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Chapter 9: Problem 21

Use principles established in this chapter to arrange the following atoms in order of increasing value of the first ionization energy: $\mathrm{Sr}, \mathrm{Cs}, \mathrm{S}, \mathrm{F}, \mathrm{As}.$

Short Answer

Expert verified
The correct order of the atoms from lowest to highest first ionization energy is Cs, Sr, As, S, F.

Step by step solution

01

Locate elements on the periodic table

Identify the positions of Sr (Strontium), Cs (Cesium), S (Sulphur), F (Fluorine), and As (Arsenic) on the periodic table. Strontium and Cesium are in the same group (Group 2), and Cesium is below Strontium. Sulphur, Fluorine, and Arsenic are in different periods but in the same group (Group 16), with Arsenic below Sulphur and Fluorine.
02

Rank based on groups

According to the ionization energy trend in the periodic table, ionization energy decreases down a group. Therefore, between Strontium and Cesium, Cesium will have the lower ionization energy. Similarly, between Sulphur, Fluorine, and Arsenic, Arsenic will have the lowest ionization energy.
03

Rank based on periods

Ionization energy increases across a period from left to right. Therefore, between Strontium and Arsenic, Strontium will have lower ionization energy since it is further left on the periodic table. Similarly, between Sulphur and Fluorine, Sulphur will have lower ionization energy since it is further left on the periodic table.
04

Final arrangement

Based on the steps above, the final arrangement of the atoms from lowest to highest first ionization energy should be: Cs, Sr, As, S, F.

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Most popular questions from this chapter

For the atom \(^{119} \mathrm{Sn}\), indicate the number of (a) protons in the nucleus; (b) neutrons in the nucleus; (c) \(4 d\) electrons; (d) 3s electrons; (e) 5 \(p\) electrons; (f) electrons in the valence shell.

Compare the elements \(\mathrm{Na}, \mathrm{Mg}, \mathrm{O},\) and \(\mathrm{P}.\) (a) Place the elements in order of increasing ionization energy. (b) Place the elements in order of increasing electron affinity.

Two elements, \(A\) and \(B\), have the electron configurations shown. $$ \mathrm{A}=\left[\begin{array}{ll} \mathrm{Ar} & 4 s^{1} \end{array} \quad \mathrm{B}=\left[\begin{array}{l} \mathrm{Ar} \end{array}\right] 3 d^{10} 4 \mathrm{s}^{2} 4 p^{3}\right. $$ (a) Which element is a metal? (b) Which element has the greater ionization energy? (c) Which element has the larger atomic radius? (d) Which element has the greater electron affinity?

Plot a graph of the square roots of the ionization energies versus the nuclear charge for the two series \(\mathrm{Li}, \mathrm{Be}^{+}, \mathrm{B}^{2+}, \mathrm{C}^{3+},\) and \(\mathrm{Na}, \mathrm{Mg}^{+}, \mathrm{Al}^{2+}, \mathrm{Si}^{3+} .\) Explain the observed relationship with the aid of Bohr's expression for the binding energy of an electron in a one electron atom.

Mendeleev's periodic table did not preclude the possibility of a new group of elements that would fit within the existing table, as was the case with the noble gases. Moseley's work did preclude this possibility. Explain this difference.
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