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Chapter 9: Problem 10

Indicate the smallest and the largest species (atom or ion) in the following group: Al atom, F atom, As atom, \(\mathrm{Cs}^{+}\) ion, \(\mathrm{I}^{-}\) ion, \(\mathrm{N}\) atom.

Short Answer

Expert verified
From the given group the smallest species is the F atom and the largest is the \(\mathrm{I}^{-}\) ion.

Step by step solution

01

Understanding Atomic Radius Trend

Before diving into solving this problem, you need to understand how atomic radius behaves in the Periodic Table. The atomic radius generally increases as we go down a group due to the increase in quantum level (number of electron shells). It decreases as we move from left to right across a period on the Periodic Table due to higher effective nuclear charge.
02

Identify the Smallest Species

Given the atomic or ionic species: Al atom, F atom, As atom, \(\mathrm{Cs}^{+}\) ion, \(\mathrm{I}^{-}\) ion, \(\mathrm{N}\) atom. When referring to the Periodic Table, it is seen that Fluorine(F) atom is the furthest to the right and Nitrogen (N) atom is in the same period as Fluorine but to its left. Thus, Fluorine atom will be smaller than Nitrogen. The \(\mathrm{Cs}^{+}\) ion is a cation, meaning it has lost an electron, making it smaller than its parent atom, but still larger than atoms in upper periods. Therefore, the F atom is the smallest among them.
03

Identify the Largest Species

From the given species, Cesium (\(\mathrm{Cs}^{+}\) ion) belongs to the group which is at the bottom of the Periodic Table, implying it has the largest atomic radius among given atoms. However, the \(\mathrm{I}^{-}\) ion is an anion which means it has gained an electron, increasing repulsion and making it larger than its parent atom. It is larger than the \(\mathrm{Cs}^{+}\) ion, thus making it the largest species in this group.

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Most popular questions from this chapter

Give the symbol of the element (a) in group 14 that has the smallest atoms; (b) in period 5 that has the largest atoms; (c) in group 17 that has the lowest first ionization energy.

Arrange the following in expected order of increasing radius: $\mathrm{Br}, \mathrm{Li}^{+}, \mathrm{Se}, \mathrm{I}^{-} .$ Explain your answer.

The highest first ionization energy of the following is that of (a) \(\mathrm{Cs}_{i}\) (b) \(\mathrm{Cl}_{i}\) (c) I; (d) Li.

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

Neither \(\mathrm{Co}^{2+}\) nor \(\mathrm{Co}^{3+}\) has \(4 \mathrm{s}\) electrons in its electron configuration. How many unpaired electrons would you expect to find in each of these ions? Explain.
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