Ionic radii follow similar trends to atomic radii with one critical difference. Cations have very different radii than anions. Cations are all smaller than their neutral analogs while anions are all larger. This is easy to understand since cations have lost electrons. As a result they have both fewer electrons in the highest energy atomic orbitals farthest from the nuclei and the remaining electrons feel a stronger pull from the nucleus. Look at Na+ which is isoelectronic with Ne. It goes from being one of the "largest" atoms on the left-hand side of the periodic table to effectively one of the smallest with an electron configuration that is the same a neon (all the way on the right hand side). Moreover, Ne has a nuclear charge of Z=10 and Na+ has a Z=11. Thus the Na+ should be smaller than Ne. However as atomic radii and ionic radii are often defined differently then this comparison is difficult to make.
Conversely adding an electron to F to make F- also generates an ion that isoelectron with Ne. However, now you have added electrons and kept the number of protons constant. Thus F- will be larger than Ne (and larger than Na+).The general trends will continue to hold. From top to bottom of the periodic table ions will increase in radii. However, now left to right the radius is more of a function of the number of electrons. Mg2+ is smaller than Na+. Both have 10 electrons but Z=12 for Mg and Z=11 for Na. Similarly, O2- will be larger than F- as both have 10 electrons but Z=8 for oxygen and Z=9 for fluorine.
The following chart so both the atomic and ionic radii for some common ions. You can see the most dramatic differences are generally found in the trend top to bottom.
