The elements Ga, Ge, As, Se, and Br lie in the same period 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.

The periodic table is more than just a tabular arrangement of chemical elements; it's a powerful tool for predicting the chemical behavior and properties of each element. Elements are ordered by increasing atomic number, which is the number of protons in the nucleus, and are arranged into rows called 'periods' and columns known as 'groups' or 'families.' Each period corresponds to the highest energy level that is occupied by electrons in the ground state of the atoms of the elements in that period.

For instance, the elements Ga, Ge, As, Se, and Br all lie in Period 4 of the periodic table. This means that the principal quantum level 'n' being filled with electrons is the fourth one. Knowing this, we can begin to write out the electron configuration, as was done in the exercise, by filling each sublevel in order of increasing energy until we reach the atomic number specific to each element. The sublevels are filled according to the Aufbau principle, Hund's rule, and the Pauli exclusion principle, which help explain the distribution and spin of electrons in an atom's orbitals.

Unpaired electrons play a crucial role in the magnetic and chemical properties of an atom. An unpaired electron is one that occupies an orbital alone, not having a paired electron with opposite spin. According to Hund's rule, electrons will fill degenerate orbitals singly first, with parallel spins, before any pairing occurs. This rule is integral in predicting the number of unpaired electrons in an atom.

For example, the 4p sublevel can hold up to 6 electrons and consists of three orbitals. Looking at the ground-state electron configurations for elements Ga through Br, we identify their unpaired electrons by observing the 4p electrons. As shown in the solution, Ga has one 4p electron, which is unpaired, while Ge has two, both unpaired in different orbitals. As has three unpaired electrons, one in each of the three 4p orbitals. Se follows with one pair and two unpaired electrons, and Br has two pairs and one unpaired electron. Recognizing the patterns of how electrons fill orbitals allows us to predict magnetic properties and reactivity, as elements with unpaired electrons tend to have magnetic moments and are usually more chemically reactive.

The electron configuration of an atom describes the distribution of electrons in the atom's orbitals when the atom is in its lowest energy state, known as the ground state. The 'ground state' implies that the electrons have filled the available orbitals starting with the lowest energy levels first, without any extra energy being added to the system.

In the case of the Period 4 elements Ga, Ge, As, Se, and Br, their electron configurations start with the noble gas from the previous period, [Ar], as shorthand for the configuration of argon. This is followed by a systematic filling of the 4s and 3d sublevels, which are lower in energy than the 4p sublevel within the same principal quantum level. The configurations given for the atoms in the exercise reflect the stable arrangements of electrons that minimize the energy of the atom. This foundational understanding is essential for studying chemistry further because the ground-state electron configuration is the starting point for understanding how atoms interact, bond, and form compounds.