An element has atomic number 79. Predict the group and period in which the element is placed. (a) \(2^{\text {nd }}\) group, \(7^{\text {th }}\) period (b) \(11^{\text {th }}\) group, \(6^{\text {th }}\) period (c) \(13^{\text {th }}\) group, \(6^{\text {th }}\) period (d) \(12^{\text {th }}\) group, \(6^{\text {th }}\) period

The atomic number is foundational to the organization of the periodic table and has significant implications on an element's properties. It is denoted as 'Z' and it precisely quantifies the number of protons contained within the nucleus of an atom. Since protons carry a positive charge, the atomic number also helps in identifying the charge of the nucleus.

For instance, with an atomic number of 79, an element possesses 79 protons and, in a neutral atom, an equal number of electrons. The atomic number essentially dictates the chemical identity of the element; elements are defined and differentiated by their atomic numbers. Knowing that the atomic number is linked to the element's position on the periodic table, we can use it to predict the behavior and characteristics of the element, such as reactivity and the types of bonds it forms.

Electron shells are layers that surround the nucleus of an atom where electrons are most likely to be found. These shells are arranged in levels of increasing energy, each of which can hold a certain maximum number of electrons. For example, the first shell can hold up to 2 electrons, the second up to 8, and so on, following a pattern described by the formula 2n^2, where 'n' is the shell level number.

The period of an element in the periodic table indicates the highest energy level of its electron shells that contains electrons. In simple terms, the period number matches the number of electron shells in an atom. This is why elements with atomic numbers from 58 to 71 are placed in the 6th period; they all have electron shells that extend to the 6th level. Periods also show a progression of properties, such as melting points and reactivities, influenced by the number of electron shells and their configurations.

Valence electrons are the outermost electrons of an atom and are crucial in determining an element's chemical properties and the way it will bond with other elements. These electrons are the 'frontline' participants in chemical reactions as they can be gained, lost, or shared in chemical bonds.

Each element strives for a stable electronic configuration, frequently achieved by having a complete outer shell, similar to the noble gases; this concept is known as the octet rule. The number of valence electrons not only determines the type of compound that an element will form (ionic or covalent) but also the group it belongs to in the periodic table. For example, elements in Group 1 have one valence electron, making them highly reactive and ready to lose that electron to achieve stability.

Transition metals encompass a series of elements that reside in the middle of the periodic table in groups 3 through 12. These metals have incompletely filled d-subshells or readily give rise to cations with incomplete d-subshells. They are known for their distinctive properties such as the ability to form colorful compounds, exhibit variable oxidation states, and possess a magnetic nature.

These elements play an integral role in modern industry and biological systems. For example, iron is a critical element for blood, while catalysts for chemical reactions often involve transition metals. With a unique combination of durability and conductive properties, they are pivotal in numerous applications from electronics to architecture. The element with atomic number 79, which we've identified to be in the 6th period, is also part of the transition metals, exhibiting these typical properties and group characteristics.