Trends in Atomic Radius, Ionization Energy, and Electronegativity

Atomic Radius

The radius of an atom is the distance from the nucleus to the outskirts of the electron cloud.

Trend of Atomic Radius Down a Family

As one moves down a family (or column), the atomic radius increases due to the increase in the number of electrons occupying more and higher energy levels and the valence electrons being further away from the nucleus.

Shielding Effect

Inner electrons shield outer electrons from the strong nucleus pull.

Trend of Atomic Radius Across a Period

As one moves across a period (or row) to the right, the atomic radius generally decreases because the extra electrons are being pulled closer to the nucleus without the shielding effect.

Larger Radii

In general, metals (on the left) have larger radii than nonmetals (on the right) of the periodic table.

Ionization Energy (IE)

Ionization energy is the amount of energy required to remove an electron from an atom.

High Ionization Energy

If the ionization energy is high, then it is relatively hard to remove an electron from the atom.

Trend of Ionization Energy Down a Family

As one moves down a family (or column), the ionization energy decreases because it is easier to remove an electron as the valence electrons are further away from the nucleus and the inner electrons shield the outer electrons from the nucleus pull.

Trend of Ionization Energy Across a Period

As one moves across a period (or row) to the right, the ionization energy increases because it is harder to remove an electron as the valence electrons are closer to the nucleus without the shielding effect.

Higher Ionization Energies

In general, nonmetals have higher ionization energies than metals.

Noble Gases and Ionization Energies

Noble gases have particularly high ionization energies because they are the most stable due to completely filled orbitals and smaller radii.

Electronegativity (EN)

Electronegativity is the relative tendency of an atom to attract electrons in chemical bonding.

High Electronegativity

If an atom has a high electronegativity, the atom strongly attracts the electrons.

Trend of Electronegativity Down a Family

As one moves down a family (or column), the electronegativity decreases because the attraction for electrons is weaker due to inner electrons shielding more electrons from the nucleus pull.

Trend of Electronegativity Across a Period

As one moves across a period (or row) to the right, the electronegativity increases because the attraction for electrons is stronger as the extra electron can be pulled by the nucleus without the shielding effect.

Higher Electronegativity

In general, nonmetals have higher electronegativities than metals.

Elements with Highest Electronegativity

The three elements with the highest electronegativities are fluorine, oxygen, and nitrogen.

Electron Attraction in NaCl

In NaCl, Cl attracts the electrons more strongly.

Electron Attraction in CO

In CO, oxygen attracts the electrons more strongly.

Atomic Radius Size

The larger the atomic radius, the larger the size of the atom.

Atom Size Dependence

The size of an atom mainly depends on the space the electrons take up.

Ionization Energy and Atom Size

The higher the ionization energy, the harder it is to remove an electron from an atom.

Electronegativity and Electron Attraction

The higher the electronegativity of an atom, the more the atom attracts electrons.

Connection Between Ionization Energy and Electronegativity

Ionization energy and electronegativity tend to change in the same fashion because both properties are a measure of how much an atom 'wants' electrons.

Lower Electronegativity Elements

In general, metals tend to have lower electronegativities.

Higher Electronegativity Elements

In general, nonmetals tend to have higher electronegativities.

Elements Forming Positive Ions

Metals tend to lose electrons and form positive ions because they have larger radii, lower ionization energy, lower electron affinity, and lower electronegativity values.

Electron Sharing in HCl

When hydrogen and chlorine bond together by sharing electrons, Cl attracts the electrons more strongly.

Hydrogen's Properties

Most periodic tables place Hydrogen in Group I (alkali metals) because it has a s1 electron, but none of its other atomic properties align with these alkali metals.