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Chapter 9: Problem 95

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

Short Answer

Expert verified
The ionization energies of the ions He-, Li-, Be-, B-, C-, N-, O-, and F- increase with increasing atomic number. This is because the increased number of protons in the nucleus of larger atoms more strongly attracts the electrons, thus requiring more energy to remove one.

Step by step solution

01

Understand Ionization Energy

Ionization energy is the energy required to remove an electron from an atom or ion in its gaseous state. It is also important to remember that Ionization energy generally increases moving from left to right across a period (the horizontal rows in the periodic table of elements).
02

Consider Atomic Structures

The atomic number determines the number of protons in an atom, and generally, the number of electrons in a neutral atom. All the ions mentioned are negatively charged, meaning they have one extra electron than their neutral state. Therefore, they all have a similar electronic configuration.
03

Rule of Thumb

As a rule of thumb, ionization energy increases with atomic number because the increasing number of protons attracts the electrons stronger, making it harder to remove one. The ions He-, Li-, Be-, B-, C-, N-, O-, and F- will thus show an ascending trend in ionization energy as their atomic number increase. Exceptions to this rule can occur due to the stability of certain electron configurations.

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Most popular questions from this chapter

All the isoelectronic species illustrated in the text had the electron configurations of noble gases. Can two ions be isoelectronic without having noble-gas electron configurations? Explain.

Complete and balance the following equations. If no reaction occurs, so state. (a) \(\operatorname{Rb}(\mathrm{s})+\mathrm{H}_{2} \mathrm{O}(1) \longrightarrow\) (b) \(\mathrm{I}_{2}(\mathrm{s})+\mathrm{Na}^{+}(\mathrm{aq})+\mathrm{Br}^{-}(\mathrm{aq}) \longrightarrow\) (c) \(\operatorname{SrO}(\mathrm{s})+\mathrm{H}_{2} \mathrm{O}(1) \longrightarrow\) (d) \(\mathrm{SO}_{3}(\mathrm{g})+\mathrm{H}_{2} \mathrm{O}(1) \longrightarrow\)

Refer only to the periodic table on the inside front cover and indicate which of the atoms, \(\mathrm{Bi}, \mathrm{S}, \mathrm{Ba}, \mathrm{As}\) and \(\mathrm{Ca},\) (a) is most metallic; (b) is most nonmetallic; (c) has the intermediate value when the five are arranged in order of increasing first ionization energy.

Neither \(\mathrm{Co}^{2+}\) nor \(\mathrm{Co}^{3+}\) has \(4 \mathrm{s}\) electrons in its electron configuration. How many unpaired electrons would you expect to find in each of these ions? Explain.

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