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Chapter 9: Problem 33

Unpaired electrons are found in only one of the following species. Indicate which one, and explain why: \(\mathrm{F}^{-}, \mathrm{Ca}^{2+}, \mathrm{Fe}^{2+}, \mathrm{S}^{2-}\)

Short Answer

Expert verified
Among the given species, only the ion \(\mathrm{Fe}^{2+}\) has unpaired electrons.

Step by step solution

01

Electron Configuration of \(\mathrm{F}^{-}\)

The electron configuration of neutral Flourine atom (F) is \(1s^{2} 2s^{2} 2p^{5}\). When Flourine gains an electron to become \(\mathrm{F}^{-}\), the electron configuration becomes \(1s^{2} 2s^{2} 2p^{6}\). There are no unpaired electrons, hence \(\mathrm{F}^{-}\) does not have unpaired electrons.
02

Electron Configuration of \(\mathrm{Ca}^{2+}\)

The electron configuration of neutral Calcium atom (Ca) is \(1s^{2} 2s^{2} 2p^{6} 3s^{2} 3p^{6} 4s^{2}\). When Calcium loses two electrons to become \(\mathrm{Ca}^{2+}\), the electron configuration becomes \(1s^{2} 2s^{2} 2p^{6} 3s^{2} 3p^{6}\). There are no unpaired electrons, hence \(\mathrm{Ca}^{2+}\) does not have unpaired electrons.
03

Electron Configuration of \(\mathrm{Fe}^{2+}\)

The electron configuration of neutral Iron atom (Fe) is \(1s^{2} 2s^{2} 2p^{6} 3s^{2} 3p^{6} 4s^{2} 3d^{6}\). When Iron loses two electrons to form \(\mathrm{Fe}^{2+}\), the electron configuration becomes \(1s^{2} 2s^{2} 2p^{6} 3s^{2} 3p^{6} 3d^{6}\). Thus, there are unpaired electrons in \(\mathrm{Fe}^{2+}\).
04

Electron Configuration of \(\mathrm{S}^{2-}\)

The electron configuration of neutral Sulfur atom (S) is \(1s^{2} 2s^{2} 2p^{6} 3s^{2} 3p^{4}\). When Sulfur atom gains two electrons to form \(\mathrm{S}^{2-}\), the electron configuration becomes \(1s^{2} 2s^{2} 2p^{6} 3s^{2} 3p^{6}\). There are no unpaired electrons, hence \(\mathrm{S}^{2-}\) does not have unpaired electrons.

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Most popular questions from this chapter

Among the following ions, several pairs are isoelectronic. Identify these pairs. \(\mathrm{Fe}^{2+}, \mathrm{Sc}^{3+}, \mathrm{Ca}^{2+}, \mathrm{F}^{-}\) \(\mathrm{Co}^{2+}, \mathrm{Co}^{3+}, \mathrm{Sr}^{2+}, \mathrm{Cu}^{+}, \mathrm{Zn}^{2+}, \mathrm{Al}^{3+}.\)

For each of the following pairs, indicate the atom that has the larger size: (a) Te or Br; (b) \(\mathrm{K}\) or \(\mathrm{Ca} ;\) (c) Ca or Cs; (d) \(\mathrm{N}\) or \(\mathrm{O} ;\) (e) \(\mathrm{O}\) or \(\mathrm{P} ;\) (f) \(\mathrm{Al}\) or \(\mathrm{Au}.\)

Describe how the ionization energies of the ions \(\mathrm{He}^{-}, \mathrm{Li}^{-}, \mathrm{Be}^{-}, \mathrm{B}^{-}, \mathrm{C}^{-}, \mathrm{N}^{-}, \mathrm{O}^{-},\) and \(\mathrm{F}^{-}\) vary with atomic number.

Which of the following ions are unlikely to be found in chemical compounds: \(\mathrm{K}^{+}, \mathrm{Ga}^{4+}, \mathrm{Fe}^{6+} \mathrm{S}^{2-}, \mathrm{Ge}^{5+},\) or \(\mathrm{Br}^{-} ?\) Explain briefly.

An ion that is isoelectronic with \(\mathrm{Se}^{2-}\) is (a) \(\mathrm{S}^{2-} ;\) (b) I \(^{-}\) (c) \(\mathrm{Xe} ;\) (d) \(\mathrm{Sr}^{2+}.\)
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