There's something odd about your data. The website I consulted gave $c_v$, the molar heat capacity at constant volume as 12.4717 $\text J\ \text{mol}^{-1}\ \text {K}^{-1}$ for both helium and neon (I couldn't find Xenon). Both of these were measured at 25 °C, though for inert gases there is very little change of $c_v$ with temperature.

This value of $c_v$ agrees to five sig figs with the theoretical value of $\frac32 R$.

So at equal temperatures the same amount (number of moles) of the gases have the same internal energy, given by $$U=\tfrac32 nRT.$$ This works for all monatomic gases at lowish densities (so they behave as ideal gases).

(a) There's something odd about your data. The website I consulted gave $c_v$, the molar heat capacity at constant volume as 12.4717 $\text J\ \text{mol}^{-1}\ \text {K}^{-1}$ for both helium and neon (I couldn't find Xenon). Both of these were measured at 25 °C, though for inert gases there is very little change of $c_v$ with temperature.

This value of $c_v$ agrees to five sig figs with the theoretical value of $\frac32 R$.

So at equal temperatures the same amount (number of moles) of the gases have the same internal energy, given by $$U=\tfrac32 nRT.$$ This works for all monatomic gases at lowish densities (so they behave as ideal gases).

(b) However for diatomic gases, such as oxygen, nitrogen, hydrogen, a good approximation for the molar heat capacity at constant volume is $c_v = \frac52 RT$, so $$U=\tfrac52 nRT.$$

The reason is that the molecules of these gases have kinetic energy of rotation as well as translation (moving about!). We say that diatomic molecules at ordinary temperatures have 5 "degrees of freedom", 3 translational and 2 rotational. Kelvin temperature is proportional to the average kinetic energy per degree of freedom, so a mole of diatomic molecules has 5/3 times the internal energy of a mole of monatomic molecules at the same temperature!