Periodic Classification of Elements – Long Answer Type Questions
Class 10
Chemistry — Chapter 14
Board: CBSE
Chapter Standard: NCERT Class 10
Topic index
Sections:
History & Development (Q1–Q6)
Q1. Describe Dobereiner’s triads and explain one example in detail. How did this idea contribute to the development of the periodic table?
Döbereiner (early 19th century) grouped elements with similar chemical properties into sets of three called triads. In a triad, the atomic mass of the middle element is approximately the average of the other two. Example: The triad Ca (40), Sr (88), Ba (137) — the atomic mass of Sr (~88) is roughly the average of Ca and Ba. Although triads worked for only a few elements, Döbereiner's idea highlighted patterns among elements and suggested that properties depend on atomic masses — an early step toward systematic classification.
Significance: Triads inspired later scientists to search for broader periodic patterns and supported the idea that elements could be grouped by properties, paving the way for Newlands and Mendeleev.
Significance: Triads inspired later scientists to search for broader periodic patterns and supported the idea that elements could be grouped by properties, paving the way for Newlands and Mendeleev.
Q2. Explain Newlands’ law of octaves and discuss its successes and limitations.
Newlands (1864) observed that when elements were arranged by increasing atomic mass, every eighth element exhibited similar properties — analogous to musical octaves. This worked reasonably for lighter elements (up to calcium) and highlighted recurring patterns.
Limitations: Newlands' rule failed for heavier elements, forced dissimilar elements into the same group, and left no room for undiscovered elements. It also did not account for elements with different properties despite similar positions. Thus, while insightful, the law was insufficient as a general classification scheme.
Limitations: Newlands' rule failed for heavier elements, forced dissimilar elements into the same group, and left no room for undiscovered elements. It also did not account for elements with different properties despite similar positions. Thus, while insightful, the law was insufficient as a general classification scheme.
Q3. Discuss Mendeleev’s periodic table: how he arranged elements, the role of gaps, and how his predictions validated the table.
Mendeleev arranged elements in order of increasing atomic mass but grouped elements with similar chemical properties in vertical columns (groups). Crucially, he left gaps for elements not yet discovered and predicted their properties (for example, he predicted 'eka-aluminium' and 'eka-silicon', later discovered as gallium and germanium) with remarkable accuracy.
Role of gaps: Leaving spaces demonstrated confidence in the periodic pattern and allowed Mendeleev to predict properties of undiscovered elements. When these elements were later found with properties close to his predictions, the credibility of the periodic table rose significantly.
Role of gaps: Leaving spaces demonstrated confidence in the periodic pattern and allowed Mendeleev to predict properties of undiscovered elements. When these elements were later found with properties close to his predictions, the credibility of the periodic table rose significantly.
Q4. Why was there a need to modify Mendeleev’s table and what discovery led to the modern periodic law?
Mendeleev’s table, based on atomic mass, could not explain anomalies like the placement of isotopes and occasional mismatches between atomic mass order and chemical properties (e.g., Te and I). The discovery of the proton and the definition of atomic number (number of protons) by Moseley showed that an element’s chemical properties correlate more directly with its atomic number. This led to the modern periodic law: properties of elements are periodic functions of their atomic numbers, resolving many earlier inconsistencies.
Q5. Compare and contrast the strengths and limitations of Mendeleev's periodic table and the modern periodic table.
Mendeleev's strengths: predicted unknown elements and their properties, grouped elements by chemical properties, and provided a functional classification that matched many observed patterns.
Mendeleev's limitations: relied on atomic mass causing placement anomalies and could not accommodate isotopes properly.
Modern periodic table strengths: arranges elements by atomic number, explains periodicity using electronic configuration, resolves isotope issues, includes noble gases, and provides a systematic long-form table with clear blocks (s, p, d, f). Limitations: while comprehensive, modern table is descriptive; detailed behaviour still requires quantum mechanics for full explanation.
Mendeleev's limitations: relied on atomic mass causing placement anomalies and could not accommodate isotopes properly.
Modern periodic table strengths: arranges elements by atomic number, explains periodicity using electronic configuration, resolves isotope issues, includes noble gases, and provides a systematic long-form table with clear blocks (s, p, d, f). Limitations: while comprehensive, modern table is descriptive; detailed behaviour still requires quantum mechanics for full explanation.
Q6. How did the discovery of noble gases affect the structure of the periodic table?
The discovery of noble gases (late 19th century) added an entirely new group of chemically inert elements with complete valence shells. They did not fit into Mendeleev’s original table, so a new group (Group 18) was added to accommodate them. Their inclusion emphasized the importance of electronic configuration in determining chemical properties and contributed to the shift from atomic mass to atomic number as the organizing principle.
Modern Periodic Table & Structure (Q7–Q13)
Q7. Explain the structure of the long-form (modern) periodic table, including groups, periods and blocks (s, p, d, f). Give examples of elements in each block.
The long-form periodic table organizes elements by increasing atomic number into 18 vertical groups and 7 horizontal periods. Blocks correspond to the type of atomic orbital being filled:
- s-block: Groups 1–2 and helium; outer electrons in s-orbitals (e.g., Na, Mg).
- p-block: Groups 13–18; filling p-orbitals (e.g., C, N, O, F, Ne).
- d-block: Transition metals, Groups 3–12; filling d-orbitals (e.g., Fe, Cu, Zn).
- f-block: Lanthanoids and actinoids; filling f-orbitals, usually placed separately below the main table (e.g., La, Ce, U).
Q8. Define atomic number and mass number. How are they used to identify elements and isotopes?
Atomic number (Z) is the number of protons in the nucleus and uniquely identifies an element. Mass number (A) is the total number of protons and neutrons in the nucleus. Isotopes are atoms of the same element (same Z) with different mass numbers (different numbers of neutrons). For example, carbon-12 and carbon-14 both have Z=6 but A=12 and A=14 respectively. In the periodic table, elements are placed by atomic number, so isotopes occupy the same position.
Q9. Describe why elements in the same group exhibit similar chemical properties, using electronic configuration as the basis.
Elements in the same group have the same number of valence electrons and similar outer electronic configurations. For instance, Group 1 elements have configuration ns¹, making them all prone to lose one electron and thus show similar chemistry (form M⁺ ions, react vigorously with water). Because chemical reactions primarily involve valence electrons, similar valence structures translate into similar chemical behaviour across a group.
Q10. Explain how the periodic table reflects the periodicity of properties using examples of atomic radius and ionization energy.
Periodicity means that many properties show regular trends across periods and groups. Atomic radius decreases across a period due to increasing nuclear charge pulling electrons closer, while ionization energy increases across a period because electrons are held more tightly. Down a group, atomic radius increases (addition of shells) and ionization energy decreases (valence electrons are farther and shielded). Example: In Period 3, Na has a larger atomic radius and lower ionization energy than Cl.
Q11. How are metals, non-metals and metalloids positioned in the periodic table? Provide reasons for their location with reference to electronic structure and properties.
Metals lie on the left and centre (s and d blocks) where atoms have few valence electrons and tend to lose electrons easily, resulting in metallic bonding, conductivity, malleability, and luster. Non-metals are on the upper right (p-block) with more valence electrons and a tendency to gain/share electrons, forming covalent bonds and showing poor conductivity. Metalloids lie along the zig-zag boundary between metals and non-metals and have intermediate properties (e.g., Si, Ge) because their electronic configurations allow both metallic and non-metallic behaviour depending on conditions.
Q12. Discuss the significance of Group 18 (noble gases) in terms of electronic configuration and chemical inertness.
Noble gases have complete valence shells (He: 1s²; Ne: 1s²2s²2p⁶; Ar: …3p⁶), giving them maximum stability and minimal tendency to gain, lose, or share electrons. This makes them largely chemically inert under normal conditions. Their full valence shells explain their extremely low reactivity and absence of stable compounds in ordinary chemistry (though heavier noble gases can form some compounds under extreme conditions).
Q13. What are transition metals? Discuss two distinctive properties of transition metals and explain the reason behind each.
Transition metals are d-block elements (Groups 3–12) characterized by partially filled d-orbitals. Two distinctive properties:
- Variable oxidation states: Because d-electrons can be lost in addition to s-electrons, transition metals exhibit multiple oxidation states (e.g., Fe²⁺/Fe³⁺).
- Coloured compounds: d–d electronic transitions within partially filled d-orbitals absorb visible light, producing coloured ions and complexes (e.g., Cu²⁺ blue, Cr₂O₇²⁻ orange).
Electronic Configuration & Periodicity (Q14–Q19)
Q14. Explain how electronic configuration determines an element’s position in the periodic table and give three examples.
The highest occupied energy level (principal quantum number n) determines the period, while the type of orbital being filled (s, p, d, f) determines the block. Examples:
- Sodium (Na): 1s²2s²2p⁶3s¹ → highest n=3 and s-orbital → Period 3, s-block, Group 1.
- Chlorine (Cl): …3s²3p⁵ → highest n=3 and p-orbital → Period 3, p-block, Group 17.
- Iron (Fe): [Ar]4s²3d⁶ → d-orbital filling → d-block, transition metal (Period 4).
Q15. Describe and explain the trend in atomic radius across a period and down a group. Use diagrams or comparisons in your answer.
Across a period: Atomic radius decreases from left to right due to increasing nuclear charge (more protons) which pulls electrons closer without significantly increasing shielding, so atomic size shrinks (e.g., Na → Mg → Al → Si → P → S → Cl → Ar).
Down a group: Atomic radius increases because each successive element has an additional electron shell, increasing distance between nucleus and valence electrons despite increasing nuclear charge; shielding by inner electrons reduces effective nuclear attraction (e.g., Li → Na → K → Rb). Diagrams: a horizontal arrow (→) for decrease across period and vertical arrow (↓) for increase down group help visualise these trends.
Down a group: Atomic radius increases because each successive element has an additional electron shell, increasing distance between nucleus and valence electrons despite increasing nuclear charge; shielding by inner electrons reduces effective nuclear attraction (e.g., Li → Na → K → Rb). Diagrams: a horizontal arrow (→) for decrease across period and vertical arrow (↓) for increase down group help visualise these trends.
Q16. Define ionization energy and explain why first ionization energy generally increases across a period but decreases down a group.
Ionization energy (IE) is energy required to remove an electron from a gaseous atom or ion. Across a period, IE increases because higher nuclear charge tightly holds electrons, making them harder to remove. Down a group, IE decreases because valence electrons are farther from nucleus and more shielded by inner shells, so they are held less strongly and are easier to remove.
Q17. What is electron affinity? How does it change across the periodic table and why are halogens notable in this property?
Electron affinity is the energy change when an atom gains an electron (often releasing energy). Across a period, electron affinity generally becomes more negative (atoms more readily accept electrons) due to increasing nuclear charge. Halogens have very high (more negative) electron affinities because they need only one electron to achieve a stable noble gas configuration, making them highly eager to gain electrons (e.g., Cl has high electron affinity). Down a group, electron affinity generally decreases (less negative) because added shells reduce nuclear attraction for incoming electrons.
Q18. Explain electronegativity. How does it vary in the periodic table and what practical predictions can it help make?
Electronegativity is a measure of an atom’s ability to attract shared electrons in a chemical bond. It increases across a period (toward fluorine) and decreases down a group. Practical predictions: Electronegativity differences predict bond type—large differences → ionic bonds (e.g., NaCl); small differences → covalent bonds (e.g., H₂), and intermediate differences → polar covalent bonds. It also helps predict acidity of oxides and tendencies to form negative/positive ions.
Q19. Using examples, discuss anomalies in periodic trends and briefly explain why they occur (for instance, irregularities in ionization energies or atomic radii).
Anomalies occur due to electron-electron interactions, subshell configurations, and half-filled or fully filled subshell stability. Example: First ionization energy of oxygen (O) is slightly less than that of nitrogen (N) despite N being to its left. Reason: N has half-filled 2p³ (stable), while O’s 2p⁴ introduces electron pairing, increasing repulsion and slightly lowering IE. Another anomaly: atomic radii across d-block elements don’t decrease smoothly due to poor shielding by d-electrons. These exceptions are explained by finer electronic structure and repulsion effects.
Groups and Periods: Detailed Behaviour (Q20–Q24)
Q20. Describe in detail the trends and chemical behaviour of alkali metals (Group 1). Include reactivity, reaction with water, and common compounds.
Alkali metals (Li, Na, K, Rb, Cs) have ns¹ configuration and are highly reactive, especially with water. Trends: Reactivity increases down the group as the single valence electron becomes easier to lose (larger atomic radius, more shielding). Reaction with water: 2M + 2H₂O → 2MOH + H₂ (vigorous exothermic reactions producing hydroxides and hydrogen). Common compounds: sodium chloride (NaCl), potassium nitrate (KNO₃), sodium hydroxide (NaOH). They form ionic compounds, are soft, have low melting points relative to other metals, and exhibit strong reducing behaviour.
Q21. Discuss Group 17 (halogens): physical properties, chemical reactivity trends, and typical reactions with metals.
Halogens (F, Cl, Br, I) are diatomic non-metals with seven valence electrons (ns² np⁵). Physical properties: F and Cl are gases, Br is a liquid, and I is a solid at room temperature; colours darken down the group. Chemical reactivity: highly reactive non-metals that tend to gain one electron to form X⁻ ions; reactivity decreases down the group due to larger atomic size and reduced ability to attract electrons. Typical reaction with metals: form ionic halides (e.g., 2Na + Cl₂ → 2NaCl). Halogens also participate in displacement reactions where a more reactive halogen displaces a less reactive halide from solution (e.g., Cl₂ + 2KBr → 2KCl + Br₂).
Q22. Explain the general properties and compounds of Group 2 elements (alkaline earth metals). How do they differ from Group 1?
Group 2 elements (Be, Mg, Ca, Sr, Ba) have ns² configuration and commonly exhibit +2 oxidation state. Properties: harder and denser than alkali metals, higher melting points, and less reactive. Reactivity increases down the group but is generally lower than Group 1 because two electrons must be removed. Common compounds include CaCO₃ (limestone), MgO (magnesium oxide), and BaSO₄ (barium sulfate). They form ionic compounds and their hydroxides become more basic and more soluble down the group (e.g., Mg(OH)₂ is less soluble than Ba(OH)₂).
Q23. How does the chemistry of transition metals differ from s- and p-block elements? Discuss catalytic behaviour and complex formation.
Transition metals differ due to partially filled d-orbitals which allow variable oxidation states, formation of coloured compounds, and the ability to form coordination complexes. Catalytic behaviour: transition metals and their compounds often act as catalysts (e.g., Fe in Haber process for ammonia synthesis, Ni in hydrogenation) because they can lend and take electrons during reactions and provide surfaces or intermediate complexes that lower activation energy. Complex formation: transition metals form coordination complexes with ligands (e.g., [Fe(CN)₆]³⁻), stabilising unusual oxidation states and enabling diverse chemistry. These features contrast with s- and p-block elements which have more predictable oxidation states and fewer complexation behaviours.
Q24. Give a detailed account of the variation in metallic character across the periodic table and its practical implications.
Metallic character refers to the tendency to lose electrons and exhibit metallic properties (conductivity, malleability). It increases down a group (atoms more willing to lose electrons as they are further from nucleus) and decreases across a period (left → right) as atoms hold electrons more tightly. Practical implications: Elements on the left-bottom (e.g., Cs, Fr) are highly metallic and used where strong reducing agents are needed; elements on the right-top (e.g., F, O) are non-metallic and used as oxidising agents or in covalent chemistry. Understanding metallic character helps predict bonding types, electrical conductivity, and material selection in engineering (e.g., metals for wiring, ceramics for insulation).
Metals, Non-metals, Transition Elements (Q25–Q28)
Q25. Compare the chemical behaviour of metals and non-metals with emphasis on oxide formation and acid-base nature of oxides.
Metals generally form basic oxides (e.g., Na₂O, CaO) which react with water to form bases (Na₂O + H₂O → 2NaOH). Non-metals form acidic oxides (e.g., CO₂, SO₂) that react with water to form acids (CO₂ + H₂O → H₂CO₃). Some oxides are amphoteric (e.g., Al₂O₃) and can react both with acids and bases. This difference is rooted in electronegativity: metals donate electrons to oxygen forming ionic oxides, while non-metals form covalent oxides that produce acidic behaviour.
Q26. Describe how the periodic table helps predict the types of ions an element will form, giving three detailed examples.
The group number and electronic configuration help predict ion formation. Examples:
- Sodium (Group 1): Na (…3s¹) loses one electron to form Na⁺, achieving noble gas configuration.
- Oxygen (Group 16): O (…2s²2p⁴) tends to gain two electrons to form O²⁻ (oxide ion) completing octet.
- Aluminium (Group 13): Al (…3s²3p¹) commonly loses three electrons to form Al³⁺ due to lower energy requirement relative to gaining five electrons.
Q27. Explain why elements show multiple oxidation states and provide examples where this is important in real-world chemistry.
Multiple oxidation states arise when electrons from different shells (s and d) can be lost or shared. Transition metals commonly show this due to accessible d-orbitals (e.g., Fe²⁺ and Fe³⁺ in biological systems and industrial processes). Importance: Variable oxidation states are crucial in redox chemistry, catalysis (e.g., vanadium oxides in contact process for SO₂ oxidation), and biological electron transport (iron in haem proteins). They also affect compound colours, magnetic properties, and reactivity.
Q28. Discuss the role of periodic trends in selecting elements for technological applications, with specific examples.
Periodic trends guide material selection: high electrical conductivity and ductility (metals like Cu, Ag) make them ideal for wiring and electronics; high melting point and corrosion resistance (transition metals like Ti) suit aerospace components; semiconducting properties of metalloids (Si, Ge) enable their use in electronics and solar cells. Understanding band structure, influenced by atomic and electronic properties, allows engineers to tailor materials for conductivity, reactivity, and mechanical strength.
Applications & Important Examples (Q29–Q30)
Q29. Provide a comprehensive explanation of why fluorine is the most electronegative element and discuss its chemical consequences.
Fluorine has the highest electronegativity due to its small atomic radius and high effective nuclear charge, which strongly attracts electrons in a bond. Chemical consequences: fluorine forms very polar bonds and can oxidise many substances; it forms acids (like HF, though HF is weak due to hydrogen bonding) and highly stable fluoride salts (e.g., CaF₂). Its high electronegativity also makes F₂ extremely reactive and hazardous, reacting with many elements and compounds often explosively.
Q30. Using periodic table reasoning, explain the differences in chemical behaviour between sodium and magnesium despite being in the same period.
Sodium (Na, Group 1, configuration …3s¹) tends to lose one electron to form Na⁺, whereas magnesium (Mg, Group 2, …3s²) typically loses two electrons to form Mg²⁺. Na forms ionic compounds like NaCl and is more reactive with water (2Na + 2H₂O → 2NaOH + H₂) than Mg which reacts slowly due to higher ionization energy. Differences stem from group position and valence electrons: Mg’s additional valence electron increases nuclear attraction for remaining electrons and raises the energy required to remove them compared to Na. Consequently, Mg compounds often have higher lattice energies and melting points than analogous Na compounds.