Using only a periodic table as a guide, arrange each of the...

Using only a periodic table as a guide, arrange each of the following series of species in order of increasing first ionization energy.
(a) $\mathrm{K}, \mathrm{Se}, \mathrm{S}$
(b) $\mathrm{Cr}, \mathrm{As}, \mathrm{O}$
(c) $\mathrm{O}, \mathrm{O}^{-}, \mathrm{F}$

Key Concepts

Electron-Electron Repulsion

Electron-electron repulsion involves the interactions between electrons within the same orbital or shell, which can influence ionization energy. In cases where extra electrons are added, increased repulsion may reduce the energy required to remove one electron, an effect that must be considered when comparing species with different electron configurations.

Atomic Radius

Atomic radius refers to the size of an atom, which generally increases down a group and decreases across a period. A larger atomic radius means that the outer electrons are further from the nucleus and are less strongly attracted, resulting in a lower ionization energy. This factor is important when examining the periodic trends in ionization energy.

Ionization Energy

Ionization energy is the measure of the energy required to remove an electron from the outermost shell of an atom in its gaseous state. It is a fundamental property that reflects how strongly an electron is held by the nucleus, and this concept is crucial for understanding the chemical reactivity and the periodic behavior of elements.

Periodic Trends

Periodic trends describe how properties such as ionization energy change as one moves across periods and down groups in the periodic table. These trends arise due to variations in atomic structure, such as atomic size and effective nuclear charge, and are essential for predicting and rationalizing the behavior of different elements.

Effective Nuclear Charge

Effective nuclear charge is the net positive charge experienced by an electron in an atom, after accounting for the shielding effect of other electrons. A higher effective nuclear charge typically results in a tighter hold on electrons, leading to an increase in ionization energy. This concept is key when comparing elements in different parts of the periodic table.

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