Write the electron configuration for each of the following ions: (a) \(\mathrm{As}^{3-}\) (b) I\(^{-}\) (c) \(\mathrm{Be}^{2+}\) (d) \(\mathrm{Cd}^{2+}\) (e) \(\mathrm{O}^{2-}\) (f) \(\mathrm{Ga}^{3+}\) (g) \(\mathrm{Li}^{+}\) (h) \(\mathrm{N}^{3-}\) (i) \(\mathrm{Sn}^{2+}\) (j) \(\mathrm{Co}^{2+}\) (k) \(\mathrm{Fe}^{2+}\) (l) \(\mathrm{As}^{3+}\)

The Aufbau principle is a crucial guideline for understanding the order in which electrons are arranged in an atom's orbitals. The word 'Aufbau' is German for 'building up' or 'construction.' According to this principle, electrons fill the atomic orbitals starting from the lowest energy level and moving towards higher energy levels. This means that before a higher energy orbital can be filled, all lower energy orbitals must be occupied by electrons.

When applying the Aufbau principle to ions, adjustments must be made based on whether electrons are being added to an anion (a negatively charged ion) or removed from a cation (a positively charged ion). The electron configuration for an ion reflects these changes. For instance, when an electron is removed from an atom to form a cation, the last electron added (according to the Aufbau principle) is the first to be removed.

Hund's rule deals with the distribution of electrons across orbitals of the same energy—known as 'degenerate orbitals.' It states that electrons will fill an empty orbital before they pair up in an already occupied one. This is because electrons are negatively charged and repel each other; they will occupy separate orbitals when possible to minimize repulsion.

When looking at ion electron configurations, this rule still applies. For each sublevel being filled or emptied, Hund’s rule guides us to add electrons to empty orbitals first before pairing them, or conversely, to remove electrons from singly occupied orbitals before those that are paired.

The Pauli exclusion principle is another fundamental concept in quantum mechanics. It states that no two electrons in an atom can have the same set of four quantum numbers. In simpler terms, it means that each orbital can hold a maximum of two electrons, and they must have opposite spins.

This principle influences the writing of the electron configuration for ions by dictating that when an electron is added to or removed from an orbital, it respects the maximum capacity and the need for opposite spins. The principle ensures that the right number of electrons is allocated and that they are properly distributed according to their spin.

The atomic number is a fundamental characteristic of an element. It represents the number of protons found in the nucleus of an atom and, in a neutral atom, it also equals the number of electrons surrounding the nucleus. This number is crucial because it defines the element and determines the order of the periodic table.

In terms of electron configurations for ions, the atomic number tells us the starting point—the configuration of the neutral atom—before considering the gain or loss of electrons due to ionization. After we determine the atomic number, we then adjust the electron count to reflect the ion's charge to get the correct electron configuration for the ion.

Orbital distribution refers to how electrons are allocated among the different orbitals surrounding an atom's nucleus. Each orbital type (s, p, d, f) has a distinctive shape and a specific maximum capacity of electrons it can hold. Understanding this arrangement is vital for writing the electron configuration of both neutral atoms and ions.

For ions, the process begins just as it does for neutral atoms: by distributing electrons in order of increasing energy levels. However, when ions are formed, their orbital distribution changes—in cations, electrons are removed usually starting with the highest energy level orbital, and in anions, electrons are added starting with the lowest available energy level orbital not fully occupied. This influences the chemical properties and reactivity of the ions formed.