The Periodic Table Decoded: Periods, Groups and Trends
Learn how the modern periodic table is built from atomic number and electron arrangement, how to locate any element by its period, group and block, and how atomic radius, ionisation enthalpy and electronegativity change in predictable patterns.
By the PadhoDost Team · 📖 7 min read · Updated 4 August 2026
Part of Class 11 (CBSE) prep🧠 Think of a giant apartment building
Imagine a huge apartment building in your city. Every family is given a flat by one simple rule: their house number. Families on the same floor share the same 'level', and families stacked one above another in the same column turn out to have surprisingly similar lifestyles. The periodic table is exactly this kind of building for the 118 known elements. Each element gets its address from its atomic number (Z), the floors are called periods, and the vertical columns are called groups. Once you know the address, you can predict how an element behaves without memorising each one by heart. Padho, dost, let's learn to read this building.
The Modern Periodic Law
Mendeleev first arranged elements in order of increasing atomic mass, which was brilliant but had a few gaps and reversals. In 1913, Henry Moseley showed that the real basis of an element's identity is its atomic number (the number of protons), not its mass. This gave us the Modern Periodic Law: 'The physical and chemical properties of elements are periodic functions of their atomic numbers.' In plain words, if you line up elements by increasing atomic number, similar properties come back again and again at regular intervals.
The modern (long form) periodic table is the result. It has 7 horizontal rows called periods and 18 vertical columns called groups. Its structure is not random at all: it is built directly from the electronic configurations of atoms, which is why the table is such a powerful memory-saving tool.
Periods, Groups and Blocks
A period number tells you the highest principal quantum number (n) in the atom, that is, how many electron shells the atom uses. Elements in the same group have the same outermost-shell (valence) electronic configuration, and that shared configuration is exactly why they show similar chemistry. Finally, based on which subshell the last (differentiating) electron enters, every element falls into one of four blocks: s, p, d or f.
The four blocks
- ✓s-block: Groups 1 and 2 (last electron enters an s-subshell) — reactive metals like Na and Ca.
- ✓p-block: Groups 13 to 18 (last electron enters a p-subshell) — includes non-metals, metalloids and the noble gases.
- ✓d-block: Groups 3 to 12, the transition elements (last electron enters an (n-1)d subshell), e.g. Fe, Cu.
- ✓f-block: the lanthanoids and actinoids, shown separately at the bottom (last electron enters an (n-2)f subshell).
How to locate any element
- 1Write the full electronic configuration for its atomic number (Z).
- 2Find the highest value of n present — that is the period number.
- 3See which subshell received the last electron — that tells you the block (s, p, d or f).
- 4Apply the matching group rule to get the group number.
- 5Cross-check: the group should match the element's known family (e.g. Group 17 = halogens).
📝 Worked example: place two elements
Element A, Z = 17: configuration 1s² 2s² 2p⁶ 3s² 3p⁵.
Highest n = 3, so Period = 3.
The last electron enters the 3p subshell → p-block.
Outermost s + p electrons = 2 + 5 = 7, so Group = 10 + 7 = 17.
Answer: Chlorine — Period 3, Group 17 (a halogen). Correct.
Element B, Z = 20: configuration 1s² 2s² 2p⁶ 3s² 3p⁶ 4s².
Highest n = 4 → Period 4; last electron in 4s → s-block; valence electrons = 2 → Group 2.
Answer: Calcium — Period 4, Group 2 (an alkaline earth metal). Correct.
Periodic Trends: the patterns that repeat
The key idea behind every trend is effective nuclear charge (Zeff) — the net pull the outer electrons actually feel after the inner electrons 'shield' or screen part of the nucleus. Atomic radius is the size of the atom, measured in picometres (pm). Across a period, electrons keep filling the same shell while protons keep increasing, so Zeff rises and the electron cloud is pulled in tighter: radius decreases (across Period 2 it falls from Li at about 152 pm toward F at about 72 pm). Down a group, each new period adds a whole new shell, so the radius increases even though the nuclear charge is larger.
Ionisation enthalpy is the minimum energy needed to remove the most loosely bound electron from one mole of isolated gaseous atoms in their ground state: M(g) → M⁺(g) + e⁻. It is measured in kJ/mol. A smaller atom holds its outer electron more tightly, so ionisation enthalpy increases across a period and decreases down a group. For example, the first ionisation enthalpy of sodium is about 496 kJ/mol, while the much larger caesium needs only about 376 kJ/mol.
Electronegativity is the tendency of an atom in a molecule to attract the shared pair of bonding electrons towards itself. It is not directly measured, has no units, and is usually quoted on the Pauling scale. Smaller atoms with higher Zeff grip the shared electrons more strongly, so electronegativity increases across a period and decreases down a group. Fluorine (4.0) is the most electronegative element; the heavy Group 1 metals are among the least.
| Property | Across a period (left → right) | Down a group (top → bottom) |
|---|---|---|
| Atomic radius | Decreases | Increases |
| Ionisation enthalpy | Increases | Decreases |
| Electronegativity | Increases | Decreases |
| Metallic character | Decreases | Increases |
⚡ Quick check
An element has the electronic configuration 1s² 2s² 2p⁶ 3s² 3p³. What are its period and group?
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