ideal gas
Picture a room full of tiny, hard balls flying about at enormous speeds, bouncing off the walls and off each other but otherwise paying no attention to one another. They are so small compared with the empty space between them that, most of the time, each ball travels as if it were utterly alone. That cartoon is an ideal gas — a deliberately simplified stand-in for a real gas.
Formally, an ideal gas is a hypothetical gas whose molecules occupy no volume themselves and exert no forces on one another except during brief, perfectly elastic collisions. Under these assumptions its pressure, volume, temperature and amount are tied together by a single, beautifully simple equation, and many of its properties can be calculated exactly.
No real gas is truly ideal, but the picture is astonishingly good when a gas is hot and dilute — that is, when the molecules are far apart and moving fast, so the time they spend close enough to feel each other is negligible. The ideal gas is the starting point of nearly all thermodynamics, and the corrections needed to describe real gases are small departures from it.
The air in a bicycle tyre at everyday temperatures and pressures behaves almost exactly like an ideal gas — accurate enough that pumps, weather charts and chemistry classes all treat it that way.
Common air at room conditions is a near-perfect ideal gas.
Ideal does not mean perfect or best — it means simplified. The model breaks down when a gas is cold or highly compressed, conditions under which molecules crowd together and their mutual attractions can no longer be ignored.