How 1 mol of gas always takes 22,4 liters of space? It's definitely not coincidence.?

This is derived from the Ideal Gas Law:

PV = nRT

P = pressure (Pa)

V = volume (L)

n = number of moles

R = gas constant = 8.3144598 J/(mol K)

T = temperature (K)

At Standard Temperature and Pressure (STP),

P = 101.325 kPa (1 atm)

T = 0°C or 273.15 K

We substitute the values:

PV = nRT

V = nRT/P = (1 mol)(8.3144598 J/(mol K))(273.15 K) / 101 325 Pa = 0.0224 m³ or 22.4 L

As Preston correctly stated, the molar volume is 22.4 L only at STP. It must be adjusted if temperature or pressure is anything else.

You've left out part of the specification - 1 mol of any gas at 0°C and 1 bar of pressure takes always takes up approximately 22.4 L of volume.

Or stated better: 1 mol of any gas at 0°C confined to 22.4 L of volume always exerts 1 bar of pressure.

If you put that same gas in a smaller volume but only exert 1 bar of pressure on it, its higher pressure will cause it to expand until the pressures equalize, which occurs at 22.4 L.

The relationship is because temperature and pressure are closely related properties. Temperature is the energy of the molecules; pressure is how hard and how often they collide with the walls.

Unlike in a solid or liquid, there are almost no forces between the molecules of a gas. So the energy of the molecules is sufficient to describe how much pressure they exert. It doesn't matter whether that energy is from being heavy or from being fast, and it definitely doesn't matter what shape the molecules are.

When there ARE interactions between the gas molecules, they do not take up exactly 22.4 L. For example, the hydrogen bonds between water molecules compress it to a smaller volume.

Gas expands and is spread far apart so it taking up huge space makes sense

It is called "standard atmospheric pressure" and when you are closer to the ground, the air is thicker, and less liters for every mol of gas.