Classification Of Elements And Periodicity In Properties

The periodic table is the single most economical piece of information in chemistry — memorise a few trends here and you can predict the behaviour of an element you have never studied. NEET reliably asks 1–2 questions from this chapter, usually on comparing atomic/ionic sizes, ionization enthalpy order, or identifying the block/group from an electronic configuration.

From Early Attempts to the Modern Periodic Law

The urge to arrange elements began when only a few dozen were known. Döbereiner noticed that in certain sets of three chemically similar elements (his "triads"), the atomic weight of the middle one was roughly the average of the other two — for instance Li, Na, K. Newlands then proposed the Law of Octaves: on arranging elements in increasing order of atomic weight, every eighth element resembled the first, much like musical notes. His idea collapsed beyond calcium.

The decisive step came from Mendeleev, who stated that the properties of elements are a periodic function of their atomic weights. His strength was that he trusted chemical behaviour over numerical order:

• He left gaps for undiscovered elements and predicted their properties, naming them eka-boron, eka-aluminium and eka-silicon — later identified as scandium, gallium and germanium, with predictions strikingly close to the measured values.
• He reversed pairs where atomic weights demanded a wrong chemical placement (e.g. placing Te before I).

Two defects remained: the position of hydrogen was ambiguous, and isotopes plus certain weight anomalies could not be explained. Moseley's X-ray work showed that atomic number, not atomic weight, is the fundamental ordering property. This gives the Modern Periodic Law: the physical and chemical properties of elements are a periodic function of their atomic numbers. The real cause of periodicity is the recurrence of similar outer-shell electronic configurations after regular intervals.

For elements with atomic number above 100, IUPAC prescribes temporary names built from numerical roots (nil-, un-, bi-, tri-, quad-, pent-, hex-, sept-, oct-, enn-) with the suffix -ium. Thus element 106 was unnilhexium (Unh) before being named seaborgium.

Architecture of the Long Form Periodic Table

The modern table has 18 vertical groups and 7 horizontal periods. A period corresponds to the filling of one principal shell (with its accessible subshells), so the number of elements in a period equals the number of electrons that can be accommodated by the orbitals being filled: 2, 8, 8, 18, 18, 32, 32.

Elements are grouped into four blocks based on the subshell receiving the last electron:

1. s-block (Groups 1 and 2): outer configuration ns¹⁻². Highly reactive metals, low ionization enthalpies, form ionic compounds, lose electrons readily. Reactivity increases down the group.
2. p-block (Groups 13–18): ns²np¹⁻⁶. Includes metals, non-metals and metalloids. Group 18 is the noble gases with completely filled shells, hence very high ionization enthalpies and near-zero reactivity. Group 17 halogens and Group 16 chalcogens are the most reactive non-metals.
3. d-block (Groups 3–12), the transition elements: (n−1)d¹⁻¹⁰ns¹⁻². They bridge the highly reactive s-block metals and the less reactive p-block, show variable oxidation states, and are typically hard, high-melting metals that form coloured, often paramagnetic compounds. Zn, Cd, Hg have a filled d¹⁰ shell in all their common states and are not regarded as typical transition metals.
4. f-block, the inner transition elements: lanthanoids (Ce–Lu, filling 4f) and actinoids (Th–Lr, filling 5f), placed as two separate rows at the foot of the table. All are metals; actinoids are radioactive.

A practical skill for the exam is reading the position directly off a configuration. For s- and p-block elements, the period = highest principal quantum number of the valence shell; the group = number of valence electrons for the s-block, and 10 + (number of ns + np electrons) for the p-block. For d-block elements, group number = (number of (n−1)d electrons) + (number of ns electrons). Example: [Ar]3d⁵4s¹ → period 4, group 5 + 1 = 6, i.e. chromium.

Metals occupy over three-quarters of the table (left and centre); non-metals crowd the upper right. Along the dividing staircase lie the metalloids — Si, Ge, As, Sb, Te — whose properties are intermediate. Metallic character increases down a group and decreases across a period; non-metallic character does the opposite.

Atomic and Ionic Sizes

Since an atom has no sharp boundary, size is expressed operationally. Covalent radius is half the internuclear distance in a homonuclear diatomic molecule (Cl–Cl distance 198 pm → radius 99 pm). Metallic radius is half the distance between adjacent nuclei in a metallic crystal. Van der Waals radius is used for noble gases and is much larger, which is why the noble gas "radius" appears anomalously big in some data tables.

Two competing factors govern size:

• Across a period (left → right): effective nuclear charge rises because each added electron enters the same shell and screens poorly. The nucleus grips the electron cloud harder, so radius decreases (Na 186 pm → Cl 99 pm).
• Down a group: a new principal shell is added each time, and this outweighs the extra nuclear charge, so radius increases.

For ions:

• A cation is always smaller than its parent atom — electrons are removed, often emptying the outermost shell, and the remaining electrons feel a greater nuclear pull per electron.
• An anion is always larger than its parent atom — added electrons increase inter-electronic repulsion while nuclear charge stays fixed.
• Isoelectronic species have the same number of electrons but different nuclear charge. Size falls as nuclear charge rises: N³⁻ > O²⁻ > F⁻ > Ne > Na⁺ > Mg²⁺ > Al³⁺.

Ionization Enthalpy and Electron Gain Enthalpy

Ionization enthalpy (ΔᵢH) is the energy needed to remove the most loosely held electron from an isolated gaseous atom: M(g) → M⁺(g) + e⁻. It is always endothermic (positive). Successive values increase (ΔᵢH₂ > ΔᵢH₁) because pulling an electron off an already positive species is harder.

Trends and their causes:

• Increases across a period (higher effective nuclear charge, smaller size) and decreases down a group (valence electron farther out and better shielded).
• Two famous breaks in period 2: ΔᵢH of Be > B, and N > O. Boron's electron comes from a 2p orbital that is better shielded and higher in energy than Be's filled 2s²; oxygen loses an electron from a 2p⁴ set, relieving repulsion in a doubly occupied orbital, whereas nitrogen's half-filled 2p³ is extra stable. The same pattern repeats as Mg > Al and P > S in period 3.
• Noble gases have the highest values in their periods; alkali metals the lowest.

Electron gain enthalpy (Δ_egH) accompanies X(g) + e⁻ → X⁻(g). It is negative when energy is released (the usual case for non-metals) and positive when energy must be supplied (noble gases, and Be, Mg, N). Halogens have the most negative values; Group 18 the most positive.

An important anomaly: chlorine, not fluorine, has the most negative electron gain enthalpy. Fluorine's 2p subshell is so compact that the incoming electron suffers strong repulsion from the electrons already present. Similarly O < S and N < P in magnitude of electron affinity. Second electron gain enthalpy is always positive, since an anion repels an approaching electron (O⁻ + e⁻ → O²⁻ needs energy).

Electronegativity, Valence and Periodicity of Chemical Behaviour

Electronegativity is the tendency of an atom in a molecule to attract shared electrons. Unlike electron gain enthalpy it is not a measurable energy but a relative number (Pauling scale: F = 4.0, the highest; Cs and Fr lowest). It increases across a period and decreases down a group, following effective nuclear charge and size. Non-metallic character parallels electronegativity; metallic character is its inverse.

Valence of representative elements equals the number of valence electrons, or 8 minus that number, whichever is smaller — hence the familiar rise and fall 1, 2, 3, 4, 3, 2, 1, 0 across a period. The highest oxidation state with oxygen or fluorine rises across a period (Na₂O, MgO, Al₂O₃, SiO₂, P₄O₁₀, SO₃, Cl₂O₇).

Consequences for compound behaviour:

• Oxides shift from basic on the left (Na₂O, MgO) through amphoteric in the middle (Al₂O₃, which reacts with both acid and alkali) to acidic on the right (P₄O₁₀, SO₃, Cl₂O₇).
• Hydrides: ionic with s-block metals, covalent in the p-block; acidity of the hydrides of a group increases downward (HF < HCl < HBr < HI).
• Anomalous first members: Li, Be, B, C, N, O, F differ markedly from the rest of their groups because they are small, have high electronegativity and lack available d-orbitals. This produces the diagonal relationship — Li resembles Mg, Be resembles Al, B resembles Si.

Common Mistakes and Exam Traps

1. Assuming fluorine has the most negative electron gain enthalpy. It has the highest electronegativity, but chlorine releases more energy on gaining an electron. Do not confuse electronegativity (a relative tendency in a bond) with electron gain enthalpy (a measured energy change for a gaseous atom).
2. Ignoring the Be/B and N/O breaks while writing ionization enthalpy orders. A period-2 ordering question is almost always testing these two exceptions, not the plain left-to-right increase.
3. Ranking isoelectronic ions by group position instead of nuclear charge. For species with equal electron counts, more protons means smaller size, so the anion with the smallest nuclear charge is largest.
4. Mis-assigning group numbers in the p- and d-blocks. A p-block element's group is 10 + valence electrons (so ns²np⁵ is Group 17, not Group 7), and for d-block elements you must add the (n−1)d electrons to the ns electrons. Also remember that noble-gas radii quoted as van der Waals radii cannot be compared directly with covalent radii of other elements.