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Chapter 7: Problem 27

Identify each statement as true or false: (a) Cations are larger than their corresponding neutral atoms. (b) \(\mathrm{Li}^{+}\) is smaller than Li. (c) \(\mathrm{Cl}^{-}\) is bigger than I .

Short Answer

Expert verified
(a) False. Cations are smaller than their corresponding neutral atoms, as the loss of electrons results in a stronger attraction towards the nucleus. (b) True. \(\mathrm{Li}^{+}\) is smaller than Li because the ion has lost one electron, resulting in a smaller size. (c) False. The neutral iodine atom has a larger atomic size compared to the \(\mathrm{Cl}^{-}\) anion due to their positions in the Periodic Table.

Step by step solution

01

(Statement a: Cations are larger than their corresponding neutral atoms)

False. Cations are formed when neutral atoms lose one or more electrons, resulting in a positively charged ion. The loss of electrons causes the remaining electrons to experience a stronger attraction towards the nucleus, shrinking the atomic size. Therefore, cations are typically smaller than their corresponding neutral atoms.
02

(Statement b: \(\mathrm{Li}^{+}\) is smaller than Li)

True. Losing an electron in the process of forming a cation generally results in the ion being smaller than the neutral atom. In the case of \(\mathrm{Li}^{+}\), this Li ion has lost one electron, resulting in a smaller size compared to its corresponding neutral state.
03

(Statement c: \(\mathrm{Cl}^{-}\) is bigger than I)

False. This statement is concerning two different elements: chlorine (Cl) and iodine (I). The statement refers to the size of the anion \(\mathrm{Cl}^{-}\) and compares it to the neutral atom of iodine (I). As we move down a group in the Periodic Table, the atomic size increases due to the increased number of electron shells. Since iodine is located below the chlorine in Group 17 (halogens), the neutral iodine atom would have a larger atomic size compared to the neutral chlorine atom. As for the \(\mathrm{Cl}^{-}\) anion, it has gained an extra electron, making it larger than the neutral Cl atom. However, even with the increased size of the anion, the neutral iodine atom would still be larger than the \(\mathrm{Cl}^{-}\) anion, making this statement false.

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Most popular questions from this chapter

Arrange each of the following sets of atoms and ions, in order of increasing size: (a) \(\mathrm{Pb}, \mathrm{Pb}^{2+}, \mathrm{Pb}^{4+}\) (b) \(\mathrm{V}^{3+}, \mathrm{Co}^{2+}, \mathrm{Co}^{3+}\) (c) \(\mathrm{Se}^{2-}, \mathrm{S}^{2-}, \mathrm{Sn}^{2+}\) (d) \(\mathrm{K}^{+}, \mathrm{Rb}^{+}, \mathrm{Br}^{-}\)

Write the electron configurations for the following ions, and determine which have noble-gas configurations: (a) \(\mathrm{Cu}^{2+}\) (b) \(\mathrm{Ca}^{2+},(\mathbf{c}) \mathrm{N}^{3-}\) (d) \(\mathrm{Ru}^{2+}\), (e) \(\mathrm{H}^{-}\).

Mercury in the environment can exist in oxidation states \(0,\) \(+1,\) and \(+2 .\) One major question in environmental chemistry research is how to best measure the oxidation state of mercury in natural systems; this is made more complicated by the fact that mercury can be reduced or oxidized on surfaces differently than it would be if it were free in solution. XPS, X-ray photoelectron spectroscopy, is a technique related to PES (see Exercise 7.111 ), but instead of using ultraviolet light to eject valence electrons, X rays are used to eject core electrons. The energies of the core electrons are different for different oxidation states of the element. In one set of experiments, researchers examined mercury contamination of minerals in water. They measured the XPS signals that corresponded to electrons ejected from mercury's 4 forbitals at \(105 \mathrm{eV}\), from an X-ray source that provided \(1253.6 \mathrm{eV}\) of energy $\left(1 \mathrm{ev}=1.602 \times 10^{-19} \mathrm{~J}\right)$ The oxygen on the mineral surface gave emitted electron energies at $531 \mathrm{eV}\(, corresponding to the \)1 s$ orbital of oxygen. Overall the researchers concluded that oxidation states were +2 for \(\mathrm{Hg}\) and -2 for O. (a) Calculate the wavelength of the \(\mathrm{X}\) rays used in this experiment. (b) Compare the energies of the \(4 f\) electrons in mercury and the \(1 s\) electrons in oxygen from these data to the first ionization energies of mercury and oxygen from the data in this chapter. (c) Write out the ground- state electron configurations for \(\mathrm{Hg}^{2+}\) and \(\mathrm{O}^{2-}\); which electrons are the valence electrons in each case?

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?

(a) Use orbital diagrams to illustrate what happens when an oxygen atom gains two electrons. (b) Why does \(\mathrm{O}^{3-}\) not exist?
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