Classification of Elements and Periodicity in Properties

🎯 Key Points

• 118 elements arranged in 7 periods (rows) and 18 groups (columns), ordered by increasing atomic number
• Across a period (left→right): atomic radius ↓, ionization energy ↑, electronegativity ↑, metallic character ↓
• Down a group (top→bottom): atomic radius ↑, ionization energy ↓, electronegativity ↓, metallic character ↑
• s-block = Groups 1-2; p-block = Groups 13-18; d-block = Groups 3-12; f-block = lanthanides & actinides
• Anomaly: IE(N) > IE(O) because N's half-filled 2p³ is extra stable
• Diagonal relationship: Li ~ Mg, Be ~ Al (similar charge density despite being in different groups)
• Inert pair effect makes heavy p-block elements (Tl, Pb, Bi) prefer an oxidation state 2 less than the group's maximum

Modern Periodic Law

Properties of elements are periodic functions of their atomic numbers (Mendeleev originally arranged elements by atomic mass; Moseley's X-ray studies later established atomic number as the correct basis). The table has 118 confirmed elements in 7 periods (rows) and 18 groups (columns).

Periodic Trends Overview

The complete modern periodic table (118 elements). Image: Cepheus et al., Public Domain, via Wikimedia Commons.

• Atomic radius: Decreases across a period (nuclear charge increases, pulling electrons closer); increases down a group (extra electron shells dominate)
• Ionic radius: Cations are always smaller than their parent atom (lost outer shell or lost electron density); anions are always larger (gained electron density, same nuclear charge spread over more electrons)
• Ionization energy (IE): Energy needed to remove the outermost electron from a gaseous atom. Increases across a period; decreases down a group. Successive ionization energies (IE₁ < IE₂ < IE₃...) always increase, with a sharp jump once a stable noble-gas-like configuration is reached
• Electronegativity: Tendency of an atom to attract shared electrons in a bond. Increases across a period; decreases down a group. Fluorine is the most electronegative element (3.98 on the Pauling scale)
• Electron gain enthalpy / Electron affinity: Generally becomes more negative (more favourable) across a period, with the most negative values in Group 17 (halogens)
• Metallic / non-metallic character: Metallic character decreases across a period and increases down a group (opposite to electronegativity) - this is why the bottom-left of the table is most metallic and the top-right is most non-metallic
• Acidic/basic nature of oxides: Oxides become more acidic across a period (e.g. Na₂O basic → MgO basic → Al₂O₃ amphoteric → SiO₂ weakly acidic → Cl₂O₇ strongly acidic) and more basic down a group

Blocks of the Periodic Table

• s-block: Groups 1-2 (alkali and alkaline earth metals) - valence electrons fill the s-subshell
• p-block: Groups 13-18 (non-metals, metalloids, noble gases) - valence electrons fill the p-subshell
• d-block: Groups 3-12 (transition metals) - valence electrons fill the d-subshell; show variable oxidation states and form coloured complexes
• f-block: Lanthanides and actinides (inner transition) - valence electrons fill the f-subshell; lanthanides show the "lanthanide contraction" (steady, gradual decrease in atomic/ionic radius across the series due to poor shielding by f-electrons)

Important Groups

• Group 1: Alkali metals (Li, Na, K) - highly reactive, soft, low density, always +1 ion, react vigorously with water
• Group 2: Alkaline earth metals (Be, Mg, Ca) - less reactive than alkali metals, always +2 ion
• Group 17: Halogens (F, Cl, Br, I) - most electronegative group, exist as diatomic molecules, always −1 ion in simple compounds
• Group 18: Noble gases (He, Ne, Ar) - fully filled valence shell, chemically inert (though Xe and Kr can form compounds under special conditions)

Effective Nuclear Charge and Shielding

• Effective nuclear charge (Z_eff): Net positive charge actually felt by valence electrons after accounting for shielding/screening by inner electrons. Approximated by Z_eff = Z − σ (σ = shielding constant)
• Z_eff increases across a period (electrons added to the same outer shell shield each other poorly), causing atomic radius to shrink and ionization energy to rise
• Down a group, extra inner shells shield the nucleus more effectively, so Z_eff felt by the outermost electron stays roughly constant while the radius increases due to the added shells

Anomalous Periodic Trends (Common Exceptions)

• Ionization energy of O is lower than N, because N has a stable half-filled 2p³ configuration that resists electron removal, while O's 2p⁴ has one paired set that's easier to remove
• Ionization energy of Be and Mg are higher than their neighbours B and Al, due to the extra stability of a fully filled s-orbital (ns²)
• Electron gain enthalpy of F is less negative than Cl, because of F's very small size causing strong electron-electron repulsion in the compact 2p shell - the incoming electron is repelled more than expected

Diagonal Relationship

• Elements diagonally placed in the periodic table show similar properties because of comparable charge density (charge/size ratio), even though they belong to different groups
• Li (Group 1) resembles Mg (Group 2): both form a nitride directly with N₂, their carbonates decompose easily on heating, and their fluorides/carbonates are poorly soluble
• Be (Group 2) resembles Al (Group 13): both form amphoteric oxides, covalent chlorides that exist as dimers (BeCl₂, AlCl₃), and are passivated by a protective oxide layer

Periodicity in Valence and Oxidation States

• Maximum covalency/oxidation state of a p-block element generally equals the group number minus 10 (e.g., Group 17 max state = +7, Group 15 max state = +5)
• Inert pair effect: Heavier p-block elements (like Tl, Pb, Bi) prefer an oxidation state two units lower than the group oxidation state, because their ns² valence electron pair is held tightly (poor shielding from intervening d/f electrons) and resists participating in bonding

Common Exam Question

Q: Arrange Na, Mg, Al in order of increasing first ionization energy and explain.
A: Na < Al < Mg. Across the period, IE generally rises (Na < Mg), but Al's IE dips slightly below Mg's because removing Al's single 3p electron is easier than disturbing Mg's stable, fully-filled 3s² configuration.

Genesis of Periodic Classification

The modern table is the result of several earlier attempts to organise the elements:

• Döbereiner's Triads: Grouped elements in threes where the atomic mass of the middle element was roughly the average of the other two (e.g. Li, Na, K). Limited - very few such triads exist.
• Newlands' Law of Octaves: Arranged by increasing atomic mass, every eighth element had similar properties (like musical notes). Worked only up to calcium.
• Mendeleev's Periodic Table: Arranged by increasing atomic mass into periods and groups based on similar properties. His great success was leaving gaps and predicting undiscovered elements (eka-aluminium = Ga, eka-silicon = Ge) with remarkable accuracy. Its defects: no fixed place for isotopes, and some pairs (Ar/K, Co/Ni, Te/I) had to be placed out of mass order to fit their properties.
• Moseley's contribution: X-ray studies showed atomic number (not mass) is the fundamental property, resolving Mendeleev's anomalies and giving the modern periodic law.

IUPAC Nomenclature of Elements (Z > 100)

Elements with atomic number above 100 are given temporary systematic names until officially confirmed, using numerical roots for each digit of the atomic number, with the suffix "-ium".

• Roots: 0 = nil, 1 = un, 2 = bi, 3 = tri, 4 = quad, 5 = pent, 6 = hex, 7 = sept, 8 = oct, 9 = enn.
• Example: element 120 = un-bi-nil-ium = "Unbinilium" (symbol Ubn); element 118 = un-un-oct-ium = "Ununoctium" (now officially Oganesson, Og).
• Once confirmed, the discoverers propose a permanent name (e.g. 101 = Mendelevium, 118 = Oganesson).

Types of Atomic Radius

Because an atom has no sharp boundary, "atomic radius" is defined operationally in different ways depending on bonding:

• Covalent radius: Half the distance between the nuclei of two identical atoms joined by a single covalent bond (e.g. half the Cl–Cl bond length in Cl₂).
• Metallic radius: Half the internuclear distance between two adjacent atoms in a metallic crystal lattice. Slightly larger than covalent radius (metallic bonding is weaker/looser than covalent).
• Van der Waals radius: Half the distance between the nuclei of two non-bonded atoms of adjacent molecules in the solid state. It is the largest of the three, since no bonding pulls the atoms close.
• Order for the same element: van der Waals radius > metallic radius > covalent radius.

Electronegativity Scales

• Pauling scale: The most widely used; based on bond dissociation energies. Fluorine is assigned the maximum value of 3.98 (highest of all elements).
• Mulliken scale: Defines electronegativity as the average of ionization energy and electron affinity: EN = (IE + EA)/2. Mulliken values are about 2.8 times the Pauling values.
• Electronegativity is a property of an atom in a bonded state (not an isolated atom), and it increases with higher oxidation state and greater s-character of the hybrid orbital (sp > sp² > sp³).

Isoelectronic Species

Isoelectronic species are atoms and ions that contain the same number of electrons (and the same electronic configuration).

• Example set (10 electrons each): N³⁻, O²⁻, F⁻, Ne, Na⁺, Mg²⁺, Al³⁺.
• Within an isoelectronic series, size decreases as the nuclear charge (Z) increases, because the same number of electrons is pulled in more tightly: N³⁻ > O²⁻ > F⁻ > Ne > Na⁺ > Mg²⁺ > Al³⁺.
• General rule: the greater the positive charge, the smaller the ion; the greater the negative charge, the larger the ion.

Cause of Periodicity & Types of Elements

Periodicity - the recurrence of similar properties at regular intervals - arises because elements with similar valence-shell (outermost) electronic configurations recur at fixed intervals of atomic number. It is the outer-shell configuration, not the total electron count, that governs chemical behaviour.

On the basis of electronic configuration, elements are grouped into four types:

• Representative (main-group) elements: s-block and p-block (Groups 1, 2 and 13–17); outer s/p shells being filled, showing the most varied chemistry.
• Noble gases: Group 18; completely filled valence shells (ns²np⁶, except He 1s²), chemically very stable.
• Transition elements: d-block (Groups 3–12); the (n−1)d orbitals are being filled.
• Inner-transition elements: f-block (lanthanoids and actinoids); the (n−2)f orbitals are being filled.

Periodic Trends in Chemical Reactivity

• Chemical reactivity is highest at the two extremes of a period and lowest in the centre. On the left, reactivity of metals increases down the group (easier to lose electrons); on the right, reactivity of non-metals decreases down the group (harder to gain electrons).
• The most reactive metal is caesium (bottom-left region) and the most reactive non-metal is fluorine (top-right).
• Reactivity links directly to ionization enthalpy and electron gain enthalpy: low IE → reactive metal; highly negative electron gain enthalpy → reactive non-metal.
• Metallic character (and basic nature of oxides) rises down a group and falls across a period; non-metallic character (and acidic nature of oxides) does the reverse.

🚀 JEE Advanced Edge

Slater's rules (approximate Z_eff calculation): Electrons in the same group as the electron of interest contribute 0.35 each to shielding; electrons in the shell just below contribute 0.85 each; electrons further inside contribute 1.00 each. Z_eff = Z − (sum of these contributions).

Successive ionization energy jumps: A very large jump between IE_n and IE_n+1 reveals the number of valence electrons - e.g. if IE₃ ≫ IE₂ for an element, it has 2 valence electrons (removing the 3rd electron means breaking into a stable, lower noble-gas-like shell).

Lanthanide contraction's knock-on effect: Because of lanthanide contraction, the atomic radii of the 5d transition series elements (right after the lanthanides) are almost identical to their 4d counterparts (e.g. Zr and Hf, Nb and Ta have nearly the same radius), causing these element pairs to have very similar chemical properties and making them notoriously hard to separate.

Worked comparison: Explain why the ionic radius order is N³⁻ > O²⁻ > F⁻ > Na⁺ > Mg²⁺ > Al³⁺ even though all are isoelectronic (10 electrons each). Answer: All have the same electron count, but nuclear charge increases steadily from N (Z=7) to Al (Z=13); higher nuclear charge pulls the same number of electrons in tighter, so radius shrinks steadily as Z increases across this isoelectronic series.

Ionisation Enthalpy

• Ionisation enthalpy is the energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state; it is always endothermic (positive)
• It increases across a period (rising nuclear charge, decreasing size) and decreases down a group (larger size, greater shielding)
• Successive ionisation enthalpies always increase (IE₁ < IE₂ < IE₃) because removing an electron from an increasingly positive ion needs more energy
• Anomalies: boron has a lower IE₁ than beryllium (removal of a less stable 2p electron versus a filled 2s), and oxygen is lower than nitrogen (loss of a paired electron relieving repulsion, versus a stable half-filled 2p³)
• Elements with stable fully-filled and half-filled configurations (noble gases, group 2, group 15) show exceptionally high ionisation enthalpies

Electron Gain Enthalpy

• Electron gain enthalpy is the enthalpy change when an electron is added to an isolated gaseous atom; it is usually negative (energy released) but positive for noble gases and group 2
• Across a period it becomes more negative (greater tendency to accept an electron), and down a group it generally becomes less negative
• Halogens have the most negative electron gain enthalpies because gaining one electron gives them a stable noble-gas configuration
• Notable exception: the electron gain enthalpy of chlorine is more negative than that of fluorine, because the small size of the F atom causes strong electron-electron repulsion in the compact 2p subshell
• Similarly O and N (period 2) have less negative values than S and P (period 3) due to their compact size and higher electron density

Metallic and Non-metallic Character

• Metallic character is the tendency of an atom to lose electrons and form cations; it relates directly to low ionisation enthalpy
• It decreases across a period and increases down a group, so the most metallic elements lie at the bottom-left of the periodic table (caesium, francium)
• Non-metallic character is the tendency to gain electrons and form anions; it increases across a period and decreases down a group, peaking at the top-right (fluorine, oxygen)
• A diagonal band of metalloids (B, Si, Ge, As, Sb, Te) separates metals from non-metals and shows intermediate properties
• Metals form basic oxides, non-metals form acidic oxides, and metalloids or borderline elements often form amphoteric oxides (such as Al₂O₃ and ZnO)

Structure of the Long Form Periodic Table

• The modern long-form table has 7 horizontal periods and 18 vertical groups, arranged in order of increasing atomic number
• The period number equals the principal quantum number (n) of the outermost shell being filled; the number of elements in each period is fixed by the orbitals available (2, 8, 8, 18, 18, 32)
• Elements in the same group have the same outer-shell (valence) electronic configuration, which explains their similar chemical behaviour
• The f-block elements (lanthanoids and actinoids) are placed separately at the bottom to keep the table compact
• The first period holds only 2 elements (only the 1s orbital is available), making it the shortest period
• Group 18 elements (noble gases) have completely filled shells and are placed at the far right