Moseley's experiments on \(X\) rays emitted from atoms led to the concept of atomic numbers. (a) If arranged in order of increasing atomic mass, which element would come after chlorine? (b) Describe two ways in which the properties of this element differ from the other elements in group 8A.

The periodic table is a fundamental tool in chemistry, organizing all known elements in an informative array based on their atomic numbers, which represent the number of protons in an atom's nucleus. This ingenious chart classifies elements into groups and periods, revealing patterns in their chemical behavior.

As you move from left to right across a period, the atomic numbers increase. The modern periodic table arranges elements in order of increasing atomic number rather than atomic mass, a concept stemming from the pioneering work of Henry Moseley and his experiments with X-rays. This ordering not only reflects the elements' physical and chemical properties but also their electron configurations, leading to the periodic nature of the table.

For instance, if we consider chlorine with the atomic number 17, the next element in order is argon, with the atomic number 18. Students often find this arrangement helpful for predicting how an element will react chemically because elements that belong to the same group (vertical columns on the table) typically exhibit similar chemical behavior.

Chemical properties of elements are deeply interconnected with their position on the periodic table. Two such properties worth noting are atomic radius and ionization energy, which can vastly differ even among elements within the same group.

Atomic radius refers to the size of an atom, generally measured from the nucleus to the outer edge of the electron cloud. This size can affect an element's reactivity and bonding with others. As elements progress down a group, the radius typically increases due to additional electron shells.

Another critical property is ionization energy—the energy required to remove an electron from an atom in its gaseous state. Elements with higher ionization energies tend to be less reactive, as they are more reluctant to lose electrons. Across a period, ionization energy generally increases due to the rising nuclear charge, pulling electrons closer and making their removal more energy-intensive.

By understanding these properties, students can predict or explain the reactivity and other chemical behaviors of elements, enhancing their grasp of chemistry beyond rote memorization.

Noble gases, located in group 8A—or group 18, using the IUPAC nomenclature—of the periodic table, are characterized by their remarkably low reactivity. This inert nature is due to their full valence electron shells, which make noble gases quite stable and generally resistant to forming bonds with other elements.

Argon, the noble gas that follows chlorine on the periodic table, serves as an exemplary case. Despite sharing the same group with other noble gases such as helium, neon, krypton, xenon, and radon, argon exhibits distinct chemical properties. Its atomic radius and ionization energy are pivotal in determining its stand within the group. Compared to its heavier counterparts, argon's relatively smaller atomic radius and higher ionization energy suggest that while still inert, its behaviors in certain situations can differ from other noble gases.

Understanding the nuances of noble gases, including argon's unique properties, not only sheds light on the incredible order of the periodic table but also assists students in grasping broader chemical concepts. It enables them to see the relative changes between each noble gas, rather than viewing the group as a monolith of inertness.