In general, an element cannot have electrons in an outer shell while having empty orbitals in the inner shells. This is due to the way electrons fill the available energy levels in an atom, following the Aufbau principle. According to this principle, electrons fill the lowest available energy orbitals first before occupying higher energy levels. Therefore, inner shells (which have lower energy) are filled before the electrons start filling the outer shells. There are a few exceptions, particularly in transition metals and heavier elements, where electron configurations may not strictly follow the Aufbau principle due to very small differences in energy levels. However, in these cases, it's not that the inner orbitals are left empty, but rather that electrons are distributed in a way that may seem counterintuitive, such as in the case of chromium and copper.

Elements in the same group of the periodic table have similar chemical properties because they have the same number of valence electrons. The valence electrons are primarily responsible for an element's chemical behaviour, as they are the electrons involved in bonding and reactions. For instance, all alkali metals in Group 1 have one valence electron, which they tend to lose easily, leading to similar reactivity patterns like forming ionic bonds with halogens and reacting vigorously with water. Similarly, noble gases in Group 18 have a full valence shell, making them generally unreactive. This consistency in valence electron configuration across a group leads to similarities in chemical properties, despite differences in atomic size, electronegativity, and other atomic characteristics.

Ions form as atoms seek to achieve a more stable electron configuration, often resembling that of the nearest noble gas. Atoms can achieve this stability by either losing or gaining electrons to fill or empty their outermost shell. The electron configuration of an atom directly influences the charge of the ion it forms. Atoms that have only a few electrons in their outermost shell (like alkali metals) tend to lose these electrons to achieve stability, forming positively charged ions (cations). Conversely, atoms with nearly full outer shells (like halogens) tend to gain electrons to complete their outer shell, forming negatively charged ions (anions). The number of electrons lost or gained determines the charge of the ion. For example, sodium (Na), which has one valence electron, loses this electron to form Na⁺, while chlorine (Cl), which needs one electron to complete its outer shell, gains an electron to form Cl⁻. This process is governed by the tendency of elements to achieve a stable, noble gas-like electron configuration.

Across a period, as the atomic number increases, electrons are added to the same principal energy level while the number of protons in the nucleus also increases. This increased nuclear charge attracts the electrons more strongly, pulling them closer to the nucleus. As a result, despite the increase in the number of electrons, the atomic radius decreases across a period. This is because the increased nuclear charge outweighs the addition of electrons in the same shell. In contrast, down a group, new electron shells are added, placing the outermost electrons further from the nucleus. Even though the nuclear charge increases down a group, the effect of increased distance and the added electron shielding from inner shells means that the atomic radius generally increases down a group. Thus, electron configuration plays a critical role in determining the size of atoms, both across periods and down groups in the periodic table.

The filling of the 4s orbital before the 3d orbital is a consequence of the subtle differences in energy levels within the atom. In terms of principal quantum numbers, the 4s orbital belongs to the fourth shell, while the 3d orbital belongs to the third shell. It might seem logical for the 3d orbital to have a lower energy level, but due to the complexities in the atomic structure, the 4s orbital is actually lower in energy when it's empty or partially filled. Therefore, electrons fill the 4s orbital first. Once the 4s orbital is filled, the electrons then begin to populate the 3d orbital. This is important because the energy levels of orbitals are not fixed and can change depending on the electron configuration. As more electrons are added, the relative energy levels can shift, causing the 3d orbitals to become lower in energy than the 4s orbitals in larger atoms.