AP Chemistry: Periodic Trends To Know

Periodic Trends to Know

Periodic Table

A structured arrangement of elements highlighting trends and properties based on increasing atomic number.

Periodic Trends

Recurring patterns in element properties observed as you move across or down the periodic table.

Effective Nuclear Charge (Zeff)

The net positive charge experienced by an electron, calculated as actual nuclear charge (Z) minus shielding effect (S).

Atomic Radius

The distance from the nucleus to the outermost electron; decreases across a period and increases down a group.

Ionic Radius

The size of an ion; cations are smaller than their neutral atoms, while anions are larger.

Electronegativity

The ability of an atom to attract electrons in a chemical bond; increases across a period and decreases down a group.

Ionization Energy

The energy required to remove an electron from an atom; increases across a period and decreases down a group.

Electron Affinity

The energy change when an atom gains an electron; becomes more negative across a period and less negative down a group.

Groups

Vertical columns on the periodic table where elements share the same number of valence electrons.

Periods

Horizontal rows on the periodic table where elements share the same number of electron shells.

Valence Electrons

The outermost electrons of an atom that determine its chemical reactivity and bonding.

Group 18 (Noble Gases)

Elements with a full octet of valence electrons, making them stable and largely unreactive.

Shielding Effect

The repulsion of valence electrons by inner electrons, reducing the effective nuclear charge felt by the outermost electrons.

Atomic Radius Across a Period

Decreases due to increased nuclear charge pulling electrons closer to the nucleus.

Atomic Radius Down a Group

Increases due to additional electron shells, increasing the distance from the nucleus.

Ionization Energy Across a Period

Increases as nuclear charge increases, making it harder to remove an electron.

Ionization Energy Down a Group

Decreases as electrons are farther from the nucleus and more easily removed.

Electronegativity Across a Period

Increases due to greater nuclear charge attracting bonding electrons more strongly.

Electronegativity Down a Group

Decreases as the nucleus is farther from bonding electrons.

Electron Affinity Across a Period

Becomes more negative as atoms are more eager to gain electrons to achieve stability.

Electron Affinity Down a Group

Becomes less negative as larger atoms have weaker attraction for additional electrons.

Neon vs. Xenon

Both are noble gases with full octets, but neon has 2 electron shells while xenon has 5, affecting atomic size.

Sodium vs. Argon

Both are in Period 3, but argon has a greater nuclear charge, resulting in a smaller atomic radius than sodium.

Be vs. B Exception

Boron has lower ionization energy than beryllium because its 2p electron is less tightly bound than Be’s 2s electrons.

Fluorine vs. Chlorine Electron Affinity

Chlorine has a higher magnitude of electron affinity than fluorine due to less electron repulsion in its outer shell.

Trends in Periods

Properties change across periods due to increasing nuclear charge and decreasing atomic radius.

Trends in Groups

Properties vary due to increasing electron shells and shielding effect down a group.

Key Trend Comparison: Atomic Radius

Decreases across a period but increases down a group.

Key Trend Comparison: Ionization Energy

Increases across a period but decreases down a group.

Key Trend Comparison: Electronegativity

Increases across a period but decreases down a group.