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Chapter 7: Problem 84

In Table 7.8 , the bonding atomic radius of neon is listed as $58 \mathrm{pm},\( whereas that for xenon is listed as \)140 \mathrm{pm}$. A classmate of yours states that the value for Xe is more realistic than the one for Ne. Is she correct? If so, what is the basis for her statement?

Short Answer

Expert verified
The classmate's statement is correct. The basis for her statement is the understanding of periodic trends and the relationship between the atomic radius and the position of elements in the periodic table. Xenon (Xe) has a larger atomic radius (140 pm) compared to Neon (Ne) with an atomic radius of 58 pm, due to its lower position in the periodic table, having more electron shells.

Step by step solution

01

The atomic radius is the distance from the nucleus to the outermost electron shell in an atom. As we move across a period in the periodic table, the atomic radius tends to decrease because the number of protons in the nucleus increases, leading to an increased attraction between the electrons and the nucleus. On the other hand, as we move down a group in the periodic table, the atomic radius increases due to additional electron shells. #Step 2: Compare Neon and Xenon in the periodic table#

Neon (Ne) and Xenon (Xe) are both noble gases. Neon is in the second period of the periodic table, while Xenon is in the fifth period. Since Xenon is farther down the group than Neon, it has more electron shells, which results in a larger atomic radius. #Step 3: Evaluate the classmate's statement#
02

The classmate's statement suggests that the bonding atomic radius value for Xenon (Xe) is more realistic than that of Neon (Ne). Considering the periodic trends, Xenon's atomic radius of 140 pm is expected to be larger than Neon's atomic radius of 58 pm, since Xenon is lower in the group and has more electron shells. #Step 4: Determine the correctness of the classmate's statement#

Since the periodic trends justify the larger bonding atomic radius for Xenon compared to Neon, we can conclude that the classmate's statement is correct. The basis for her statement is likely the understanding of periodic trends and the relationship between the atomic radius and the position of elements in the periodic table.

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Most popular questions from this chapter

Detailed calculations show that the value of \(Z_{\text {eff }}\) for the outermost electrons in \(\mathrm{Na}\) and \(\mathrm{K}\) atoms is \(2.51+\) and \(3.49+\), respectively. (a) What value do you estimate for \(Z_{\text {eff }}\) experienced by the outermost electron in both \(\mathrm{Na}\) and \(\mathrm{K}\) by assuming core electrons contribute 1.00 and valence electrons contribute 0.00 to the screening constant? (b) What values do you estimate for \(Z_{\text {eff }}\) using Slater's rules? (c) Which approach gives a more accurate estimate of \(Z_{\text {eff }}\) ? (d) Does either method of approximation account for the gradual increase in \(Z_{\text {eff }}\) that occurs upon moving down a group? (e) Predict \(Z_{\text {eff }}\) for the outermost electrons in the \(\mathrm{Rb}\) atom based on the calculations for \(\mathrm{Na}\) and \(\mathrm{K}\).

In the chemical process called electron transfer, an electron is transferred from one atom or molecule to another. (We will talk about electron transfer extensively in Chapter 20.) A simple electron transfer reaction is $$ \mathrm{A}(g)+\mathrm{A}(g) \longrightarrow \mathrm{A}^{+}(g)+\mathrm{A}^{-}(g) $$ In terms of the ionization energy and electron affinity of atom A, what is the energy change for this reaction? For a representative nonmetal such as chlorine, is this process exothermic? For a representative metal such as sodium, is this process exothermic?

Write an equation for the first electron affinity of helium. Would you predict a positive or a negative energy value for this process? Is it possible to directly measure the first electron affinity of helium?

One way to measure ionization energies is ultraviolet photoelectron spectroscopy (PES), a technique based on the photoelectric effect. exo (Section 6.2 ) In PES, monochromatic light is directed onto a sample, causing electrons to be emitted. The kinetic energy of the emitted electrons is measured. The difference between the energy of the photons and the kinetic energy of the electrons corresponds to the energy needed to remove the electrons (that is, the ionization energy). Suppose that a PES experiment is performed in which mercury vapor is irradiated with ultraviolet light of wavelength \(58.4 \mathrm{nm} .\) (a) What is the energy of a photon of this light, in joules? (b) Write an equation that shows the process corresponding to the first ionization energy of \(\mathrm{Hg}\). (c) The kinetic energy of the emitted electrons is measured to be \(1.72 \times 10^{-18} \mathrm{~J}\). What is the first ionization energy of \(\mathrm{Hg}\), in $\mathrm{kJ} / \mathrm{mol} ?$ (d) Using Figure 7.10 , determine which of the halogen elements has a first ionization energy closest to that of mercury.

Write equations that show the processes that describe the first, second, and third ionization energies of a chlorine atom. Which process would require the least amount of energy?
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