Structure of Long Form Table

The modern periodic table, also called the long form or extended periodic table, is the definitive arrangement of all 118 known chemical elements organised by increasing atomic number. Proposed in its current 18-column format largely through the work of Dmitri Mendeleev and later refined by Henry Moseley and Glenn Seaborg, the table reflects the underlying quantum mechanical structure of the atom. Every element's position encodes critical information: its electron configuration, chemical behaviour, metallic character, and the type of atomic orbital being filled by its outermost electrons. The table is divided into horizontal rows called periods (numbered 1 through 7) and vertical columns called groups (numbered 1 through 18 under the IUPAC system). A period corresponds to a principal quantum number n, meaning all elements in Period 3, for example, have their outermost electrons in orbitals where n = 3. Groups collect elements that share the same valence electron configuration and therefore exhibit strikingly similar chemical properties — the alkali metals (Group 1), halogens (Group 17), and noble gases (Group 18) being the most familiar families. Beyond the grid of groups and periods, the table is further subdivided into four blocks — s, p, d, and f — each named after the type of atomic subshell being filled by the differentiating electron. The s-block occupies the far left (Groups 1–2), the p-block the far right (Groups 13–18), the d-block the central transition metals (Groups 3–12), and the f-block forms the two rows of lanthanides and actinides usually displayed below the main table. Understanding this block structure is essential for predicting electron configurations, oxidation states, and periodic trends such as ionisation energy and atomic radius.

Key takeaways

Periods 1–7: Principal Quantum Levels
Groups 1–18: Vertical Families
The s-Block (Groups 1–2)
The p-Block (Groups 13–18)
The d-Block (Groups 3–12): Transition Metals
The f-Block: Lanthanides and Actinides

Periods 1–7: Principal Quantum Levels

Each of the seven periods corresponds to one value of the principal quantum number n (1 through 7). Period 1 contains only 2 elements (H and He) because the first shell holds a maximum of 2 electrons (only the 1s subshell). Period 2 and Period 3 each contain 8 elements, as the s and p subshells together accommodate 2 + 6 = 8 electrons. Periods 4 and 5 each hold 18 elements because the d subshell (10 electrons) becomes available. Periods 6 and 7 each hold 32 elements because the f subshell (14 electrons) also fills. The formula for maximum electrons in a shell is 2n², giving: n=1 → 2, n=2 → 8, n=3 → 18, n=4 → 32.

Groups 1–18: Vertical Families

Groups are numbered 1–18 (IUPAC). Elements in the same group have the same number of valence electrons and similar chemical properties. Group 1 (alkali metals: Li, Na, K, Rb, Cs, Fr) all have one s-electron in their outermost shell (ns¹) and readily form +1 ions. Group 17 (halogens: F, Cl, Br, I, At) have the configuration ns²np⁵ and readily accept one electron to form −1 ions. Group 18 (noble gases: He, Ne, Ar, Kr, Xe, Rn) have full valence shells and are chemically inert. The older Roman numeral system (IA–VIIIA, IB–VIIIB) is still encountered in older textbooks.

The s-Block (Groups 1–2)

The s-block comprises Groups 1 and 2, plus helium (which, despite being in Group 18, fills the 1s orbital). In these elements, the differentiating electron enters an s subshell (l = 0). Their electron configurations end in ns¹ (Group 1) or ns² (Group 2). The s-block elements are highly reactive metals (except H and He), have low ionisation energies, and form ionic compounds with non-metals. Examples: Na (1s²2s²2p⁶3s¹), Mg (1s²2s²2p⁶3s²).

The p-Block (Groups 13–18)

The p-block spans Groups 13–18 (six groups, matching the six electrons the p subshell can hold: l = 1, m_l = −1, 0, +1, each with spin up and down). Elements here have configurations ending in ns²np¹ through ns²np⁶. The p-block includes non-metals (C, N, O, F), metalloids (Si, Ge, As), and metals (Al, Sn, Pb), as well as the noble gases. Chemical behaviour transitions from metallic (left of p-block) to non-metallic (right), illustrating the diagonal trend in the table.

The d-Block (Groups 3–12): Transition Metals

The d-block occupies Groups 3–12. Here the (n−1)d subshell fills (l = 2, holding up to 10 electrons). These are all metals characterised by variable oxidation states, coloured ions, and catalytic activity due to partially filled d orbitals. For example, iron (Fe) has configuration [Ar]3d⁶4s² and forms both Fe²⁺ and Fe³⁺ ions. Chromium (Cr) and copper (Cu) are exceptions, adopting [Ar]3d⁵4s¹ and [Ar]3d¹⁰4s¹ respectively, to achieve extra stability from half-filled or fully filled d subshells.

The f-Block: Lanthanides and Actinides

The f-block (14 elements in each of two rows) is typically shown below the main table to keep it manageable in width. The lanthanides (Ce–Lu, Period 6) fill the 4f subshell; the actinides (Th–Lr, Period 7) fill the 5f subshell (l = 3, m_l = −3 to +3, holding 14 electrons). All actinides are radioactive. The f-block elements show very similar properties to each other because the f electrons are buried inside inner shells and shield nuclear charge poorly, producing the lanthanide contraction — a steady decrease in atomic radius across the lanthanide series that has cascading effects on the chemistry of Period 6 d-block metals.

The Key Equation

N_{\text{max}} = 2n^2

N_{\text{max}} — Maximum electron capacity (dimensionless (count)): The total number of electrons that can fit in all subshells of principal quantum level n.
n — Principal quantum number (dimensionless integer (1, 2, 3, …)): Defines the energy level (shell) of an electron; corresponds directly to the period number for the outermost electron.

Worked Example

Problem. Germanium (Ge, atomic number 32) is placed in Period 4, Group 14 of the periodic table. (a) Write its full electron configuration. (b) Identify which block it belongs to and justify your answer. (c) State the maximum number of electrons in the n = 4 shell.

\text{Ge: } [\text{Ar}]\,3d^{10}4s^24p^2;\quad N_{\text{max}}(n=4) = 2(4)^2 = 32

Answer. Ge configuration: [Ar]3d¹⁰4s²4p²; it is a p-block element (Group 14, Period 4); max electrons in n=4 shell = 32.

Figure

Diagram

Origins · 1 minDmitri Ivanovich Mendeleev and how this was discovered

Dmitri Ivanovich Mendeleev

1834–1907 · Russian

Mendeleev published his periodic law in 1869, arranging the 63 known elements by increasing atomic mass and repeating chemical properties. Crucially, he left gaps for undiscovered elements and correctly predicted the properties of gallium (eka-aluminium), scandium (eka-boron), and germanium (eka-silicon). Although the modern table is ordered by atomic number (Moseley, 1913) rather than mass, Mendeleev's organisational framework — periods and groups — remains the foundation of the long form table used today. Mendeleev reportedly said he saw the complete arrangement of elements in a dream, waking to write it down immediately. Whether apocryphal or not, he famously remarked that he had been thinking about the problem intensely for years before the solution became clear.

Electron Configuration and Aufbau Principle — Block classification depends entirely on which subshell the last electron occupies; understanding Aufbau, Pauli exclusion, and Hund's rule is prerequisite knowledge.
Periodic Trends: Ionisation Energy — Ionisation energy varies predictably across periods and down groups; knowing the table's structure is essential to explain and predict these trends.
Periodic Trends: Atomic Radius — Atomic radius decreases across a period and increases down a group — a direct consequence of how the long form table is organised by period (shell number) and group (nuclear charge).
Properties of Transition Metals (d-Block) — The d-block's position in Groups 3–12 explains why transition metals show variable oxidation states and form coloured complexes — topics that build directly on the block structure concept.
Lanthanide Contraction and the f-Block — The lanthanide contraction is a direct consequence of the poor shielding of 4f electrons and explains anomalous properties of Period 6 d-block metals, linking f-block chemistry to the broader table structure.