Which of the following statements are true?

A.    Unlike Ga that has a very high melting point, Cs has a very low melting point.
B.     On Pauling scale, the electronegativity values of N and Cl are not the same.

C.     Ar, K⁺, Cl^–, Ca²⁺, and S^2– are all isoelectronic species.

D.     The correct order of the first ionization enthalpies of Na, Mg, Al, and Si is Si >Al > Mg > Na.

E.      The atomic radius of Cs is greater than that of Li and Rb.

The correct option is C and E only.

Let us analyze each statement individually to determine whether it is true or false:

Statement A: Unlike Ga that has a very high melting point, Cs has a very low melting point.
This statement is false. Gallium (Ga) has an unusually low melting point of approximately 303 K (29.76 °C), meaning it can melt in the palm of a hand. Cesium (Cs) also has a very low melting point of approximately 301.6 K (28.4 °C). Therefore, both elements have very low melting points, making the statement incorrect.

Statement B: On Pauling scale, the electronegativity values of N and Cl are not the same.
This statement is false. On the Pauling scale of electronegativity, both Nitrogen (N) and Chlorine (Cl) have the same electronegativity value, which is approximately 3.0.

Statement C: Ar, K⁺, Cl^-, Ca²⁺, and S^2- are all isoelectronic species.
This statement is true. Isoelectronic species are atoms, ions, or molecules that have the same number of electrons. Let us count the electrons for each species:
- Argon (Ar): Atomic number = 18. Number of electrons = 18.
- Potassium ion (K⁺): Atomic number of K = 19. Since it carries a +1 charge, it has 19 - 1 = 18 electrons.
- Chloride ion (Cl^-): Atomic number of Cl = 17. Since it carries a -1 charge, it has 17 + 1 = 18 electrons.
- Calcium ion (Ca²⁺): Atomic number of Ca = 20. Since it carries a +2 charge, it has 20 - 2 = 18 electrons.
- Sulfide ion (S^2-): Atomic number of S = 16. Since it carries a -2 charge, it has 16 + 2 = 18 electrons.
All these species contain exactly 18 electrons, so they are indeed isoelectronic.

Statement D: The correct order of the first ionization enthalpies of Na, Mg, Al, and Si is Si > Al > Mg > Na.
This statement is false. Generally, first ionization enthalpy increases across a period from left to right due to an increase in effective nuclear charge. However, there is an exception between Magnesium (Mg) and Aluminum (Al).
The electronic configuration of Mg is [Ne] 3s² (fully filled, stable s-orbital), while that of Al is [Ne] 3s² 3p¹. It is easier to remove an electron from the higher energy 3p orbital of Al than from the stable, fully-filled 3s orbital of Mg. As a result, the first ionization enthalpy of Mg is higher than that of Al.
Therefore, the correct increasing order of first ionization enthalpy is:
$\mathrm{Na}<\mathrm{Al}<\mathrm{Mg}<\mathrm{Si}$
Or in decreasing order:
$\mathrm{Si}>\mathrm{Mg}>\mathrm{Al}>\mathrm{Na}$

Statement E: The atomic radius of Cs is greater than that of Li and Rb.
This statement is true. In the periodic table, Lithium (Li), Rubidium (Rb), and Cesium (Cs) all belong to Group 1 (alkali metals). The atomic radius increases as we go down the group because new electronic shells are added.
The order down the group is:
$\mathrm{Li}<\mathrm{Na}<K<\mathrm{Rb}<\mathrm{Cs}$
Hence, Cesium (Cs) has a larger atomic radius than both Lithium (Li) and Rubidium (Rb).

Thus, only statements C and E are true.