Why does a gas fill up an entire balloon evenly, while a solid just sits there holding its shape? Why can you squeeze a syringe full of air but not one full of water? To answer questions like these, chemists use a mental model called the kinetic molecular theory (KMT) — a set of assumptions about how tiny particles (atoms or molecules) behave that, together, explain almost everything gases do. You don't need to see individual molecules to use this model; you just need to accept its postulates and follow where they lead.
The three states, particle-view
- Particles vibrate in fixed positions
- Strong attractive forces hold particles close together
- Fixed shape and fixed volume
- Very low compressibility
- Particles move freely in straight lines between collisions
- Attractive forces are essentially negligible
- No fixed shape or volume — takes the shape and volume of its container
- Highly compressible; low density
Liquids sit in between: particles are close together like a solid (so liquids barely compress) but are free to slide past one another (so liquids flow and take the shape of their container, but keep a fixed volume). This unit focuses on gases, the state where the kinetic molecular theory has the most explanatory power.
| Feature | Solid | Liquid | Gas |
|---|---|---|---|
| Particle spacing | Touching, in a fixed pattern | Touching, but disordered | Very far apart (mostly empty space) |
| Particle motion | Vibrate in place | Slide past one another | Fast, straight-line, random |
| Attractive forces | Strong enough to lock positions | Strong enough to keep particles together | Negligible (ideal model) |
| Shape / volume | Fixed / fixed | Takes container shape / fixed | Takes container shape / fills container |
| Compressibility | Almost none | Very slight | High |
| Density of water in that state | Ice: about 0.92 g/mL | Liquid: about 1.00 g/mL | Steam at 100°C: about 0.0006 g/mL |
The postulates of the kinetic molecular theory (for an ideal gas)
- Gases consist of a huge number of tiny particles (atoms or molecules) in constant, random, straight-line motion, changing direction only when they collide.
- The volume of the particles themselves is negligible compared to the volume of empty space between them — a gas is mostly empty space.
- There are no attractive or repulsive forces between particles (or between particles and the container walls) — particles act completely independently except during a collision.
- Collisions between particles, and between particles and the container walls, are perfectly elastic — no kinetic energy is lost overall, it's just transferred between particles.
- The average kinetic energy of the particles is directly proportional to the absolute (Kelvin) temperature of the gas — every gas at the same temperature has the same average kinetic energy, regardless of identity.
These postulates were not pulled out of thin air; each one is backed by something you can observe. Gases compress easily, so there must be a lot of empty space between particles (postulate 2). A gas never settles into a layer at the bottom of a sealed jar, so its particles must keep moving (postulate 1). A sealed tire does not slowly lose pressure just from sitting still, so collisions cannot be draining away energy (postulate 4). And when chemists heated many different gases, they all expanded by the same fraction for each degree, which only makes sense if temperature controls particle energy the same way for every gas (postulate 5). A good model is a short list of assumptions that explains a long list of observations.
Temperature is average kinetic energy
At any moment, the particles in a gas do not all move at the same speed. A few are nearly stopped after an unlucky collision, a few are moving unusually fast, and most are somewhere in between. Temperature measures the average kinetic energy of this whole crowd. Kinetic energy depends on both mass and speed (), so if two gases are at the same temperature (same average ), the lighter particles must be moving faster. A helium atom has about one-seventh the mass of a nitrogen molecule, so at room temperature it moves roughly 2.6 times faster on average. Nitrogen molecules in the room around you already average close to 500 m/s, yet a smell released across the room takes seconds to reach you: each particle collides billions of times per second, so its path is a zig-zag random walk rather than a straight dash. Because average kinetic energy is proportional to Kelvin temperature, doubling the Kelvin temperature doubles the average kinetic energy: with both temperatures in kelvins.
Explaining gas properties from KMT
| Observed property | KMT explanation |
|---|---|
| Gases are easily compressed | Particles are far apart with mostly empty space between them, so pushing them closer together is easy — there isn't much "stuff" in the way. |
| Gases expand to fill any container | Particles move independently in straight lines with no attractive forces holding them together, so they spread out until they hit walls on every side. |
| Gases have low density compared to solids/liquids | The same number of particles occupies a much larger volume, since most of that volume is empty space. |
| Gases diffuse (mix) into each other | Random particle motion carries particles from a region of high concentration into a region of low concentration over time. |
| Gases exert pressure on their container | Particles are constantly colliding with the container walls; each tiny collision delivers a tiny push, and billions of collisions per second add up to measurable pressure. |
That last row is the key idea that the next topic builds on: gas pressure is not some separate mysterious property — it is a direct, measurable consequence of particles slamming into the walls of whatever is holding them. When a test question says "use the kinetic molecular theory to explain", a full answer has three links: name what changed (for example, the temperature rose), describe what the particles do differently (they move faster, so they hit the walls harder and more often), and connect that particle behaviour to the observation (so the pressure rises). Skipping the middle link is the most common way to lose marks on these questions.