Answer:option d==> Si.
Explanation:
The energy required to remove electron from a gaseous atom or ion is what is called an ionization energy. As we remove electrons continually in a gaseous atom or ion, the ionization energy increases which are know as the first ionization energy, the second ionization energy, third ionization energy and so on.
Looking at the electronic configuration of Silicon, Si; Ne 3s2 3p2. We can can see that the first four ionization energies are from the removal of the 3p2 and 3s2 electrons and the fifth ionization energy, which is the highest ionization energy of 14800 kJ/mol is the the electron removed from the core shell.
The element having these ionization energy values is Al.
Ionization energy is the energy required to remove an electron from an atom. The first ionization energy is always less than the second ionization energy and so on.
The ionization energy continues to increase smoothly depending on the number of electrons on the valence shell.
As we go past the valence shell, we experience a quantum jump. The ionization energy suddenly increases by a very large value. This owes to the fact that "inner" electrons are now being removed and this required a very large amount of energy.
If we look at the first three values of the ionization energy in the question; 577.9 kJ/mol , 1820 kJ/mol, 2750 kJ/mol, we will notice that ionization energy increases smoothly. This implies that these three electrons reside on the same shell.
The fourth ionization energy suddenly becomes outrageously high (11600 kJ/mol) indicating that the electron is being removed from an inner shell.
Al has three electrons on its valence shell, therefore, the element must be Al.
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