Trends & Properties

One of the greatest strengths of the periodic table is its ability to reveal predictable patterns in the properties of elements. By understanding these periodic trends, you can anticipate an element's behaviour based solely on its position in the table — without needing to memorise the properties of every individual element.

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Key Periodic Trends

The major trends across the periodic table are governed by two opposing forces:

1. Effective nuclear charge (Z_eff) — the net positive charge experienced by valence electrons, after accounting for shielding by inner electrons
2. Electron shielding — the repulsion between inner electrons and valence electrons, which reduces the pull of the nucleus on the outermost electrons

As you move across a period (left to right), Z_eff increases because protons are added but electrons go into the same shell (minimal additional shielding). As you move down a group, new electron shells are added, increasing shielding and moving valence electrons further from the nucleus.

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1. Atomic Radius

Atomic radius is a measure of the size of an atom, typically defined as half the distance between the nuclei of two bonded atoms of the same element.

Trend Across a Period (Left → Right)

Atomic radius decreases. Why?

• Each successive element adds a proton to the nucleus and an electron to the same shell
• The increased nuclear charge pulls all the electrons closer to the nucleus
• Shielding remains roughly constant because the added electrons are in the same energy level

Trend Down a Group (Top → Bottom)

Atomic radius increases. Why?

• Each successive element adds a new electron shell
• Valence electrons are further from the nucleus
• Increased shielding from inner electron shells reduces the effective nuclear charge on the outermost electrons

Period 2 Element	Atomic Radius (pm)	Period 3 Element	Atomic Radius (pm)
Li	152	Na	186
Be	112	Mg	160
B	87	Al	143
C	77	Si	117
N	75	P	110
O	73	S	104
F	71	Cl	99

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2. Ionisation Energy

The first ionisation energy is the energy required to remove one electron from a gaseous atom in its ground state:

X(g) → X⁺(g) + e⁻

Trend Across a Period (Left → Right)

Ionisation energy generally increases. Why?

• Increasing nuclear charge holds electrons more tightly
• The atomic radius decreases, so electrons are closer to the nucleus
• More energy is required to remove an electron

Trend Down a Group (Top → Bottom)

Ionisation energy decreases. Why?

• Valence electrons are further from the nucleus
• Greater shielding from inner electrons
• Less energy is required to remove the outermost electron

Anomalies to Note

The general trend across a period is not perfectly smooth. Two notable exceptions occur:

• Group 2 to Group 13 — ionisation energy drops slightly (e.g. Be → B). The Group 13 element loses an electron from a p-orbital, which is higher in energy and easier to remove than the s-orbital electron of Group 2.
• Group 15 to Group 16 — ionisation energy drops slightly (e.g. N → O). In Group 16, two electrons are paired in one p-orbital, and the electron-electron repulsion makes one easier to remove.

Element	First Ionisation Energy (kJ/mol)
Li	520
Be	900
B	801
C	1086
N	1402
O	1314
F	1681
Ne	2081

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3. Electronegativity

Electronegativity is a measure of the ability of an atom to attract bonding electrons in a covalent bond. The most commonly used scale is the Pauling scale, developed by Linus Pauling.

Trend Across a Period (Left → Right)