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) $\mathrm{O}, \mathrm{O}^{2-}, \mathrm{F}$
(b) $\mathrm{C}, \mathrm{Si}, \mathrm{N}$
(c) $\mathrm{Te}, \mathrm{Ru}, \mathrm{Sr}$

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Key Concepts

Atomic Radius

Atomic radius refers to the overall size of an atom. Atoms with larger radii typically have lower ionization energies because their outer electrons are farther from the nucleus and are less strongly attracted, while smaller atoms tend to have higher ionization energies due to a stronger pull from the nucleus.

Effective Nuclear Charge

Effective nuclear charge is the net positive charge experienced by an electron in an atom, taking into account the actual nuclear charge and the screening effect from inner electrons. This concept explains why electrons in different atoms require different amounts of energy to be removed.

Ionization Energy

Ionization energy is the amount of energy required to remove the most loosely held electron from a gaseous atom or ion. This concept is central to understanding how strongly an electron is bound to an atom’s nucleus and is influenced by various factors including nuclear charge and electron configuration.

Periodic Trends

The periodic table is arranged such that trends in properties, including ionization energy, can be observed. Generally, ionization energy increases across a period as the effective nuclear charge increases, and it decreases down a group as atoms become larger and electron shielding increases.

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