The gases CO₂ (g) and NH₃ (g) can be liquefied at 20°C by compressing them to sufficiently high pressures. A student claims that NH₃ (g) can be liquefied at a lower pressure than CO₂ (g) can be liquefied. What is the best justification for this claim?

NH₃ has a lower molecular mass and a stronger intermolecular bond.

Explanation:

The physical change depends on the temperature and the pressure of the substance. In liquefaction, the substance will go from gas to liquid, so its molecules must have less energy, and be more close.

The energy depends on the temperature, has higher the temperature, higher the internal energy. And the distance of the molecules depends on the strength of the bonds, and also the pressure. As higher the pressure, more close will be them, and as stronger the bond between them, as close they will be.

When a compound has a great molar mass, more energy is necessary to it to gain or to lose, to change its physical state. NH₃ has a molar mass equal to 17 g/mol, and CO₂ has a molar mass equal to 44 g/mol. So, CO₂ needs more energy to change its physical state.

Because temperature is constant, it will be necessary more pressure at CO₂ to give them the necessary energy to liquefy. Besides, NH₃ has hydrogen bonds between molecules, and CO₂ has dipole-dipole bonds.

Hydrogen bonds are formed when hydrogen bond with N, O, or F, and dipole-dipole bonds are formed in polar compounds, which are formed by atoms with large difference of electronegativity, such as C and O. Hydrogen bonds are stronger than dipole-dipole bonds because the difference in electronegativity is large. So, the molecules of NH₃ intend to be together, then it's easier to liquefy it.