Ionization energy trends on the periodic table refer to the systematic variation in the energy required to remove a valence electron from an atom. Generally, ionization energy tends to increase across a period from left to right, decrease down a group from top to bottom, and is determined by atomic size and effective nuclear charge. Across a period, as one moves from left to right, the atomic size decreases due to an increase in the number of protons in the nucleus. This results in a stronger attractive force between the positively charged nucleus and the electrons in the outermost shell. Down a group, from top to bottom, the atomic size increases due to the addition of new energy levels. The outermost electrons are farther from the nucleus, experiencing weaker attractive forces and requiring less energy to be removed. Consequently, elements lower in a group exhibit lower ionization energies. To understand increases, jumps and dips in ionization energy, it's important to have a grasp on how strongly the nucleus pulls on shells as well as an understanding of the order of orbital filling and how this affects stability.
I'M HAVING A TOUGH TIME RECOGNIZING PERIODIC TRENDS
The German term "Aufbau" means "building up," and the Aufbau Principle guides the arrangement of electrons in an atom. According to this principle, electrons fill the lowest energy orbitals first before moving to higher energy levels. Within a given energy level, sublevels or orbitals of different shapes (s, p, d, and f orbitals) are filled in a specific order. Specifically, the s orbital is the first to be occupied, followed by the p, d, and f orbitals in increasing order of energy. This systematic filling of orbitals reflects the progressive "building up" of the electron configuration of an atom, creating a stable arrangement that minimizes the overall energy of the system.
This principle also pertains to the filling of electron orbitals in an atom. However, it focuses specifically on the order of suborbital filling. This rule dictates that electrons will occupy degenerate orbitals (orbitals with the same energy level) in a way that maximizes the total spin of the system, referred to as the "multiplicity" or "spin multiplicity." Specifically, when filling degenerate orbitals, electrons will first fill each orbital with parallel spins before pairing up. This means that electrons in the same sub-level (e.g., the three p orbitals in a given energy level) will each initially have one electron before any of them pair up.
This principle asserts that no two electrons in an atom can have the same set of quantum numbers, which include the principal quantum number, azimuthal quantum number, magnetic quantum number, and spin quantum number. Importantly, it means that electrons within an atom must differ in at least one of these quantum numbers. The most relevant aspect of the Pauli Exclusion Principle is its impact on electron spin. Each electron has a unique spin quantum number, which can have one of two values: +1/2 or -1/2. Therefore, when electrons are filling the orbitals of an atom, they must have opposite spins if they are occupying the same orbital. This exclusion of identical quantum states ensures the distinctiveness and individuality of each electron within an atom.
Sometimes within the pattern of increasing from left to right, there are ionization energy exceptions for some elements on the periodic table. Specifically, there are instances where an unexpected dip or decrease in ionization energy occurs. One notable example is the irregularity in ionization energy between the Group 2 and Group 13 elements. The ionization energy of beryllium is higher than that of boron even though it is further to the left. This discrepancy is attributed to the partially filled p sub-shell in the Group 13 elements, leading to a slightly lower ionization energy compared to the preceding Group 2 elements which have full s orbitals. Another dip occurs between Groups 15 and 16. Moving from left to right, oxygen would be expected to have a higher ionization energy than nitrogen, but it does not. The ionization energy of oxygen is lower than that of nitrogen. According to Hund’s rule, each of the three p sub-orbitals must get one electron before they start pairing. Nitrogen’s electron configuration has exactly one electron in each suborbital making it more stable than oxygen’s configuration which has only one suborbital with paired electrons while the other two are not. This configuration is slightly less stable than nitrogen’s and therefore it is easier to remove one electron.