Periodicity Quick Kit: quick revision, 15 bites

The Periodic Law

Properties repeat with atomic number.

The **Modern Periodic Law** states that the properties of elements are a periodic function of their **atomic number** (not atomic mass). This is the foundation of the periodic table.

Takeaway. Atomic number, not mass, governs periodicity.

Periods and Groups

Rows and columns organize elements.

A **period** is a horizontal row (7 periods). A **group** is a vertical column (18 groups). Elements in the same group have the same **valence electron configuration**.

Takeaway. Groups share chemical behavior due to valence electrons.

Why Atomic Size Decreases Across a Period

More protons, same shell, smaller radius.

Across a period, **effective nuclear charge** increases because protons increase while shielding by inner electrons stays constant. This pulls valence electrons closer, **decreasing atomic radius**.

Takeaway. Increased nuclear charge shrinks atomic size across a period.

Atomic Radius Increases Down a Group

New shells add size.

Down a group, new **energy shells** are added, increasing the principal quantum number. The outermost electrons are farther from the nucleus, so **atomic radius increases**.

Takeaway. Adding shells down a group enlarges atomic radius.

Ionization Enthalpy Trend

Energy to remove an electron.

**Ionization enthalpy** is the energy required to remove the most loosely bound electron from a gaseous atom. It **increases across a period** (due to higher nuclear charge) and **decreases down a group** (due to larger size and shielding).

Takeaway. IE increases across, decreases down.

Electron Gain Enthalpy

Energy change when adding an electron.

**Electron gain enthalpy** is the energy released when an electron is added to a gaseous atom. Generally, it becomes more negative across a period (easier to add electron) and less negative down a group (due to larger size).

Takeaway. Halogens love electrons; noble gases don't.

Electronegativity: The Pulling Power

How strongly an atom attracts shared electrons.

**Electronegativity** is the tendency of an atom to attract a shared pair of electrons in a chemical bond. It **increases across a period** and **decreases down a group**.

Takeaway. Electronegativity follows the same trend as ionization enthalpy.

Metallic vs Non-metallic Character

Metals lose, non-metals gain.

Metals tend to lose electrons (low ionization enthalpy), while non-metals gain electrons (high electron gain enthalpy).

Takeaway. Metallic character is opposite to non-metallic character.

Valence Electrons and Group Number

Outer electrons decide group.

For main-group elements, the **group number** equals the number of valence electrons (for groups 1-2 and 13-18). For example, group 17 elements (halogens) have 7 valence electrons.

Takeaway. Group number = valence electrons for main groups.

Confusing Atomic Radius with Ionic Radius

Cations shrink, anions swell.

A **cation** (positive ion) is smaller than its parent atom because it loses an electron shell. An **anion** (negative ion) is larger because added electrons increase electron-electron repulsion.

Takeaway. Cations are smaller, anions are larger than their atoms.

Trends in One Glance

Memorize the periodic trends.

Takeaway. Use the periodic table to predict trends quickly.

Remembering the Trends

A simple mnemonic.

For properties that increase across a period (like ionization enthalpy and electronegativity), think **'Right is Tight'** – atoms hold electrons more tightly. For down a group, **'Down is Loose'** – electrons are less tightly held.

Takeaway. Right is tight, down is loose.

Mendeleev's Original Table

Before atomic numbers.

Dmitri Mendeleev arranged elements by **atomic mass** and left gaps for undiscovered elements. Later, Henry Moseley showed that atomic number is the fundamental property, leading to the modern periodic law.

Takeaway. Moseley corrected Mendeleev's order.

Predicting Chemical Behavior

Periodic table as a predictor.

Knowing an element's position lets you predict its **reactivity**, **bonding type**, and **oxidation states**. For example, alkali metals (group 1) are highly reactive and form +1 ions, while halogens (group 17) form -1 ions.

Takeaway. Position in the table dictates chemical personality.

Exceptions in Trends

Not all trends are perfect.

Some irregularities occur due to **half-filled and fully-filled orbitals**. For example, ionization enthalpy of nitrogen is higher than oxygen because nitrogen has a stable half-filled p-subshell.

Takeaway. Stability of half-filled orbitals can break trends.