What is the relationship between electron affinity and periodicity?

In the born haber cycle a gaseous non metal atom will gain one or more electrons to become a negative ion; if in group 7 it will gain 1 electron to become a -1 ion, if in group 6 it will gain 2 electrons to become a -2 ion and if in group 5 it will gain 3 electrons to become a 3- ion). Gaining one or more electrons will release heat energy because the positive nuclear charge is attracted towards the negative electrons and a bond forms so the reaction is exothermic. As you go down a group, the number of protons increases which strengthens the nuclear charge, however, the number of electrons shells increase and this increases the electron shielding so weakens the attraction to the electron so electron affinity weakens down the group.

The relationship between electron affinity and periodicity is governed by trends across the periodic table. Electron affinity, the energy change when an atom gains an electron, generally increases across a period from left to right due to increasing nuclear charge. However, there are exceptions such as the noble gases. Down a group, electron affinity typically decreases due to increased atomic size, which leads to weaker attraction for additional electrons. These trends provide insights into the chemical behavior of elements, aiding in predictive modeling and data analysis.

As you go across a period the number of protons increases which strengthens the nuclear charge. The number of electrons also increase but they are added to the same shell so there is a similar affect in the electron shielding. This means the nuclear attraction increases which increases the affinity for an electron to be gained and added to the outer shell, therefore electron affinity increases across a period. The noble gases have a stable outer shell and do not need to gain any more electrons so group 8/0 is exempt from this description.

As you go down a group in the periodic table, the number of protons in the nucleus increases. However, the number of electrons and electron shells increases which greatly shields the stronger nuclear charge. This weakens the nuclear attraction of the positive nucleus to the negative electrons. Therefore the affinity for an electron weakens down a group.

The number of electrons increases across a period. A S orbital hold 2 electrons (one pair). A P orbital holds 6 electrons (3 pairs). An oxygen atom has a total of 8 electrons; 1s2 2s2 2p4, whereas a sulphur atom holds 16 electrons; 1s2 2s2 2p6 3s2 3p4. An oxygen and sulphur atom both have 6 electrons in their outer electron shell, however sulphur has one more electron shell than oxygen ( a third shell). This means the energy released to add an electron to a sulphur atom is greater than needed to add an electron to an oxygen atom (because the total number of electrons is greater so repels the additional electron being added) and so the electron affinity for sulphur is greater overall.

The ionisation energy generally increases across a period. The anomalies lie between the atoms in groups 2 and 3 and 5 and 6. Example 1-Mg (grp 2) has a higher IE than Al (grp 3) because the outer e in Mg exist in the 3s orbital (1s2 2s2 2p6 3s2) which is lower in energy than the 3p orbtial of which the outer e in Al exist (1s2 2s2 2p6 3s2 3p1). Example 2-P in group 5 has a higher electron affinity then S in group 6. This is because the outer electrons in P exist as an unpaired electron in each of the three 3p orbitals (1s2 2s2 2p6 3s2 3p3), whereas in S the additional electron is paired in the 3p orbital (1s2 2s2 2p6 3s2 3p4). This repulsion lowers the IE required to remove the outer electron and therefore is lower in S than P.