6.5 Periodic Variations in Element Properties

Periodic Trends in Atomic Properties

Periodic trends in atomic size, ionization energy, and electron affinity across periods and down groups of the periodic table

Two competing factors drive nearly every periodic trend: effective nuclear charge (how strongly the nucleus pulls on valence electrons) and electron shielding (how much inner electrons block that pull). Once you understand these two ideas, the trends below all follow logically.

Atomic size (atomic radius)

• Decreases across a period (left to right) because each added proton increases the effective nuclear charge, while electrons are added to the same shell and don't shield each other very well. The nucleus pulls the electron cloud tighter. For example, Li is larger than Ne even though Ne has more electrons.
• Increases down a group because each new period adds another electron shell. The inner shells shield the valence electrons from the nucleus, so the outermost electrons sit farther out. Li is much smaller than Cs for this reason.

Ionization energy (IE)

Ionization energy is the energy required to remove an electron from a gaseous atom in its ground state.

• Increases across a period because rising effective nuclear charge holds electrons more tightly. It takes more energy to pull an electron away from Ne than from Li.
• Decreases down a group because valence electrons are farther from the nucleus and better shielded, making them easier to remove. Li has a higher IE than Cs.

Electron affinity (EA)

Electron affinity is the energy change when an electron is added to a gaseous atom. A more negative value means the atom releases more energy and "wants" that electron more.

• Becomes more negative across a period because higher effective nuclear charge makes the atom more attractive to an incoming electron. Chlorine, for instance, has a very favorable (highly negative) electron affinity.
• Becomes less negative down a group because the added electron would sit farther from the nucleus, so the attraction is weaker. Fluorine's electron affinity is actually less negative than chlorine's due to the very small size of F causing strong electron-electron repulsion, which is a common exception worth remembering.

Electronegativity

Electronegativity measures an atom's ability to attract electrons within a bond. It follows the same pattern as ionization energy: increases across a period and decreases down a group. Fluorine is the most electronegative element.

Atomic vs ionic size comparisons

When atoms gain or lose electrons, their size changes significantly. The key is tracking what happens to the balance between protons and electrons.

Cations (positive ions)

• Smaller than their parent atoms. Removing electrons reduces electron-electron repulsion, and the same number of protons now pulls on fewer electrons, increasing the effective nuclear charge felt by the remaining electrons. For example, $\text{Na}^+$ is much smaller than Na because the entire outer shell is lost.
• The higher the positive charge, the smaller the ion. $\text{Al}^{3+}$ is smaller than $\text{Mg}^{2+}$ because more electrons have been removed relative to the nuclear charge.

Anions (negative ions)

• Larger than their parent atoms. Adding electrons increases electron-electron repulsion while the nuclear charge stays the same, so the electron cloud expands. $\text{Cl}^-$ is noticeably larger than Cl.
• The higher the negative charge, the larger the ion. $\text{O}^{2-}$ is larger than $\text{F}^-$ because it has more excess electrons relative to its nuclear charge.

A useful comparison: in an isoelectronic series (ions with the same number of electrons, like $\text{O}^{2-}$, $\text{F}^-$, $\text{Na}^+$, $\text{Mg}^{2+}$), the ion with the most protons is the smallest because its nucleus pulls harder on the same number of electrons.

Nuclear charge and periodic trends

Effective nuclear charge ($Z_{eff}$) is the net positive charge experienced by a valence electron. You can approximate it as:

$Z_{eff} = Z - S$

where $Z$ is the atomic number (total protons) and $S$ is the shielding from inner electrons.

• $Z_{eff}$ increases across a period. Each step to the right adds a proton, but the new electron enters the same shell and provides little additional shielding. From Li to Ne, $Z_{eff}$ rises steadily.
• $Z_{eff}$ felt by valence electrons does not decrease as dramatically down a group as you might expect, but the increased distance from the nucleus (more shells) outweighs the modest rise in $Z_{eff}$, which is why atoms get larger going down.

$Z_{eff}$ ties all the trends together:

• Atomic size: Higher $Z_{eff}$ pulls electrons closer, shrinking the atom.
• Ionization energy: Higher $Z_{eff}$ means electrons are held more tightly, requiring more energy to remove.
• Electron affinity: Higher $Z_{eff}$ makes the atom more attractive to incoming electrons (more negative EA).
• Ionic size: In an isoelectronic series, the ion with higher $Z_{eff}$ (more protons) is smaller.

Periodic Table Organization and Element Properties

• The periodic table is arranged by increasing atomic number. Elements in the same column (group) share similar valence electron configurations, which is why they have similar chemical properties.
• Metallic character generally increases down a group and decreases across a period. Metals lose electrons easily, so this trend mirrors decreasing ionization energy. Elements in the lower-left corner (like Cs and Fr) are the most metallic, while elements in the upper-right corner (like F and O) are the most nonmetallic.