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Chapter 9: Problem 85

Answer each of the following questions: (a) Which of the elements \(P, A s,\) and \(S\) has the largest atomic radius? (b) Which of the following has the smallest radius: \(\mathrm{Xe}, \mathrm{O}^{2-}, \mathrm{N}^{3-},\) or \(\mathrm{F}^{-} ?\) (c) Which should have the largest difference between the first and second ionization energy: \(\mathrm{Al}, \mathrm{Si}, \mathrm{P},\) or \(\mathrm{Cl} ?\) (d) Which has the largest ionization energy: \(\mathrm{C}, \mathrm{Si}\), or \(\mathrm{Sn}\) ? (e) Which has the largest electron affinity: \(\mathrm{Na}, \mathrm{B}\) \(\mathrm{Al},\) or \(\mathrm{C} ?\)

Short Answer

Expert verified
The answers are: (a) P, (b) \(F^{-}\), (c) Al, (d) C, (e) C.

Step by step solution

01

Identify the trends

In general, atomic radius increases from top to bottom in a group, and decreases from left to right across a period. Ionization energy decreases from top to bottom in a group, and increases from left to right across a period. Electron affinity decreases from top to bottom in a group, and increases from left to right across a period.
02

Answer question (a)

Given the elements P (Phosphorus), As (Arsenic), and S (Sulfur), use the atomic radius trend. As these elements are found in the same period but different groups, the atomic radius decreases from left to right. So, the element with the largest atomic radius is P (Phosphorus).
03

Answer question (b)

When looking at Xe, \(O^{2-}\), \(N^{3-}\), and \(F^{-}\), consider that ions are generally larger when they have a negative charge because they have gained electrons. Among the ions, \(N^{3-}\) has the greatest negative charge and hence the largest size. Xe is a neutral atom and is larger than an oxygen or fluorine ion. So \(F^{-}\) has the smallest radius.
04

Answer question (c)

Regarding the largest difference between the first and second ionization energies, it's important to understand that the second ionization energy is always larger than the first. However, if removing the second electron means breaking a stable electron configuration, the difference will be huge. Between Al, Si, P, and Cl, Al has an electron configuration ending in \(3p^1\). Removing one electron results in a stable \(3s^2\) configuration, so removing a second one will involve more energy. Thus, Al should have the largest difference between the first and second ionization energies.
05

Answer question (d)

For the greatest ionization energy among C, Si, and Sn, you need to consider the trend that ionization energy decreases from top to bottom and increases from left to right. C is above Si and Sn in the same group, so C has the greatest ionization energy.
06

Answer question (e)

For the largest electron affinity among Na, B, Al, and C, use the trend that electron affinity generally increases from left to right across a period. Thus, C, being furthest to the right, has the highest electron affinity.

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Most popular questions from this chapter

Is it possible for two different atoms to be isoelectronic? two different cations? two different anions? a cation and an anion? Explain.

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

Two elements, \(A\) and \(B\), have the electron configurations shown. $$ \mathrm{A}=[\mathrm{Kr}] 4 s^{2} \quad \mathrm{B}=[\mathrm{Ar}] 3 d^{10} 4 s^{2} 4 p^{5} $$ (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?

Assuming that the seventh period is 32 members long, what should be the atomic number of the noble gas following radon (Rn)? Of the alkali metal following francium (Fr)? What would you expect their approximate atomic masses to be?

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