The number of electrons in the outermost shell of alkali metals is 1. (Alkali metals belong to Group 1 and have valency 1).

Alkaline earth metals have valency 2. This means that their position in the modern periodic table is in Group 2.

Since the molecular formula of chloride is XCl, the valency of element X is 1. Among the given options, only sodium (Na) has a valency of 1 (electronic configuration 2, 8, 1) and belongs to Group 1, forming an ionic solid NaCl with a high melting point.

Reason: Beryllium (Z = 4) has an electronic configuration of (2, 2) with 2 valence electrons and belongs to Group 2, having similar chemical properties to Magnesium (2, 8, 2).

• Moving from left to right in a period, the atomic number increases by 1 at each successive element, which means the positive nuclear charge (number of protons) increases progressively.
• However, the newly added electrons are added into the same valence shell; no new electron shell is created.
• Due to the increased positive charge on the nucleus, the valence electrons experience a stronger electrostatic pull towards the center of the atom.
• Therefore, the electrons are pulled closer to the nucleus, which decreases the atomic radius across a period.

• Metallic character is defined as the tendency of an atom to lose its valence electrons and form positive ions (cations / electropositivity).
• As we move from left to right across a period, the effective nuclear charge increases while the atomic radius decreases.
• Because the valence electrons are pulled more tightly by the nucleus, the ionization energy increases and it becomes progressively harder for atoms to lose electrons.
• Instead, the tendency to gain electrons (electronegativity / nonmetallic character) increases. Therefore, metallic character decreases across a period.

• Moving down any group, a completely new electron shell is added at each successive element.
• Although the nuclear charge also increases, the addition of a new outer shell significantly increases the physical distance between the nucleus and the outermost valence electrons.
• The inner electron shells also exert an effective screening (shielding) effect, weakening the net nuclear attraction on the outer electrons.
• As a result, the size of the atom expands, and the atomic radius goes on increasing down a group.

• The valency of an element is determined by the number of valence electrons present in its outermost shell.
• In the Modern Periodic Table, all elements placed in the same group have the same number of valence electrons.
• Since the number of valence electrons governs the number of electrons lost, gained, or shared to achieve a stable octet, their combining capacity remains identical.
• Hence, all elements in a given group share the same valency (e.g., all Group 1 elements have valency 1; all Group 2 elements have valency 2).

• Group 18 elements (noble gases: Helium, Neon, Argon, Krypton, Xenon, Radon) have completely filled outermost valence electron shells.
• Helium has a stable duplet configuration (2 electrons in K shell), while all other noble gases have a stable octet configuration (8 valence electrons).
• Because their outermost shells are completely saturated, they have zero valency and no tendency to lose, gain, or share electrons under standard conditions.
• Therefore, noble gases are very stable and chemically unreactive.

Statement: "The physical and chemical properties of elements are periodic functions of their atomic masses."

Key Highlights & Merits:

• Dmitri Mendeleev organized the 63 known elements in increasing order of atomic masses in horizontal rows (periods) and vertical columns (groups).
• He placed elements with similar chemical properties (hydrides and oxides formulae) in the same group.
• He corrected atomic masses of certain elements (e.g. Beryllium corrected from 14.09 to 9.4).
• He left vacant places for undiscovered elements and remarkably predicted their properties accurately:
  • Eka-boron → Scandium (Sc)
  • Eka-aluminium → Gallium (Ga)
  • Eka-silicon → Germanium (Ge)

The Modern Periodic Table, created by Henry Moseley, organizes elements in increasing order of their atomic number (Z):

• Periods: 7 horizontal rows called periods (numbered 1 to 7). Period 1 has 2 elements; Periods 2 and 3 have 8 elements each.
• Groups: 18 vertical columns called groups (numbered 1 to 18). Elements in the same group have the same number of valence electrons.
• Four Blocks:
  • s-block: Groups 1 and 2 (Alkali and alkaline earth metals).
  • p-block: Groups 13 to 18 (Metals, nonmetals, metalloids, and noble gases).
  • d-block: Groups 3 to 12 (Transition metals).
  • f-block: Two separate series placed at the bottom: Lanthanides and Actinides (Inner transition elements).
• Zigzag Line: A border line in the p-block separating metals on the left from nonmetals on the right, with metalloids (Si, Ge, As, Sb, Te, Po) lining the border.

• In Mendeleev’s Periodic Table: Mendeleev arranged elements according to atomic mass. Since isotopes of an element have different atomic masses (e.g. Cl-35 and Cl-37), they should theoretically occupy different positions. However, they share identical chemical properties, warranting the same position. Mendeleev could not resolve this ambiguity.
• In Modern Periodic Table: Moseley classified elements according to atomic number (Z). Because all isotopes of an element have the exact same atomic number (number of protons), they all fit into the same single place in the modern periodic table.

1. Atomic masses of some elements were revised to give them proper places in accordance with their properties (e.g., Beryllium mass revised from 14.09 to 9.4 u).
2. Mendeleev left vacant spaces for undiscovered elements and accurately predicted their existence, atomic mass, and chemical formulas of their oxides and chlorides (e.g., Eka-aluminium became Gallium).
3. When noble gases (Helium, Neon, Argon) were discovered at the end of the nineteenth century, Mendeleev created the 'zero group' without disturbing the existing table structure.

1. Ambiguity of Cobalt and Nickel: Cobalt (58.9) and Nickel (58.7) have nearly identical atomic masses, creating confusion over their sequence.
2. Position of Isotopes: Isotopes have different atomic masses but identical chemical properties, presenting an insolvable placement dilemma in an atomic-mass based system.
3. Irregular rise in atomic masses: The rise in atomic masses between successive elements was not uniform, making it impossible to predict how many elements lay between two heavier metals.
4. Position of Hydrogen: Hydrogen resembles both alkali metals (Group 1) in forming halides/oxides and halogens (Group 17) in forming diatomic molecules (H₂), making its placement uncertain.