Ideal Gas Law is an Approximation
The ideal gas law describes how gases behave, but does not account for molecular size or intermolecular forces. Since molecules and atoms in all real gases have size and exert force on each other, the ideal gas law is only an approximation, albeit a very good one for many real gases. It is most accurate for monoatomic gases at high pressure and temperature, since it is for these gases that size and intermolecular forces play the most negligible role.
Strength of Intermolecular Forces
Depending on their structure, size and other properties, different compounds have different intermolecular forces--that's why water boils at a higher temperature than ethanol, for example. Unlike the other three gases, ammonia is a polar molecule and can hydrogen-bond, so it will experience stronger intermolecular attraction than the others. The other three are subject only to London dispersion forces. London dispersion forces are created by transient, short-lived redistribution of electrons that makes a molecule act as a weak temporary dipole. The molecule is then able to induce polarity in another molecule, thereby creating an attraction between the two molecules.
In general, London dispersion forces are stronger between larger molecules and weaker between smaller molecules. Helium is the only monoatomic gas in this group and hence the smallest in terms of size and diameter of the four. Since the ideal gas law is a better approximation for monoatomic gases--and since helium is subject to weaker intermolecular attractions than the others--out of these four gases, helium is the one that will behave most like an ideal gas.