – groups of atoms bonded together.
  • Microscopic particles – the broad category that includes atoms, molecules, ions, electrons, dust grains, etc.

• Recall \(pV = \text{constant}\) for a fixed mass of gas at constant temperature and use \(T(\text{K}) = \theta(^{\circ}\text{C}) + 273\) for temperature conversion.

Sample examination questions

1. Explain why a dust particle suspended in air moves randomly even though it is much larger than an air molecule.
2. State the difference between an atom and a molecule. Give one example of each.
3. In a sealed container the temperature is increased. Describe qualitatively how this affects the frequency and energy of collisions between gas molecules and a suspended microscopic particle, and the resulting pressure.
4. Using the temperature conversion formula, calculate the factor by which the pressure of a fixed mass of gas increases when it is heated from 20 °C to 70 °C at constant volume.