Gases are the only state of matter that can turn a tidy room into a “where did my perfume go?” mystery in seconds. If you have ever watched a balloon swell, a syringe compress, or a scent spread across a classroom, you have seen the same particle-level story playing out. In IB Chemistry, this story matters because exam questions reward you for explaining properties (pressure, compressibility, diffusion) using particles (spacing, motion, intermolecular forces).
The particle-level checklist (what examiners want)
When you compare gases with solids and liquids in IB Chemistry, anchor your answer to three levers:
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Particle spacing (far apart vs close packed)
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Particle motion (free, rapid, random vs restricted)
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Intermolecular forces (negligible vs significant)
If you hit those three clearly, the “why” becomes automatic.
IB Chemistry explanation: spacing is the hidden superpower
In gases, particles are very far apart relative to their own size. That means most of a gas is empty space. At the particle level, that single fact unlocks two big macroscopic behaviours:
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Compressibility: you can push particles closer together because you are mostly reducing empty space.
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Expansion: particles spread out until they occupy the full container volume.
In solids and liquids, particles are already close together, so there is far less “room” to compress. For a quick syllabus-aligned recap, see RevisionDojo’s Structure 1 hub on the particulate model: Models of the particulate nature of matter (IB Chemistry).
Motion: why gases fill containers and create pressure
Gas particles move in constant, random motion, travelling in straight lines between collisions. Those collisions with container walls create pressure.
Solids are different: particles mainly vibrate about fixed positions. Liquids sit in the middle: particles move and flow, but they remain close enough that their motion is still constrained.
If you want the clean markscheme language behind this, revise the assumptions directly: S1.5.1 The ideal gas model notes and the full topic page S1.5 Ideal gases.
Forces: why gases act “independent”
In gases, intermolecular forces are negligible most of the time because particles are so far apart. So their behaviour is dominated by kinetic energy, not attraction.
In liquids and solids, attractive forces matter much more, which is why those states have:
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higher density
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fixed volume (solids and liquids)
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much lower compressibility
To sharpen this comparison, RevisionDojo breaks the idea down clearly here: Why intermolecular forces determine physical states of matter.
Temperature changes hit gases harder
In IB Chemistry, temperature is basically a “speed control knob” for particles. Increase temperature and average kinetic energy increases, so gas particles move faster and collide more often and more forcefully.
That is why heating a gas can cause volume to increase (if pressure is constant) or pressure to increase (if volume is constant). For an exam-focused walkthrough, use: How temperature affects gas volume (IB Chemistry).
Bring it back to RevisionDojo (and your next mark)
If you can explain gases using spacing, motion, and forces, you can answer a surprising number of IB Chemistry questions with confidence. Build that confidence the fast way by pairing understanding with practice: use RevisionDojo’s IB Chemistry resources hub, drill exam-style prompts in the Questionbank, lock in definitions with Flashcards, and sanity-check your explanations with AI Chat. When you are ready to level up, the Study Notes, Mock Exams, Predicted Papers, Grading tools, Coursework Library, and Tutors help you turn particle stories into exam marks.