Periodic Trends: Atomic & Ionic Radius, Ionization Energy, and Electronegativity

Vocabulary practice flashcards reviewing periodic trends including atomic radius, ionic radius, ionization energy, and electronegativity based on homework lecture notes.

Atomic Radius (Group Trend)

The distance from the nucleus to the outer electrons, which increases moving down a group due to additional energy levels being added with each new period.

Atomic Radius (Period Trend)

The distance from the nucleus to the outer electrons, which decreases across a period because the number of protons increases while energy levels stay the same, increasing the nuclear charge and pulling electrons in closer.

Cation Radius

An ionic radius that is smaller than the neutral atom from which it formed; when metals lose valence electrons, the outer energy level is lost and the remaining protons pull electrons in closer.

Anion Radius

An ionic radius that is larger than the neutral atom from which it formed; when nonmetals gain electrons, having more electrons than protons increases repulsive forces.

Smallest Atomic Radius (Li, P, O, S)

$O$ (oxygen), because atomic radius decreases as you move across a period and increases down a group.

Largest Atomic Radius (Cl, S, Br, As)

$As$ (arsenic), because it possesses more energy levels than period 3 elements and is further left than bromine.

Barium vs. Barium Ion Radius ($Ba$ vs. $Ba^{2+}$)

$Ba$ is larger than $Ba^{2+}$ because neutral metal atoms are larger than the cations formed by losing electrons.

Bromine vs. Bromide Ion Radius ($Br$ vs. $Br^{1-}$)

$Br^{1-}$ is larger than $Br$ because nonmetal anions gain electrons, which increases electron-electron repulsive forces.

Aluminum vs. Aluminum Ion Radius ($Al$ vs. $Al^{3+}$)

$Al$ is larger than $Al^{3+}$ because forming a cation removes the outer valence energy level and pulls remaining electrons closer.

Oxygen vs. Oxide Ion Radius ($O$ vs. $O^{2-}$)

$O^{2-}$ is larger than $O$ because adding two electrons causes greater electron repulsion compared to the neutral atom.

Ionization Energy (Period Trend)

The energy required to remove an electron, which increases as you move further to the right across a period (e.g., $P$ has a higher ionization energy than $Al$ or $Si$) because more protons pull electrons closer into a smaller atomic radius.

Extremes of V, Cr, and W

$Cr$ is both the smallest atom and possesses the highest ionization energy among vanadium ($V$), chromium ($Cr$), and tungsten ($W$).

Extremes of N, P, and As

$N$ is both the smallest atom and possesses the highest ionization energy among nitrogen ($N$), phosphorus ($P$), and arsenic ($As$).

Electronegativity

The ability of an atom to attract electrons in a compound.

Electronegativity (Period Trend)

A trend that increases across a period as atomic radius decreases, driven by an increase in protons that makes the nuclear charge stronger.

Group 1 Downward Trend (IE & Electronegativity)

The decrease in ionization energy and electronegativity observed when moving down Group 1 elements, caused by increasing radius and increasing shielding.

Halogen Electronegativity Sequence

$I, Br, Cl, F$, which represents the group 17 elements listed in order of increasing electronegativity.