The atomic masses in the periodic table are relative masses and average masses. Explain.

Atoms are the basic building blocks of matter. They consist of a nucleus, which contains protons and neutrons, and electrons that orbit the nucleus. Protons and electrons have a charge, with protons being positively charged and electrons being negatively charged. Neutrons have no charge. The number of protons in the nucleus determines the element to which the atom belongs.

Isotopes are atoms of the same element that have different numbers of neutrons in their nucleus. This means they have the same number of protons and electrons, but different masses because the number of neutrons varies. Each isotope has a specific mass, and the atomic mass for a given element is generally different for varying isotopes.

Atomic mass in the periodic table is an average of the masses of all isotopes of an element, weighted by the relative abundance of each isotope. The abundant isotopes have a higher contribution to the average mass, while the less abundant isotopes contribute less to the average. As a result, the atomic mass listed for an element in the periodic table represents an average value and not the mass of a single isotope.

Atomic masses measured in atomic mass units (amu) are relative, meaning they are compared to the mass of a reference element. The reference element used is carbon-12, with a mass of exactly 12 amu. An atomic mass unit is defined as 1/12 of the mass of a carbon-12 atom. Relative masses are used because the actual masses of atoms are extremely small, making it difficult to work with them directly. By comparing them to a reference element, it becomes more convenient to express and compare the atomic masses of different elements. In summary, atomic masses in the periodic table are relative masses because they are compared to the mass of a reference element, carbon-12, and they are average masses because they represent the weighted average of the masses of all isotopes for a given element, taking into account the relative abundance of each isotope.