Periodic Table Periods and Groups

Why this topic matters in CDS

The periodic table is the single most reused idea in the chemistry portion of the CDS Science paper. Questions on the position of an element, the names of element families, periodic trends and "the discoverer of X" appear almost every year, and they are usually direct recall or single-step reasoning — exactly the kind of marks you cannot afford to lose.

The NCERT treatment builds gradually: the idea of metals and non-metals in the middle classes, then a formal chapter on the periodic classification of elements in Class 10, and a deeper structural view in Class 11. The CDS paper draws from this entire range, mixing easy historical recall ("who arranged elements in octaves?") with applied reasoning ("which element has the largest atomic radius in period 3?").

Early attempts and Mendeleev's table

Before the modern table, chemists tried several ways to bring order to the growing list of elements.

• Dobereiner's Triads (1817): groups of three elements with similar properties, where the atomic mass of the middle element was roughly the average of the other two — for example Li, Na, K. Only a few such triads could be formed, so the idea was limited.
• Newlands' Law of Octaves (1866): when elements were arranged by increasing atomic mass, every eighth element had properties like the first, much like musical notes. It worked only up to calcium and failed for heavier elements.

Mendeleev's periodic table

In 1869 the Russian chemist Dmitri Mendeleev arranged the 63 elements then known in order of increasing atomic mass, placing elements with similar properties in the same vertical group. His Periodic Law stated that the properties of elements are a periodic function of their atomic masses. His genius lay in two bold moves. First, he left gaps for elements not yet discovered and predicted their properties: his "eka-aluminium", "eka-boron" and "eka-silicon" later turned out to be gallium, scandium and germanium — a striking confirmation that earned the table wide acceptance. Second, he sometimes ignored strict mass order to keep similar elements together, placing cobalt before nickel despite their masses, and tellurium before iodine for the same reason.

Moseley and the modern periodic law

In 1913 Henry Moseley showed by X-ray studies that the fundamental property of an element is its atomic number (Z) — the number of protons — not its atomic mass. This corrected the anomalies in Mendeleev's table.

The modern long form of the periodic table has 7 horizontal periods and 18 vertical groups. Elements are arranged in increasing order of atomic number, and the table directly reflects the filling of electron shells. Because properties now depend on atomic number, isotopes automatically occupy the same place, and hydrogen's awkwardness becomes a minor exception rather than a failure.

This shift from mass to number was profound. Atomic number equals the count of protons, which also fixes the number of electrons in a neutral atom and therefore the electronic arrangement. Since chemical behaviour is decided by electrons, ordering by atomic number puts the deepest cause first and the observed properties follow naturally. The repeating, or periodic, nature of properties is simply the repeating pattern of outer-shell electron configurations as we move along the table — each new period restarts the cycle with a fresh outermost shell.

Periods, groups and what they mean

The table is read in two directions, and each tells you something different about an element.

Periods (rows)

A period is a horizontal row. The period number equals the number of the outermost shell being filled. There are 7 periods: period 1 holds 2 elements, periods 2 and 3 hold 8 each, periods 4 and 5 hold 18 each, and periods 6 and 7 hold 32 each (including the lanthanides and actinides shown separately).

Groups (columns)

A group is a vertical column. Elements in the same group have the same number of valence (outermost) electrons, which is why they show similar chemical behaviour. There are 18 groups in the modern table.

Blocks and element families

The long-form table is divided into blocks based on which sub-shell receives the last electron: the s-block (groups 1−2), p-block (groups 13−18), d-block (groups 3−12, the transition metals) and the f-block (lanthanides and actinides). The named families are favourite CDS targets.

• Group 1 — Alkali metals (Li, Na, K, Rb, Cs, Fr): soft, highly reactive, 1 valence electron, stored under kerosene.
• Group 2 — Alkaline earth metals (Be, Mg, Ca, Sr, Ba, Ra): reactive but less so than group 1, 2 valence electrons.
• Group 17 — Halogens (F, Cl, Br, I, At): very reactive non-metals, 7 valence electrons, need just one electron to complete the octet.
• Group 18 — Noble (inert) gases (He, Ne, Ar, Kr, Xe, Rn): chemically unreactive, complete octet (He has a duplet).
• Groups 3−12 — Transition metals: hard, high melting points, form coloured ions and show variable valency.

Trend 1: Atomic size

Atomic radius is the distance from the nucleus to the outermost shell. Two competing factors control it: the number of shells and the pull of the nucleus (nuclear charge).

• Across a period (left → right): atomic size decreases. Electrons are added to the same shell while protons increase, so the stronger nuclear pull draws the shell inward.
• Down a group (top → bottom): atomic size increases. A new shell is added at each step, and this outweighs the rising nuclear charge.

Trend 2 and 3: Ionisation energy and electronegativity

Ionisation energy (IE) is the energy needed to remove the outermost electron from a neutral gaseous atom. Electronegativity is the tendency of an atom to attract a shared pair of electrons in a bond.

Both depend on how tightly the nucleus holds the outer electrons, so both follow the same pattern — opposite to atomic size.

• Across a period: ionisation energy and electronegativity increase (smaller atom, stronger pull, electron harder to remove or more strongly attracted).
• Down a group: both decrease (larger atom, outer electron farther from the nucleus and shielded by inner shells).

Trend 4: Metallic and non-metallic character

Metallic character is the tendency of an atom to lose electrons and form positive ions. Because losing electrons is easier when ionisation energy is low, metallic character follows atomic size, not the energy trends.

• Across a period: metallic character decreases (elements go from metals on the left to non-metals on the right).
• Down a group: metallic character increases.

The diagonal staircase of metalloids — boron, silicon, germanium, arsenic, antimony, tellurium — separates metals (lower-left) from non-metals (upper-right). Metalloids show properties of both and are vital as semiconductors.

It helps to picture the overall layout: metals fill the left and centre of the table and make up the large majority of elements, non-metals cluster in the upper-right corner, and the thin staircase of metalloids runs between them. As you slide down any group the elements become more metallic, which is why heavier members of even traditionally non-metal groups begin to show metallic shine and rising conductivity. This single mental map answers many identify-the-metal questions without recalling individual elements at all.

Placing an element from its configuration

You can locate any main-group element using its electronic configuration, without memorising the table.

• The number of shells filled gives the period.
• The number of valence electrons gives the position within the period (and the group, for s- and p-block elements).

For example, chlorine has Z = 17 and configuration 2, 8, 7. Three shells are occupied, so it lies in period 3; it has 7 valence electrons, so it sits in group 17 — a halogen. This single skill answers a surprising number of CDS questions.

Worked example: locating an element

From the configuration alone we placed the element, named its group family and predicted its valency — exactly the chain of reasoning CDS rewards.

Previous-year style question

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