Lesson Overview

Gases differ from liquids and solids in three fundamental ways: they lack both definite shape and definite volume, they are highly compressible due to the vast empty space between particles, and they mix freely with other gases in any proportion. Pressure, defined as force per unit area, is the central variable for describing the state of a gas. Standard pressure unit conversions — including atmospheres, millimeters of mercury, torr, and pascals — provide the foundation for solving gas law problems in general chemistry.

Table of Contents

• 3 Properties That Distinguish Gases from Solids and Liquids
• Definite Shape and Volume
• Compressibility
• Mixing with Other Gases
• Pressure
• 6 Units of Pressure
• Pressure Unit Conversions
• Key Takeaways

3 Properties That Distinguish Gases from Solids and Liquids

Before diving into the behavior of gases, it is worth establishing exactly what sets them apart from liquids and solids. Three key properties draw the clearest distinctions: definite shape and volume, compressibility, and the ability to mix freely with other gases.

Definite Shape and Volume

A gas has neither a definite shape nor a definite volume. Place a gas inside any container and it will expand to fill that container completely, adopting both its shape and its volume.

A liquid behaves differently: it takes on the shape of its container, but it does not expand to fill available space. A liquid therefore has a definite volume but no definite shape. A solid goes further still, maintaining both its own shape and its own volume regardless of the container.

Compressibility

Gases are composed of molecules or atoms separated by vast regions of empty space. This stands in sharp contrast to liquids and solids, where particles are in close, direct contact with one another. Because gases are mostly empty space, they can spread out to fill a container — and, crucially, they can also be compressed: applying pressure forces the molecules closer together, reducing the volume.

Liquids and solids, by contrast, are essentially incompressible. You can certainly apply enormous pressure to a liquid or a solid, but the resulting volume change will be negligibly small. With a gas, the relationship is far more dramatic: under many conditions, pressure and volume are inversely proportional — double the pressure and the volume is cut in half. (This relationship, known as Boyle's Law, will be explored in detail in subsequent lessons.)

Mixing with Other Gases

Gases will mix with other gases in any proportion. Because gas-phase particles are separated by large distances, it makes no difference whether they are attracted to or repelled by one another — they simply coexist in the same space, each occupying its share of the empty volume.

This is not always true for liquids. Oil and water, for instance, refuse to mix: the oil molecules and water molecules have very different intermolecular forces, causing them to segregate into two distinct layers. (The full explanation, involving intermolecular forces and entropy, will come in the next chapter.)

Pressure: Definition and Formula

What Is Pressure?

Because gases are so compressible, pressure is one of the most important variables used to describe their state — far more so than for liquids or solids. While pressure can be defined for any material, the extreme incompressibility of liquids and solids makes pressure less central to their description in introductory chemistry.

Mathematically, pressure is defined as force per unit area:

Pressure Is Directly Proportional to Force, Inversely Proportional to Area

Consider this scenario: you want to pierce the skin of a watermelon. You have two options — strike it with the blunt end of a marker, or strike it with a knife. In both cases, suppose you apply exactly the same force. Which will pierce the skin?

The knife wins — not because more force is applied, but because that force is concentrated over a far smaller surface area at the blade's edge. A smaller area means higher pressure at the point of contact, and it is that elevated pressure that drives the blade through the skin.

6 Units of Pressure in Chemistry

Several units of pressure appear throughout chemistry. You should be comfortable with all of them and know how they relate to one another.

Atmosphere (atm)

The atmosphere (atm) is the most common unit used in general chemistry. By definition, 1 atm is the average air pressure at sea level on Earth — it is literally the weight of the entire column of atmosphere pressing down on each square meter of Earth's surface at that elevation. As altitude increases, less atmosphere sits overhead and pressure drops below 1 atm.

Millimeters of Mercury (mmHg)

Pressure is also expressed in millimeters of mercury (mmHg), a unit rooted in early laboratory practice. Traditional barometers and manometers used liquid mercury: at standard atmospheric pressure, the atmosphere can push down on a mercury reservoir and force a column of mercury to rise to a height of exactly 760 mm. That measurement became the unit itself.

By definition: 1 atm = 760 mmHg.

Torr

The torr is a closely related unit named in honor of Evangelista Torricelli, inventor of the barometer. At Earth's surface under standard conditions, 1 torr = 1 mmHg, and the two units are used interchangeably in introductory chemistry.

By definition: 1 atm = 760 torr.

Pascal (Pa) and Kilopascal (kPa)

The pascal is the SI unit of pressure. One atmosphere equals exactly 101,325 Pa — a value conveniently approximated as 10⁵ Pa, or 100 kPa, for estimation purposes.

Pressure Unit Conversions

Pressure conversions follow the same dimensional analysis approach used throughout chemistry. Set up a conversion factor so that the unit you want to eliminate cancels, and the unit you want to keep remains.

Converting Atmospheres to Torr

Converting Atmospheres to Pascals