Periodic Table & Trends

📈 Trends Explained with Quantum Theory

Atomic Radius

The atomic radius is the distance from the nucleus to the outermost electron.

Across a period (→): Atomic radius decreases because electrons are added to the same energy level while the nucleus gains more protons, pulling electrons closer.

Down a group (↓): Atomic radius increases as more energy levels (shells) are added, making the atom larger.

Ionization Energy

Ionization energy is the energy required to remove an electron from a neutral atom.

Across a period (→): Ionization energy increases because the attraction between electrons and the nucleus gets stronger.

Down a group (↓): Ionization energy decreases since outer electrons are farther from the nucleus and are more easily removed.

Electronegativity

Electronegativity is the ability of an atom to attract electrons in a chemical bond.

Across a period (→): Electronegativity increases as atoms become better at attracting electrons due to higher nuclear charge.

Down a group (↓): Electronegativity decreases because the bonding electrons are farther from the nucleus.

Electron Affinity

Electron affinity is the energy change when an atom gains an electron.

Across a period (→): Electron affinity becomes more negative, meaning atoms release more energy when gaining electrons.

Down a group (↓): Electron affinity becomes less negative, as the added electron is farther from the nucleus.

🧠 Explanation Factors

Effective Nuclear Charge (Zₑff)

This is the net positive charge experienced by an electron. It increases across a period because more protons are added, while inner electrons stay the same. A higher Zₑff pulls electrons closer to the nucleus, affecting atomic size and energy trends.

Shielding Effect

As more electron shells are added down a group, inner electrons block the attraction between the nucleus and the outer electrons. This weakens the nucleus’s pull on outer electrons, making it easier for atoms to lose electrons and harder to attract new ones.

Orbital Penetration

Electrons in orbitals closer to the nucleus (like s-orbitals) experience stronger attraction due to better penetration. Electrons in outer orbitals (like p, d, or f) are less tightly held, which influences ionization energy and electron affinity.