Periodic Table of Elements

Dmitri Mendeleev is recognized as the 'Father of the Periodic Table' for his pioneering work in organizing the elements based on their atomic mass and properties. In 1869, he created the first version of the periodic table, which not only arranged elements in a systematic way but also predicted the existence and properties of undiscovered elements. Mendeleev's approach allowed for a clearer understanding of chemical relationships, laying the foundation for modern chemistry and the eventual development of the periodic law.

Henry Moseley's discovery that elements are organized by atomic number rather than atomic mass revolutionized the periodic table. He established that the atomic number, which represents the number of protons in an atom's nucleus, is a more fundamental property of elements. This approach resolved inconsistencies in the arrangement of elements that arose when using atomic mass. By organizing elements based on atomic number, Moseley provided a clearer and more logical framework for understanding elemental properties and their relationships, leading to the modern periodic table we use today.

In the periodic table, horizontal rows are referred to as periods, which indicate the number of electron shells surrounding the nucleus of an atom. Each period corresponds to the highest occupied energy level of electrons in the elements of that row. As you move from left to right across a period, the atomic number increases, and the properties of the elements change in a predictable manner, reflecting their electron configurations. This organization helps in understanding the relationships and trends among different elements.

Atomic size increases from top to bottom in a group due to the addition of electron shells, which outweighs the increasing nuclear charge. Conversely, atomic size decreases from left to right across a period because the increasing positive charge in the nucleus pulls electrons closer to the nucleus, resulting in a smaller atomic radius. This trend highlights the balance between electron-electron repulsion and the attractive force of the nucleus.

A sodium atom contains 11 electrons, while a sodium ion has only 10 due to the loss of one outermost electron. This loss reduces electron-electron repulsion and allows the remaining electrons to be pulled closer to the positively charged nucleus, resulting in a smaller ionic radius. Consequently, the sodium ion is smaller than the neutral sodium atom.

Inner transition elements include lanthanoids and actinoids. Lanthanoids are indeed all metals, known for their metallic properties and use in various applications. Actinoids, on the other hand, are primarily radioactive, with many of their isotopes being unstable. Furthermore, all elements beyond uranium in the periodic table are synthetic and created through nuclear reactions, confirming their artificial production. Thus, the statements regarding the metallic nature of lanthanoids, the radioactivity of actinoids, and the artificial creation of elements after uranium are accurate.

Copernicium (Cn) is correctly matched with atomic number 112, reflecting its discovery and naming in honor of astronomer Nicolaus Copernicus. Livermorium (Lv) is accurately assigned atomic number 116, named after the Lawrence Livermore National Laboratory. Nihonium, named after Japan, corresponds to atomic number 113, recognizing Japan's contributions to chemistry. The other options, particularly Flerovium which is assigned atomic number 114, are incorrect. Thus, the matches for Copernicium, Livermorium, and the naming of Nihonium are validated by their respective atomic numbers and origins.