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Periodic Table, Chemical Bonding & Atomic Structure Principles GK Questions & Answers

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The systematic organization of chemical elements advanced from early empirical classifications, including Johann Wolfgang Döbereiner's Triads and John Newlands' Law of Octaves, to Dmitri Mendeleev's 1869 periodic system organized by ascending atomic weight. Mendeleev's table gained acceptance through its precise prediction of undiscovered elements, such as eka-aluminium (gallium) and eka-silicon (germanium). In 1913, English physicist Henry Moseley established Modern Periodic Law through characteristic X-ray emission spectroscopy, demonstrating that physical and chemical properties depend fundamentally on atomic number (nuclear charge Z) rather than atomic mass. The modern extended periodic table organizes 118 confirmed elements into seven horizontal periods and eighteen vertical groups classified across four electronic blocks: s, p, d, and f.

Elemental periodicity reflects recurring valence shell electron configurations, determining systematic trends in atomic radius, ionization enthalpy, electron gain enthalpy, and electronegativity. Electronegativity increases across a period from left to right and decreases down a group, reaching its maximum in fluorine. Chemical bonding occurs as atoms attain stable noble gas octets through electron redistribution. Ionic bonds form via complete valence electron transfer between electropositive metals and electronegative non-metals, stabilized by electrostatic lattice energies in crystalline compounds such as sodium chloride. In contrast, covalent bonds involve shared electron pairs, categorized into directional sigma and pi bonds. Coordinate covalent bonds occur when a single donor atom supplies the shared pair, while metallic bonding relies on delocalized electron pools surrounding fixed metal cations.

Chemical bond characteristics dictate macroeconomic industrial manufacturing, semiconductor physics, pharmacology, and metallurgical extraction. For instance, intermolecular hydrogen bonding explains the anomalous physical properties of water, including its elevated boiling point and maximum density at 4 degrees Celsius. In UPSC CSE, State PSC, and SSC CGL examinations, questions routinely address periodic trends across transition metals, inert pair effects in heavier p-block elements, Lewis dot structures, lattice versus hydration enthalpies, and distinctions between electrovalent and covalent compounds. Candidates must demonstrate competence in correlating electronic configurations with chemical reactivity, standard oxidation states, coordinate complexes, and real-world materials science applications such as lithium-ion battery electrochemistry.

Key Concepts & Self-Assessment15 Key Facts

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#1
Modern Periodic Law, formulated by Henry Moseley in 1913, organizes elements by increasing atomic number rather than atomic weight.
#2
The modern long-form periodic table contains 18 vertical groups and 7 horizontal periods, classifying 118 validated chemical elements.
#3
Ionic bonds involve complete electron transfer generating electrostatic attractions, whereas covalent bonds involve electron pair sharing.
#4
Electronegativity increases across a period from left to right and decreases down a group; Fluorine is the most electronegative element.
#5
Subatomic constituents comprise protons and neutrons within the nucleus (discovered by Rutherford and Chadwick) surrounded by orbital electrons.
#6
Mendeleev's periodic table correctly predicted the existence and properties of undiscovered elements, designating them with the Sanskrit prefix Eka (such as Eka-aluminium for Gallium and Eka-silicon for Germanium).
#7
Ionization enthalpy represents the minimum energy required to remove the most loosely bound valence electron from an isolated gaseous atom, peaking at Helium among all elements.
#8
Electron gain enthalpy becomes generally more negative across a period from left to right; Chlorine possesses the most negative electron gain enthalpy in the periodic table, exceeding Fluorine due to lower inter-electronic repulsion.
#9
Coordinate covalent (dative) bonds form when both shared electrons in a bond are contributed exclusively by one of the bonding atoms, as seen in the ammonium ion (NH4+) and hydronium ion (H3O+).
#10
According to the Pauli Exclusion Principle, no two electrons in an atom can have the same set of four quantum numbers (n, l, ml, ms).
#11
Hund's Rule of Maximum Multiplicity dictates that electron pairing in degenerate orbitals cannot occur until each subshell orbital is singly occupied with parallel spins.
#12
Noble gases (Group 18 elements) possess stable closed-shell valence electron configurations (ns2 np6, or 1s2 for Helium), explaining their chemical inertness and zero valency.
#13
The Lanthanide Contraction describes the steady decrease in atomic and ionic radii from Lanthanum (Z=57) to Lutetium (Z=71), caused by poor shielding of nuclear charge by diffuse 4f electrons.
#14
Metallic bonding involves electrostatic attraction between a lattice of positively charged metallic cations and a delocalized sea of mobile conduction electrons, explaining electrical conductivity and ductility.
#15
Oganesson (element 118), synthesized through heavy-ion fusion reactions, completes the 7th period of the modern periodic table and is classified in Group 18.

Subject Specialist Commentary

Analytical perspective & practical exam advice from the Master10 academic board

Educator's Insight
The modern periodic table organizes all 118 known chemical elements according to their atomic number, following Henry Moseley's 1913 discovery. Arranged into eighteen vertical groups and seven horizontal periods, elements in the same group display similar chemical behaviour due to having identical valence electron configurations. Atoms interact to achieve stable octets by either transferring electrons to form ionic bonds or sharing electron pairs to create covalent bonds, following principles of electronegativity and atomic structure.
In competitive exams like SSC CGL, NDA, and State PSCs, examiners frequently test periodic trends and atomic rules. A favourite question trap tests electronegativity versus electron gain enthalpy: Fluorine is the most electronegative element, but Chlorine has the most negative electron gain enthalpy. Remember that noble gases in Group 18 possess stable closed valence shells and zero valency. In your exam revision, make sure you can distinguish Pauli's Exclusion Principle from Hund's Rule of Maximum Multiplicity.

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