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Chapter 9: Problem 60

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?

Short Answer

Expert verified
a) Element A (Potassium) is a metal, b) Element B (Phosphorus) has greater ionization energy, c) Element A (Potassium) has a larger atomic radius, d) Element B (Phosphorus) has a greater electron affinity.

Step by step solution

01

Identification of Element

Identify elements A and B based on the given electron configurations. Element A's configuration ends at 4s¹. It occupies a position of Group 1 (Period 4) in the Periodic Table. This shows that element A is Potassium (K). Element B's configuration ends at 4p³. It occupies a position of Group 15 (Period 4) in the Periodic Table, indicating that B is Phosphorus (P).
02

Determination of Character

Identify whether the elements are metals or non-metals. The Group 1 elements are known to be highly reactive metals, including potassium. Therefore, element A is metal.
03

Ionization Energy

Next, we analyze the ionization energy, which is the energy required to remove an electron from an atom. The higher the group number, the higher the ionization energy generally is. Hence, Phosphorus (B) has a greater ionization energy.
04

Atomic Radius

Review the atomic radius, which depends on the location of the element in the Periodic Table. Elements on the left and lower part of the table have larger atomic radii due to increase in energy levels. Hence, Potassium (A) has a larger atomic radius.
05

Electron Affinity

Finally, determine the electron affinity, which refers to the ability of the atom to accept an additional electron. Non-metals usually have a higher electron affinity since they readily accept electrons to achieve a full outer energy level. Therefore, Phosphorus (B) has the greater electron affinity.

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Most popular questions from this chapter

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

Write electron configurations to show the first two ionizations for Cs. Explain why the second ionization energy is much greater than the first.

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.

Use ionization energies and electron affinities listed in the text to determine whether the following reaction is endothermic or exothermic. $$ \mathrm{Mg}(\mathrm{g})+2 \mathrm{F}(\mathrm{g}) \longrightarrow \mathrm{Mg}^{2+}(\mathrm{g})+2 \mathrm{F}^{-}(\mathrm{g}) $$

The most negative electron affinity of the following elements is that of (a) \(\mathrm{Br} ;\) (b) \(\mathrm{Sn} ;\) (c) \(\mathrm{Ba} ;\) (d) Li.
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