Final answer:
Electron affinity generally decreases as you move down a group due to increasing atomic size and decreases from right to left across a period due to a weaker nuclear charge, though there are exceptions related to electron subshell configurations.
Step-by-step explanation:
Electron affinity generally decreases down groups and from right to left across periods on the periodic table because of the increased atomic size and the lesser ability of the nucleus to attract additional electrons from distant orbitals. As we go down a group, the increase in principal quantum number (n) means electrons are being added to orbitals that are further from the nucleus, which decreases nuclear attraction and makes it less energetically favorable for atoms to gain electrons, thus reducing their electron affinity.
Moving from right to left across a period, the effective nuclear charge experienced by an electron being added to an atom decreases because the atomic number is lower, thus the pull from the nucleus is weaker. This again results in a decrease in electron affinity. However, there are exceptions to these general trends, particularly when considering electrons entering half-filled or completely filled subshells, which can alter the patterns due to increased electron-electron repulsion or increased stability, respectively.