Which of the following observations are incorrect for the noble gases - He, Ne, Ar, Kr, and Xe?
Which of the following observations are incorrect for the noble gases - He, Ne, Ar, Kr, and Xe?
The correct option is B The given noble gases belong to the far right of the periodic table and exhibit very high electronegativity
The electron affinity is the negative of electron gain enthalpy. Given that the valence electron configuration is ns2 np6. Adding an electron implies that the incoming electron should be filled in a new shell. So definitely electron affinity for noble gases is negative. In simple terms, we can interpret the term “electron affinity” as the tendency of an atom to attract an electron. Since the group 18 – noble gas elements have a stable valence electron configuration, these elements tend not to attract electrons.
What can we predict about removing an electron from the stable, fully filled valence orbitals? Since the ground state is very stable, there will be a strong resistance to remove any electrons from the fully filled ns2 np6 orbitals.
These elements tend to exist as monoatomic gases at room temperature. Can you think why? This is again because of the stable valence electron configuration. The orbitals are all fully filled so there is no need for any more bonding. This only explains the monoatomicity. Why are they gases? In their solid or liquid states, the dispersion forces holding these atoms together are very weak.
What can we infer about the electro negativities of these elements?
Helium Neon and Argon do not even have a value for electronegativity. In case there was a value, it would make no sense as well. Krypton and Xenon do have electronegativity values (Pauling scale) but they are lesser than that of their halogen counterparts.
The very high ionization energies can also be attributed to the fact that at the far right of the periodic table, the effective nuclear charge is progressively higher.
So only option b is incorrect. Also, these gases are colourless and odourless. They do not support combustion.
The electron affinity is the negative of electron gain enthalpy. Given that the valence electron configuration is ns2 np6. Adding an electron implies that the incoming electron should be filled in a new shell. So definitely electron affinity for noble gases is negative. In simple terms, we can interpret the term “electron affinity” as the tendency of an atom to attract an electron. Since the group 18 – noble gas elements have a stable valence electron configuration, these elements tend not to attract electrons.
What can we predict about removing an electron from the stable, fully filled valence orbitals? Since the ground state is very stable, there will be a strong resistance to remove any electrons from the fully filled ns2 np6 orbitals.
These elements tend to exist as monoatomic gases at room temperature. Can you think why? This is again because of the stable valence electron configuration. The orbitals are all fully filled so there is no need for any more bonding. This only explains the monoatomicity. Why are they gases? In their solid or liquid states, the dispersion forces holding these atoms together are very weak.
What can we infer about the electro negativities of these elements?
Helium Neon and Argon do not even have a value for electronegativity. In case there was a value, it would make no sense as well. Krypton and Xenon do have electronegativity values (Pauling scale) but they are lesser than that of their halogen counterparts.
The very high ionization energies can also be attributed to the fact that at the far right of the periodic table, the effective nuclear charge is progressively higher.
So only option b is incorrect. Also, these gases are colourless and odourless. They do not support combustion.
