Periodic Table and Periodicity in Properties

Atomic size decreases because nuclear charge increases while electrons are added to the same shell.

Step 1: As we move left to right across a period, the atomic number increases, meaning the nuclear charge (number of protons) increases. Step 2: Electrons are added to the SAME outermost shell (same period = same principal quantum number). Step 3: The increased nuclear charge pulls all electrons, especially the outermost ones, closer to the nucleus. Step 4: The shielding effect remains roughly constant within a period because electrons are being added to the same shell. Step 5: Therefore, the effective nuclear charge (Zeff) felt by the outer electrons increases, pulling them inward and reducin

Na⁺ is smaller than Na because the same nuclear charge acts on fewer electrons.

Step 1: Na has 11 electrons. When Na loses 1 electron to form Na⁺, it has only 10 electrons. Step 2: The nuclear charge remains the same (11 protons) in both Na and Na⁺. Step 3: In Na⁺, the same nuclear charge of +11 now acts on only 10 electrons instead of 11. Step 4: This means each remaining electron experiences a greater pull from the nucleus. Step 5: As a result, the electron cloud contracts and Na⁺ is significantly smaller than Na. (Na radius ≈ 1.86 Å, Na⁺ radius ≈ 1.02 Å). This is a general rule: cations are always smaller than their parent atoms.

Nitrogen has a half-filled p-subshell (2p³) which is extra stable, making it harder to remove an electron compared to oxygen (2p⁴).

Step 1: Nitrogen's electronic configuration is 1s² 2s² 2p³. Its p-subshell is exactly half-filled, with one electron in each of the three 2p orbitals. Step 2: Half-filled subshells have extra stability due to symmetrical distribution of electrons and maximum exchange energy. Step 3: Oxygen's configuration is 1s² 2s² 2p⁴. One of the 2p orbitals has a paired electron. Step 4: This paired electron in oxygen experiences inter-electronic repulsion, making it easier (requires less energy) to remove one electron from oxygen. Step 5: Therefore, despite oxygen having a higher nuclear charge, its first

K < Na < Li < He

Step 1: Ionisation enthalpy decreases down a group because atomic size increases, and the outermost electrons are farther from the nucleus and more shielded. Step 2: For Group 1 (alkali metals): Li > Na > K. So the order of increasing IE is K < Na < Li. Step 3: Noble gases have the highest ionisation enthalpies in their respective periods because they have completely filled, very stable electron configurations. Step 4: He has the second highest ionisation enthalpy of all elements (2372 kJ/mol), much higher than Li (520 kJ/mol), Na (496 kJ/mol), or K (419 kJ/mol). Step 5: Therefore the correct