Why do magnesium, phosphorus, and zinc exhibit slightly higher first ionization energies than the general trend within each of their periods?

The 1st ionization energy, (1st ie) the amount of energy needed to remove the most energetic electron, increases going bottom to top (up the family) and left to right (across the period) on the periodic table in almost all cases.  The actual amount of energy needed to remove that electron is contingent upon four factors: Proximity to the nucleus (proximity), number of protons or the amount of positive charge on the nucleus (charge),  number of electrons between the  electron being removed and nucleus (screening), and spin state, or the amount of negative charge in a given orbital (repulsion). These first 2 factors account for almost the whole ie pattern; the aberrations usually indicate not something unique with a given element, but a change in one of these factors altering the ie of another element around it.  So Magnesium, through the screening effect, makes Aluminum have a lower ie than expected. Sulfur, through repulsion, gives it a slightly lower ie than Phosphorus, where no repulsion effects occur. However, Zinc, relative to Copper, holds all its electrons a bit tighter with its additional proton, so the charge effect is prominent; this occurs because in zinc's configuration of electrons, there's no screening nor repulsive effects; the electrons get that much closer to the nucleus, and the ie goes up.