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Chapter 7: Problem 43

Based on their positions in the periodic table, predict which atom of the following pairs will have the smaller first ionization energy: (a) $\mathrm{Br}, \mathrm{Kr} ; \mathbf{( b )} \mathrm{C}, \mathrm{Ca} ;(\mathbf{c}) \mathrm{Li}, \mathrm{Rb} ;\(; (d) \)\mathrm{Pb}, \mathrm{Si} ;$ (e) \(\mathrm{Al}, \mathrm{B}\).

Short Answer

Expert verified
The elements with smaller first ionization energies in each pair are: (a) Br, (b) Ca, (c) Rb, (d) Pb, and (e) Al.

Step by step solution

01

(a) Comparing Br and Kr

Br (Bromine) is in Group 17 (7A) and Period 4, while Kr (Krypton) is in Group 18 (8A) and also Period 4. Since they are in the same period and ionization energy generally increases from left to right within a period, Br will have the smaller first ionization energy compared to Kr.
02

(b) Comparing C and Ca

C (Carbon) is in Group 14 (4A) and Period 2, while Ca (Calcium) is in Group 2 (2A) and Period 4. The ionization energy decreases from top to bottom within a group, so Ca (being lower in the periodic table) will have the smaller first ionization energy compared to C.
03

(c) Comparing Li and Rb

Li (Lithium) is in Group 1 (1A) and Period 2, while Rb (Rubidium) is in Group 1 (1A) and Period 5. Since they are in the same group, we will apply the trend that ionization energy decreases from top to bottom and determine that Rb will have the smaller first ionization energy compared to Li.
04

(d) Comparing Pb and Si

Pb (Lead) is in Group 14 (4A) and Period 6, while Si (Silicon) is in Group 14 (4A) and Period 3. As they are in the same group, ionization energy decreases from top to bottom within a group, so Pb will have the smaller first ionization energy compared to Si.
05

(e) Comparing Al and B

Al (Aluminum) is in Group 13 (3A) and Period 3, while B (Boron) is in Group 13 (3A) and Period 2. Since they are in the same group, we will apply the trend that ionization energy decreases from top to bottom within a group, thereby concluding that Al will have the smaller first ionization energy compared to B.

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Most popular questions from this chapter

An element X reacts with \(\mathrm{F}_{2}(g)\) to form the molecular product shown here. (a) Write a balanced equation for this reaction (do not worry about the phases for \(\mathrm{X}\) and the product). (b) Do you think that \(\mathrm{X}\) is a metal or nonmetal? [Section 7.6\(]\)

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?

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.

(a) If the core electrons were totally effective at screening the valence electrons and the valence electrons provided no screening for each other, what would be the effective nuclear charge acting on the \(3 s\) and \(3 p\) valence electrons in P? (b) Repeat these calculations using Slater's rules. (c) Detailed calculations indicate that the effective nuclear charge is \(5.6+\) for the \(3 s\) electrons and \(4.9+\) for the \(3 p\) electrons. Why are the values for the \(3 s\) and \(3 p\) electrons different? (d) If you remove a single electron from a Patom, which orbital will it come from?

Hydrogen is an unusual element because it behaves in some ways like the alkali metal elements and in other ways like nonmetals. Its properties can be explained in part by its electron configuration and by the values for its ionization energy and electron affinity. (a) Explain why the electron affinity of hydrogen is much closer to the values for the alkali elements than for the halogens. (b) Is the following statement true? "Hydrogen has the smallest bonding atomic radius of any element that forms chemical compounds." If not, correct it. If it is, explain in terms of electron configurations. (c) Explain why the ionization energy of hydrogen is closer to the values for the halogens than for the alkali metals. (d) The hydride ion is \(\mathrm{H}^{-}\). Write out the process corresponding to the first ionization energy of the hydride ion. (e) How does the process in part (d) compare to the process for the electron affinity of a neutral hydrogen atom?
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