B has a smaller first ionization enthalpy than Be. Consider the following statements:
(i) it is easier to remove electron than electron.
(ii) electron of is more shielded from the nucleus by the inner core of electrons than the electrons of Be.
(iii) electron has more penetration power than electron
(iv) atomic radius of B is more than Be (atomic number , )
The correct statements are:
(i) it is easier to remove electron than electron.
(ii) electron of is more shielded from the nucleus by the inner core of electrons than the electrons of Be.
(iii) electron has more penetration power than electron
(iv) atomic radius of B is more than Be (atomic number , )
The correct statements are:
- (A)(i), (iii) and (iv)
- (B)(i), (ii) and (iii)✓ Correct
- (C)(i), (ii) and (iv)
- (D)(ii), (iii) and (iv)
Answer: (B)
Solution
Boron's outermost electron occupies a orbital whereas beryllium's is a electron. A electron lies at higher energy and is held less tightly, so it is easier to pull off than a electron, which makes statement (i) correct. That same electron is screened more effectively from the nucleus by the filled inner core than beryllium's electron is, so statement (ii) is correct, and because a orbital penetrates closer to the nucleus than a orbital, statement (iii) is correct too. Statement (iv) is wrong: across a period the growing nuclear charge contracts the atoms, so boron is in fact smaller than beryllium, not larger. The genuine reasons for boron's lower first ionisation enthalpy are therefore statements (i), (ii) and (iii).
JEE Main · 2020 · 09 Jan 2020 · Shift 1 · Single correct option