The elements \(\mathrm{N}, \mathrm{P}, \mathrm{As}, \mathrm{Sb}\), and Bi belong to the same group in the periodic table. Write the electron configuration expected for the ground-state atoms of these elements and predict how many unpaired electrons, if any, each atom has.

P-block elements are located on the right-hand side of the periodic table, from groups 13 to 18. They are characterized by having their valence electrons in p-orbitals, which are the penultimate energy level of atoms.

These elements typically display a diverse range of properties and include metals, metalloids, and non-metals. Interestingly, the p-block encompasses more elements than any other block in the periodic table. As we climb the periodic table vertically within the groups, the elements tend to become heavier, have higher atomic numbers, and exhibit different chemical properties.

In the context of an educational exercise, we might be tasked with determining the electron configurations of p-block elements. It is important to note that the p-orbitals can hold a maximum of six electrons across three sub-orbitals (px, py, and pz). The configuration process usually starts by filling the s-orbital and then adding electrons to the p-orbital. A familiarity with the periodic table allows us to identify the block, period, and hence the principal quantum number to write the full electron configuration succinctly.

Group 15 of the periodic table is comprised of the nitrogen family, which includes nitrogen (N), phosphorus (P), arsenic (As), antimony (Sb), and bismuth (Bi). This group is part of the p-block and is also referred to as the pnictogens.

Each of these elements possesses five valence electrons (hence the original Group VA designation), with the general electron configuration of ns2 np3. This configuration suggests that these elements have filled s-orbitals and half-filled p-orbitals at their highest energy level.

When exploring Group 15 elements, students should pay particular attention to trends in electronegativity, atomic size, oxidation states, and chemical reactivity. These trends can be predicted based on the element's position within the group. For example, nitrogen, being at the top, is known for forming stable triple bonds due to its high electronegativity, whereas bismuth at the bottom displays more metallic characteristics.

The concept of unpaired electrons is fundamental to understanding magnetic properties and chemical reactivity of elements. An unpaired electron is one that occupies an atomic or molecular orbital by itself, rather than as part of an electron pair. Elements with unpaired electrons are called paramagnetic and are attracted to magnetic fields, while those without unpaired electrons are diamagnetic and are not.

According to Hund's rule, orbitals of the same energy are each occupied by a single electron before any orbital is doubly occupied. Hence, for Group 15 elements, the p-orbital hosting three electrons will have each electron in a separate sub-orbital, making them unpaired.

This concept is not only important when predicting magnetic properties but also has implications in chemical bonding and reactions. For instance, the unpaired electrons can form covalent bonds with other atoms that have unpaired electrons available for bonding. Students should become adept at determining the number of unpaired electrons through electron configurations to predict such chemical behaviors.