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) $\mathrm{N}, \mathrm{N}^{3-}, \mathrm{Ne}$
(b) $\mathrm{P}, \mathrm{Si}, \mathrm{Cl}$
(c) $\mathrm{Ga}, \mathrm{O}, \mathrm{Se}$

Key Concepts

Periodic Trends

Periodic trends are systematic variations in the properties of elements that occur as one moves across a period or down a group in the periodic table. These trends, including ionization energy, atomic radius, and electronegativity, arise due to changes in electron configurations and the effective nuclear charge experienced by the electrons.

First Ionization Energy

The first ionization energy is the energy required to remove the most loosely bound electron from a neutral gaseous atom. It is influenced by factors such as atomic size, effective nuclear charge, and electron shielding, with smaller atoms and higher effective nuclear charges generally resulting in higher ionization energies.

Atomic Radius

Atomic radius refers to the size of an atom, which tends to increase down a group and decrease across a period. A larger atomic radius means that the outer electrons are further from the nucleus and are held less tightly, leading to a lower ionization energy.

Effective Nuclear Charge

Effective nuclear charge is the net positive charge experienced by an electron in an atom, taking into account both the total nuclear charge and the shielding provided by inner electrons. A higher effective nuclear charge pulls electrons closer to the nucleus, which increases the ionization energy.

Shielding Effects

Shielding effects describe how inner, non-valence electrons reduce the full positive charge of the nucleus felt by outer electrons. This reduction in effective nuclear charge lowers the ionization energy, especially in atoms with larger electron shells or in ions with added electron repulsion.

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