Among the following, the most characteristic oxidation states for lead and tin are, respectively:

1. +4, +2

2. +2, +4

3. +4, +4

4. +2, +2

A configuration with the lowest ionization enthalpy among the following is:
1. \(1 s^2 2 s^2 2 p^5\)$\mathrm{}$
2. \(1 s^2 2 s^2 2 p^3\)
3. \(1 s^2 2 s^2 2 p^6 3 s^1\)
4. \(1 s^2 2 s^2 2 p^6\)

For the second-period elements, the correct increasing order of first ionisation enthalpy is:

The incorrect order of electronegativity is :

The correct order of the decreasing ionic radii among the following isoelectronic species is:

1. $C{a}^{2+}>K^{+}>S^{2-}>C{l}^{-}$

2. $C{l}^{-}>S^{2-}>C{a}^{2+}>K^{+}$

3. $S^{2-}>C{l}^{-}>K^{+}>C{a}^{2+}$

4. $K^{+}>C{a}^{2+}>C{l}^{-}>S^{2-}$

Which of the following represents the correct order of ionic radii?

The formation of the oxide ion, O²⁻(g), from an oxygen atom occurs in two successive steps:

O(g) + e⁻ → O⁻(g)  ΔH = −141 kJ mol⁻¹

O⁻(g) + e⁻ → O²⁻(g)  ΔH = +780 kJ mol⁻¹

Although O²⁻ is isoelectronic with neon, its formation in the gaseous state is energetically unfavourable. What is the main reason for this?

Which of the following third-period elements has a positive electron gain enthalpy?

Which of the following pairs consists of elements that belong to the same period of the Periodic Table?