The States of Matter

Matter comes in three everyday forms: solid, liquid, and gas. The differences between them are not random. They come down to how tightly particles are packed. They also depend on how freely those particles move past one another. Once you can picture the particle-level arrangement inside each state, you can make predictions. You can predict whether a substance holds its shape, flows, or fills any container it's put in. This article builds directly on an earlier idea. Matter is made of distinguishable substances, covered in Classification of Matter: Pure Substances and Mixtures. Here the classification is by physical state rather than by composition. The same substance — water, for instance — can exist as ice, liquid water, or steam. This depends on conditions. Understanding states of matter sets up everything that follows. This includes how matter behaves and changes, covered in Physical and Chemical Properties and Changes. The goal here is purely descriptive and qualitative. There are no equations of state and no kinetic theory formulas. Instead, this offers a clear, quantitatively-grounded picture of particle spacing and motion. It also covers the macroscopic consequences. These include fixed shape versus fixed volume, compressibility, and density.

Key takeaways

Solids have fixed shape and fixed volume; particles vibrate in place but do not move past each other.
Liquids have fixed volume but no fixed shape; particles stay close but slide past one another.
Gases have neither fixed shape nor fixed volume; particles are far apart and move independently.
Density typically drops sharply from solid/liquid to gas because gas particles occupy roughly 1000× more space per particle.
Heating generally increases particle motion, driving transitions solid → liquid → gas; the reverse cooling drives gas → liquid → solid.
Key idea: the three common states differ in particle spacing and freedom of motion, and these two features alone explain shape, volume, compressibility, and density behavior.

The core insight

A state of matter is defined by the balance between two things: how strongly particles attract each other and how much kinetic energy they have from thermal motion. In solids, attraction wins decisively; in gases, kinetic energy wins decisively; liquids are the compromise in between.

The physical picture

Every pure substance can, in principle, exist as a solid, a liquid, or a gas, depending on temperature and pressure. The particles themselves — atoms, ions, or molecules — do not change identity between states; a water molecule in ice is chemically identical to a water molecule in steam. What changes is the arrangement and motion of those particles. In a solid, particles are packed closely in fixed positions, usually a regular repeating pattern (as in a crystal), though some solids are amorphous with less regular ordering. Each particle vibrates about a fixed point but does not migrate through the material. Because the positions are locked, a solid has a definite shape and a definite volume: it resists both compression and deformation. Typical particle-to-particle spacing in a solid is close to the particle's own diameter, often around 0.3–0.5 nm for small molecules, with essentially no empty space between neighbors. In a liquid, particles are still close together — the spacing is only slightly larger than in the solid, typically a few percent more — but the particles have enough kinetic energy to slide past one another. This gives a liquid a definite volume (particles remain in contact, so the liquid resists compression) but no definite shape (it flows to take the shape of its container). Liquids are only slightly less dense than their corresponding solids, and in the unusual case of water, the liquid is actually denser than the solid because ice's crystal structure is unusually open. In a gas, particles are separated by distances roughly ten times their own diameter or more, meaning the particles occupy only about 1/1000th of the volume that the gas as a whole occupies at ordinary pressure — the rest is empty space. Particles move independently at high speed, colliding with each other and the container walls, with no significant attractive forces holding them in fixed relative positions. A gas therefore has neither definite shape nor definite volume: it expands to fill any container completely and is easily compressed. The following two figures summarize the arrangement and the resulting macroscopic properties.

Figure

Particle arrangement in the three states of matter

Figure

Macroscopic consequence: shape and volume behavior

The core relations

There is no single dynamical equation governing "state of matter" at this level — the topic is qualitative — but two quantitative relations organize the reasoning cleanly: the density relation used to compare states, and the ideal-gas-limit ratio that shows why gases occupy so much more space per particle.

\rho = \frac{m}{V}

where:

$\rho$ is density (commonly g/mL for liquids/solids, g/L for gases)
$m$ is the mass of the sample
$V$ is the volume the sample occupies

\frac{V_{\text{gas}}}{V_{\text{liquid}}} \approx \frac{\rho_{\text{liquid}}}{\rho_{\text{gas}}}

where:

$V_{\text{gas}}$ , $V_{\text{liquid}}$ are the volumes occupied by the same mass of substance as gas and as liquid
$\rho_{\text{liquid}}$ , $\rho_{\text{gas}}$ are the corresponding densities
this ratio quantifies "gas particles are far apart" in a way you can check with real numbers

Derivation AWhy a gas occupies ~1000× more volume than the same mass of liquid

Goal: show, using real density values for water, that converting liquid to gas expands the volume roughly a thousandfold — the quantitative signature of "particles far apart" in the gas picture.

Step 1. Start from the definition of density for a fixed mass $m$ in each state:

V_{\text{liquid}} = \frac{m}{\rho_{\text{liquid}}}, \qquad V_{\text{gas}} = \frac{m}{\rho_{\text{gas}}}

Step 2. Take the ratio of the two volumes; the mass $m$ cancels because it is the same sample of matter in both states:

\frac{V_{\text{gas}}}{V_{\text{liquid}}} = \frac{m/\rho_{\text{gas}}}{m/\rho_{\text{liquid}}} = \frac{\rho_{\text{liquid}}}{\rho_{\text{gas}}}

Step 3. Insert real values. Liquid water at 100 °C has density $\rho_{\text{liquid}} = 0.958\ \text{g/mL}$ . Steam at 100 °C and 1 atm has density $\rho_{\text{gas}} = 0.000598\ \text{g/mL}$ (equivalently 0.598 g/L).

\frac{V_{\text{gas}}}{V_{\text{liquid}}} = \frac{0.958}{0.000598} \approx 1602

Step 4. Interpret the number. One gram of liquid water at its boiling point occupies about 1.04 mL; the same gram as steam occupies about 1670 mL. The expansion factor of roughly 1600 confirms, quantitatively, that gas particles are spread over a volume more than a thousand times larger than in the liquid — consistent with the qualitative picture of gas particles separated by about ten diameters instead of touching.

Limiting case check. If a substance's gas and liquid densities were ever to become equal ( $\rho_{\text{gas}} \to \rho_{\text{liquid}}$ ), the ratio would go to 1, meaning no volume change on vaporization — this only happens infinitesimally close to the critical point, where the distinction between liquid and gas disappears. Far from the critical point, as is true for water at 1 atm, the ratio stays large, which is exactly why gases are so easily compressed and liquids are not.

Worked examples

Example 1 — Comparing densities of copper in solid and (hypothetical) gas form. Solid copper has density 8.96 g/cm³. Suppose 3.50 g of copper vapor at some condition occupies 415 cm³ (this is illustrative — copper's boiling point is far above ordinary conditions, but the arithmetic is the same idea). Step 1: Find the volume of 3.50 g of solid copper using $\rho = m/V$ :

V_{\text{solid}} = \frac{m}{\rho_{\text{solid}}} = \frac{3.50\ \text{g}}{8.96\ \text{g/cm}^3} = 0.3906\ \text{cm}^3

Step 2: Find the density of the vapor sample:

\rho_{\text{vapor}} = \frac{m}{V_{\text{vapor}}} = \frac{3.50\ \text{g}}{415\ \text{cm}^3} = 0.008434\ \text{g/cm}^3

Step 3: Compute the expansion ratio going from solid to vapor:

\frac{V_{\text{vapor}}}{V_{\text{solid}}} = \frac{415}{0.3906} \approx 1063

The vapor occupies roughly 1060 times the volume of the same mass of solid — the same order of magnitude found for water above, confirming this is a general feature of the gas state rather than a coincidence specific to water.

Example 2 — Identifying state from particle-spacing data. A substance's particles are found (by X-ray-type spacing measurements) to sit an average of 3.8 Å apart in one sample and 41 Å apart in another sample of the same substance at the same temperature, where 1 Å = $10^{-10}$ m and the particle diameter is about 3.6 Å. Step 1: Compare the 3.8 Å spacing to the particle diameter of 3.6 Å:

\frac{3.8\ \text{Å}}{3.6\ \text{Å}} = 1.06

Particles are touching almost exactly at their own diameter apart — this is characteristic of a condensed phase (solid or liquid), since there is essentially no empty space between neighbors. Step 2: Compare the 41 Å spacing to the same diameter:

\frac{41\ \text{Å}}{3.6\ \text{Å}} = 11.4

Particles are separated by over ten diameters, consistent with the gas state, where most of the volume is empty space between widely-spaced, independently-moving particles. Step 3: Note that spacing alone (1.06 vs 11.4) cannot distinguish solid from liquid — that requires checking whether the particles are fixed in a lattice (solid) or free to slide (liquid), a distinction not visible from average spacing alone. This limiting comparison shows spacing ratio is a good gas-vs-condensed test but not a solid-vs-liquid test.

Limiting cases & common traps

Situation	Correct reasoning	Common trap
Gas in a rigid sealed container	Gas fills entire container volume, however large	Thinking a gas "settles" like a liquid and only fills the bottom
Ice vs. liquid water density	Ice is less dense than liquid water (0.917 vs 0.998 g/mL) — an exception	Assuming solids are always denser than their liquids (usually true, but not for water)
Compressing a liquid	Liquids resist compression strongly; volume changes are tiny even under high pressure	Treating liquids as "almost as compressible as gases"
Very low-pressure gas	Volume expands essentially without limit as pressure drops; expansion ratio versus liquid grows	Assuming the ~1000× ratio is a fixed universal constant rather than pressure/temperature dependent
Amorphous solids (e.g., glass)	Particles are closely packed but not in a perfectly repeating lattice; still has fixed shape and volume	Assuming "solid" always implies a perfect crystal structure

The chemist behind it

The modern particle picture of states of matter grew out of the kinetic theory of gases, developed through the 19th century by scientists connecting macroscopic gas behavior (pressure, volume, temperature) to microscopic particle motion. One of the central figures in building this bridge — showing that heat and particle motion were the same underlying phenomenon, and that a distribution of particle speeds rather than a single speed governs bulk properties — was James Clerk Maxwell, whose statistical treatment of gas particle velocities gave the kinetic-molecular picture its rigorous mathematical foundation and made it possible to derive, not just assume, why gases expand to fill their containers and why increasing temperature increases pressure at constant volume.

Origins · 1 minJames Clerk Maxwell and how this was discovered

James Clerk Maxwell

1831–1879 · Developed the statistical distribution of molecular speeds in gases, providing the mathematical basis for the kinetic-molecular theory that explains why gas particles move independently and spread to fill any container, directly underpinning the modern particle-level description of the three states of matter.

Ask the tutor: If I compress a gas until its particles are as close together as in a liquid, does it automatically become a liquid?

Why does a gas take the shape of its container but a liquid doesn't fully?

A liquid's particles remain close enough together, and attract each other strongly enough, that the bulk liquid holds together as a connected mass — it flows to fill the bottom of a container (conforming to shape) but the top surface stays flat under gravity rather than expanding to fill all available space. Gas particles have negligible mutual attraction on the scale of their spacing and high kinetic energy, so they spread out until they hit a boundary, filling the entire container, top to bottom.

Is plasma a fourth state of matter covered here?

Plasma — ionized gas, found in lightning, neon lights, and stars — is often called a fourth state of matter, but it is outside the scope of an introductory treatment focused on solids, liquids, and gases. The particle picture is similar to a gas (particles far apart, independent motion) with the added feature that some particles are charged ions and free electrons rather than neutral atoms or molecules.

Why is ice less dense than liquid water when solids are usually denser than their liquids?

Water molecules in ice form a hydrogen-bonded hexagonal lattice with relatively large open channels between molecules, making the solid more open (less densely packed) than the liquid, where molecules can pack more efficiently as some hydrogen bonds break and molecules crowd closer on average. This is why ice floats — an unusual but chemically well-understood exception to the general "solid denser than liquid" pattern.

Can the same amount of substance be a solid, liquid, and gas at the same temperature?

Not typically at a single temperature and pressure, except exactly at special points. At the triple point, a substance's solid, liquid, and gas phases can coexist simultaneously in equilibrium at one specific temperature and pressure combination unique to that substance. Away from the triple point, only one or two phases can coexist stably at a given temperature/pressure.

How does temperature relate to which state a substance is in?

Raising temperature increases the average kinetic energy of particles, which tends to overcome the attractive forces holding particles in fixed positions or close contact, driving transitions solid → liquid → gas. The exact temperatures at which these transitions occur (melting point, boiling point) depend on the strength of the particle-particle attractions and the pressure, topics developed further when studying intermolecular forces and phase changes.