Classification of Elements and Periodicity in PropertiesNEET Chemistry · Class 11 · NCERT Chapter 3

Introduction

By the mid-1800s, chemists had discovered about 60 elements and needed a system to organise them. The periodic table that emerged from this effort is one of the greatest achievements in science. For NEET, this chapter focuses on understanding why elements are placed where they are and how their properties change in predictable patterns across periods and down groups. Expect 2 questions per year.

The most-tested topics are: periodic trends in ionisation enthalpy, atomic radius, electron gain enthalpy, and electronegativity, plus identifying blocks/groups/periods from electronic configurations.

Historical Development of the Periodic Table

Dobereiner's Triads (1829)

Johann Döbereiner noticed that when three chemically similar elements (a triad) are arranged in order of atomic weight, the atomic weight of the middle element is approximately the average of the other two. Example: Li (6.9), Na (23), K (39). Atomic weight of Na ≈ (6.9 + 39)/2 = 22.95. This was a first hint of periodicity but was limited to three triads.

Newlands' Law of Octaves (1865)

John Newlands arranged the known elements in order of increasing atomic weight and found that every eighth element had properties similar to the first (like musical octaves). This was called the Law of Octaves. Limitation: it worked only for lighter elements up to calcium. Heavier elements did not fit, and gaps were not left for undiscovered elements.

Mendeleev's Periodic Law (1869)

Dmitri Mendeleev stated: The properties of elements are a periodic function of their atomic weights. He arranged all 63 known elements in a table and left gaps for undiscovered elements, predicting their properties. His predictions for eka-aluminium (gallium), eka-boron (scandium), and eka-silicon (germanium) were later confirmed.

Key achievements of Mendeleev's table:

• Systematic arrangement in order of increasing atomic weight with periodic repetition of properties.
• Gaps left for undiscovered elements with predicted properties.
• Anomalies like Ar (40) before K (39) and Co (58.9) before Ni (58.7) were placed by chemical similarity, not strict atomic weight order.

Limitations: could not explain why properties are periodic; placed isotopes in the same position; could not assign correct positions to hydrogen.

Modern Periodic Law

H.G.J. Moseley (1913) determined atomic numbers from X-ray spectra. The Modern Periodic Law states: The physical and chemical properties of elements are periodic functions of their atomic numbers.

Using atomic number (Z) instead of atomic weight resolved all of Mendeleev's anomalies. The position of Ar before K is now correct (Z: Ar=18, K=19). Isotopes have the same Z and correctly occupy one position in the table.

Structure of the Periodic Table

Periods

There are 7 horizontal rows called periods. The period number equals the highest principal quantum number (n) of electrons in the element's ground state. Period 1 has 2 elements (n=1). Period 2 and 3 have 8 each (n=2, 3). Periods 4 and 5 have 18 each. Periods 6 and 7 have 32 each (include lanthanides and actinides).

Groups

There are 18 vertical columns called groups. Elements in the same group have the same valence (outermost) shell electronic configuration and therefore similar chemical properties. Groups 1 and 2 are the s-block; Groups 13–18 are the p-block; Groups 3–12 are the d-block; lanthanides and actinides are the f-block.

Blocks

Helium (He) belongs to the s-block (1s²) but is placed in Group 18 (noble gases) because of its inert properties.

Identifying Group and Period from Electronic Configuration

For s-block and p-block elements: Period = highest principal quantum number (n). Group = number of valence electrons (for Groups 1–2 from ns, Groups 13–18 = 10 + number of np electrons + 2 for the ns electrons).

For d-block elements: Group = (number of electrons in outermost s + (n−1)d). Period = highest n.

Atomic Radii

Select a property to see how it varies across the periodic table. Hover an element to see its relative value.

The atomic radius is the distance from the nucleus to the outermost electrons. Since electrons are in orbitals (probability clouds), the boundary is not sharp. Several types of radii are defined:

Types of Atomic Radii

• Covalent radius: Half the distance between the nuclei of two identical bonded atoms. Used for non-metals. Example: Cl-Cl bond length = 198 pm; covalent radius of Cl = 99 pm.
• Metallic (crystal) radius: Half the distance between the nuclei of two adjacent atoms in a metallic lattice. Always larger than covalent radius.
• Van der Waals radius: Half the distance between nuclei of non-bonded atoms of the same element in adjacent molecules. Largest of all radii. Used for noble gases.
• Ionic radius: Effective size of an ion in an ionic crystal. Cations are smaller than the parent atom; anions are larger.

Periodic Trends in Atomic Radius

Across a period (left to right): Atomic radius decreases. The atomic number increases, so the nuclear charge (Z) increases. More protons pull the same outer shell electrons closer to the nucleus. The number of shielding inner electrons increases only slightly.

Down a group (top to bottom): Atomic radius increases. More energy levels (shells) are added, so the outermost electrons are farther from the nucleus despite increasing nuclear charge. Shielding by inner shells increases too.

Ionic Radii

The ionic radius is the radius of an ion in an ionic crystal, determined from X-ray diffraction.

• Cations are smaller than the parent atom: removing electrons reduces electron-electron repulsion and increases effective nuclear charge on remaining electrons. Example: Na = 186 pm; Na⁺ = 102 pm.
• Anions are larger than the parent atom: adding electrons increases repulsion among electrons. Example: Cl = 99 pm; Cl⁻ = 181 pm.
• Isoelectronic specieshave the same number of electrons. In an isoelectronic series, the ionic radius decreases as nuclear charge increases. Example: O²⁻ (8p, 10e) > F⁻ (9p, 10e) > Na⁺ (11p, 10e) > Mg²⁺ (12p, 10e) > Al³⁺ (13p, 10e).

Ionisation Enthalpy

The first ionisation enthalpy (IE₁) is the minimum energy required to remove the most loosely held electron from one mole of isolated gaseous atoms in the ground state.

The second ionisation enthalpy(IE₂) is the energy to remove an electron from X⁺(g): always IE₂ > IE₁ because you are removing an electron from a more positively charged ion.

Factors Affecting Ionisation Enthalpy

• Nuclear charge (Z): Higher Z pulls electrons more tightly. Higher IE.
• Atomic size: Larger atoms have outer electrons farther from the nucleus and more shielded. Lower IE.
• Shielding (screening) effect: Inner electrons reduce the effective nuclear charge felt by valence electrons. More shielding → lower effective nuclear charge → lower IE.
• Penetration of orbitals:s > p > d > f in penetration. s electrons are more tightly held; f electrons are easiest to remove.
• Half-filled and fully-filled subshells: These are extra stable. Elements with half-filled p or d subshells have higher IE than expected.

Periodic Trends

Across a period: IE generally increases (decreasing atomic size, increasing nuclear charge). Notable exceptions:

• IE₁(Be) > IE₁(B): The 2s orbital is more stable than 2p (deeper penetration).
• IE₁(N) > IE₁(O): N has a half-filled 2p³ (extra stable). O has a paired electron in one 2p orbital, which is easier to remove due to increased repulsion.

Down a group: IE decreases. Atomic size increases; outer electrons are more shielded. They are easier to remove.

Identifying Missing Electrons from Successive IE

A large jump in successive ionisation energies signals the removal of electrons from an inner shell (noble gas core). Example: If the values are IE₁ = 496, IE₂ = 4562, IE₃ = 6912 kJ/mol, the large jump from IE₁ to IE₂ means 1 valence electron. The element is in Group 1 (Na).

Electron Gain Enthalpy (Electron Affinity)

The electron gain enthalpy is the enthalpy change when one mole of isolated gaseous atoms gains an electron:

When energy is released during electron addition, ${\Delta }_{eg}H$ is negative (exothermic). When energy is absorbed, it is positive (endothermic).

Periodic Trends

Across a period: Electron gain enthalpy becomes more negative (more energy released) from left to right, as nuclear charge increases and atomic size decreases.

Down a group: Electron gain enthalpy becomes less negative (less exothermic), because atomic size increases and the added electron enters a farther, more shielded orbital.

Important Anomalies

• F vs Cl: Chlorine has a more negative electron gain enthalpy than fluorine (−349 kJ/mol for Cl vs −328 kJ/mol for F), even though F is smaller and more electronegative. In F, the small atomic size causes high electron-electron repulsion in the compact 2p orbital, reducing the tendency to accept an electron.
• Noble gases (He, Ne, Ar): These have positive (endothermic) electron gain enthalpies because they have fully-filled outer shells and adding an electron would go to a new shell with higher energy.
• Alkaline earth metals (Be, Mg): Positive electron gain enthalpy because the ns² subshell is full; the added electron enters the higher np subshell.
• Nitrogen: Positive electron gain enthalpy because of the extra stable half-filled 2p³ configuration.

Electronegativity

Electronegativity is the tendency of an atom in a chemical bond to attract the shared electron pair toward itself. It is not an absolute property (unlike IE) but a relative concept comparing atoms within a bond.

Scales of Electronegativity

• Pauling scale: Fluorine is assigned 4.0 (highest). Values range from 0.7 (Fr) to 4.0 (F). Most commonly used in NEET.
• Mulliken scale: Electronegativity = (IE₁ + EA) / 2. Based on intrinsic atomic properties.
• Allred-Rochow scale: Based on electrostatic attraction between nucleus and electrons.

Periodic Trends

Across a period: Electronegativity increases from left to right as nuclear charge increases and atomic radius decreases.

Down a group: Electronegativity decreases. Atomic radius increases; the bonded electrons are farther from the nucleus.

Most electronegative: F. Least electronegative non-noble-gas element: Cs or Fr. Metals generally have low electronegativity; non-metals have high values.

Uses of Electronegativity

• Predicts polarity of bonds: larger difference in electronegativity → more ionic character.
• Difference > 1.7: ionic bond. Between 0.5 and 1.7: polar covalent. < 0.5: essentially non-polar covalent.
• Related to oxidation state determination.

Valence and Oxidation State

The valence of an element is its combining capacity (number of bonds it forms). For s and p-block elements, valence = number of valence electrons (for 1–4 electrons) or 8 minus number of valence electrons (for 5–8 electrons). Elements in the same group have the same valence.

Fluorine has only one valence (1) because it has no d orbitals and cannot expand its octet. Chlorine and heavier halogens can show valences 1, 3, 5, and 7 using d orbitals.

Worked NEET Problems

Summary Cheat Sheet