Chapter 3: Classification of Elements and Periodicity in Properties Long Questions | chemistry Class 11 CBSE

Chapter 3: Classification of Elements and Periodicity in Properties Long Questions | chemistry Class 11 CBSE | ByteNotes.in

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Question 1

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Trace the historical development of the periodic table from Dobereiner to Mendeleev, highlighting the key contributions and limitations of each classification attempt.

Dobereiner (early 1800s) identified triads of elements where the middle element's atomic weight and properties were intermediate between the other two (e.g., Li-Na-K); this was the first attempt at periodic classification but worked for only a few elements and was dismissed as coincidence.
A.E.B. de Chancourtois (1862) arranged elements in order of increasing atomic weight on a cylindrical table to display periodic recurrence of properties, but his work attracted little attention.
Newlands (1865) proposed the Law of Octaves: every eighth element resembled the first; this worked only up to calcium and failed to accommodate newly discovered elements, making it unacceptable to the scientific community.
Mendeleev and Lothar Meyer (1869), working independently, arranged elements in order of increasing atomic weight and observed periodic similarities; Mendeleev published the Periodic Law first and his table was more elaborate, using both physical and chemical properties for classification.
Mendeleev's greatest achievement was predicting the existence, properties (atomic weight, density, oxide formula, chloride formula) of undiscovered elements Eka-Aluminium and Eka-Silicon, later confirmed as gallium and germanium.
A key limitation of Mendeleev's system was that it was based on atomic weight, so some elements had to be placed out of strict weight order (e.g., iodine after tellurium) to maintain chemical similarity in groups; this anomaly was resolved only when Moseley introduced the Modern Periodic Law based on atomic number.

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Question 2

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Compare and contrast the properties, electronic configurations, and chemical behaviour of s-block, p-block, d-block, and f-block elements.

s-Block elements (Groups 1 and 2) have ns1 or ns2 outermost configurations; they are highly reactive metals with low ionization enthalpies, lose electrons readily to form +1 or +2 ions, and are never found free in nature; their compounds (except Li and Be) are predominantly ionic.
p-Block elements (Groups 13-18) have ns2np1 to ns2np6 configurations; they include metals, non-metals, metalloids, and noble gases; halogens (Group 17) and chalcogens (Group 16) have high negative electron gain enthalpies and readily form anions.
d-Block elements (Groups 3-12) have (n-1)d1-10ns0-2 configuration; they are all metals, form coloured ions, show variable oxidation states, paramagnetism, and catalytic activity; they form a bridge between reactive s-block metals and less active p-block elements.
f-Block elements (lanthanoids Ce-Lu, actinoids Th-Lr) have (n-2)f1-14(n-1)d0-1ns2 configuration; they are all metals; within each series properties are quite similar; actinoids are radioactive and have more complicated chemistry than lanthanoids due to more variable oxidation states.
Hydrogen occupies a unique position: with one s-electron it resembles alkali metals but can also gain an electron like halogens (Group 17), so it is placed separately at the top; helium is an s-block element placed with noble gases due to completely filled valence shell.
The chemical reactivity in s-block increases down the group; in p-block non-metallic character increases left to right; d-block elements show less variation in atomic radius across a period than representative elements; many actinoids have been produced only in nanogram quantities.

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Question 3

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Analyse the periodic trends in ionization enthalpy across a period and down a group, explaining the role of effective nuclear charge and shielding effect, including anomalies observed.

Ionization enthalpy generally increases across a period from left to right because electrons are added to the same principal energy level while nuclear charge increases, resulting in greater attraction of electrons to the nucleus and requiring more energy for removal.
The shielding of inner core electrons remains nearly constant across a period (electrons added to same shell do not shield each other effectively), so increasing nuclear charge dominates and ionization enthalpy rises.
Ionization enthalpy decreases down a group because valence electrons are in successively higher principal energy levels, farther from the nucleus; increasing shielding by inner electrons outweighs the increasing nuclear charge.
Anomaly 1: Boron has a lower ionization enthalpy than beryllium despite higher atomic number because boron's outermost electron is in a 2p orbital (more shielded) while beryllium's is in a 2s orbital (better penetration, more tightly held).
Anomaly 2: Oxygen has a lower ionization enthalpy than nitrogen because nitrogen has a stable half-filled 2p configuration (one electron per orbital, no repulsion), while in oxygen two electrons share one 2p orbital causing repulsion that makes electron removal easier.
Noble gases show the highest ionization enthalpies in each period due to completely filled, very stable electron configurations; alkali metals show the lowest, correlating with their high chemical reactivity and tendency to form cations.

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Question 4

Long Form

Discuss electron gain enthalpy and electronegativity as periodic properties, explaining their trends across a period and down a group, including exceptions.

Electron gain enthalpy (ΔegH) is the enthalpy change when a neutral gaseous atom gains an electron to form an anion; it is negative when energy is released (exothermic, like halogens) and positive when energy must be supplied (endothermic, like noble gases).
Across a period, electron gain enthalpy generally becomes more negative because increasing effective nuclear charge makes the atom more attractive to the added electron; at the extreme right, noble gases have positive electron gain enthalpies as added electrons must enter higher energy levels.
Down a group, electron gain enthalpy becomes less negative because the atom's size increases and the added electron is farther from the nucleus; however, the electron gain enthalpies of O and F are less negative than those of S and Cl respectively due to extra repulsion in the compact n=2 orbitals.
Electronegativity is a qualitative measure of an atom's ability to attract shared electrons in a chemical compound; it is not directly measurable but is estimated by scales like Pauling's, which assigns 4.0 to fluorine as the most electronegative element.
Electronegativity increases across a period (decreasing atomic radius means stronger attraction for bonding electrons) and decreases down a group (increasing atomic radius weakens attraction); this trend parallels ionization enthalpy.
Electronegativity is directly related to non-metallic character (high electronegativity means strong tendency to gain electrons) and inversely related to metallic character; thus increasing electronegativity across a period accompanies decreasing metallic character.

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Question 5

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Explain the anomalous properties of second period elements compared to their group members, focusing on the reasons for their different behaviour.

The first element of each group in the s- and p-blocks (Li, Be, B through F) differs significantly from subsequent members due to their exceptionally small atomic and ionic sizes, which is the primary reason for anomalous behaviour.
Second period elements have a high charge-to-radius ratio (high charge density), which leads to greater polarising power; for example, Li and Be form compounds with pronounced covalent character unlike other alkali and alkaline earth metals that form predominantly ionic compounds.
The first members of each group have only four valence orbitals (2s and 2p), restricting their maximum covalency to 4; boron can form only BF4- but cannot form BF63-, while aluminium can expand to form AlF63- using available 3d orbitals.
High electronegativity of second period elements contributes to their unique properties: lithium's high electronegativity compared to other alkali metals means its compounds have more covalent character.
First period p-block elements have a unique ability to form p-pi to p-pi multiple bonds with themselves and with other second period elements (C=C, C≡C, N=N, N≡N, C=O, C=N); heavier elements do not form such stable multiple bonds due to poor overlap of larger orbitals.
The diagonal relationship arises partly from these anomalies: Li resembles Mg, and Be resembles Al in chemical behaviour; this similarity between elements diagonally placed is due to comparable charge-to-radius ratios.

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Question 6

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Analyse how electronic configuration forms the theoretical basis for the periodic classification, with specific reference to periods, groups, and block division.

The Modern Periodic Law, established after Moseley's work (1913), states that properties of elements are periodic functions of their atomic numbers; this periodicity is essentially a consequence of the periodic variation in electronic configurations.
The period number corresponds directly to the highest principal quantum number (n) of elements in that period; as n increases from 1 to 7, each new period begins with filling of a new principal energy level, creating seven periods.
Elements in the same group share the same outer electronic configuration (valence shell type and number of electrons), which determines their chemical behaviour; for example, all Group 1 elements have ns1 and all form +1 ions.
The block classification (s, p, d, f) is based on which type of orbital last receives an electron during the aufbau filling: s-block fills ns, p-block fills np, d-block fills (n-1)d, and f-block fills (n-2)f orbitals.
The number of elements in each block equals twice the orbital capacity of the subshell: s-block has 2 elements per period, p-block has 6, d-block has 10, and f-block has 14 elements per series.
Exceptions like helium (s-block configuration but placed in p-block) and hydrogen (can resemble Group 1 or Group 17) show that while electronic configuration is the primary basis, chemical behaviour is also considered in final placement.

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Question 7

Long Form

Evaluate Mendeleev's Periodic Table in terms of its achievements and limitations compared to the Modern Periodic Table.

Mendeleev's chief achievement was formulating the Periodic Law (properties are periodic functions of atomic weight) and arranging elements in a table with horizontal rows and vertical columns where elements with similar properties appeared in the same group.
His greatest boldness was predicting the existence and detailed properties of undiscovered elements: Eka-Aluminium (gallium) and Eka-Silicon (germanium) including their atomic weights, densities, oxide and chloride formulas, which were experimentally confirmed.
Mendeleev also recognised that some elements did not fit strict weight order; he placed iodine (lower atomic weight) after tellurium in Group VII due to chemical similarity, showing he prioritised chemical behaviour over strict weight sequence.
A significant limitation was the inability to explain why some elements were out of atomic weight sequence (e.g., Ar before K, Co before Ni, Te before I); this anomaly was only resolved by the Modern Periodic Law based on atomic number.
Mendeleev's table could not accommodate isotopes (atoms of same element with different atomic masses), noble gases (discovered later), and had no theoretical basis for the observed periodicity.
The Modern Periodic Table, based on atomic number, places elements in correct sequence, accommodates all known elements including noble gases, provides a theoretical basis through electronic configuration, and uses 18 groups with IUPAC numbering, making it superior to Mendeleev's original table.

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Question 8

Long Form

Discuss the periodic trends in atomic and ionic radii, explaining why noble gas radii are not compared with covalent radii of other elements.

Atomic radius decreases across a period (e.g., Period II: Li=152 pm to F=64 pm) because electrons are added to the same principal quantum level while nuclear charge increases, pulling electrons progressively closer to the nucleus.
Atomic radius increases down a group (e.g., Group I: Li=152 pm, Na=186 pm, K=231 pm, Rb=244 pm, Cs=262 pm) because each successive element adds a new principal energy level and inner electrons shield outer electrons from nuclear attraction.
Ionic radius follows the same general trend as atomic radius; cations are smaller than parent atoms because fewer electrons experience the same nuclear charge; anions are larger due to increased electron-electron repulsion with added electrons.
For isoelectronic species, the species with the greater positive nuclear charge has the smaller radius because the nucleus attracts all electrons more strongly (e.g., among O2-, F-, Na+, Mg2+ with 10 electrons, Mg2+ is smallest and O2- is largest).
Noble gas radii are not compared with covalent radii of other elements because noble gases are monoatomic and do not form covalent bonds; their radii are measured as non-bonded (van der Waals) radii, which are much larger.
The correct comparison for noble gases is with van der Waals radii of other elements, not covalent radii; van der Waals radii represent the effective size of an atom when not bonded, whereas covalent radii represent size within a bond.

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Question 9

Long Form

Describe how the IUPAC systematic nomenclature for elements with atomic numbers above 100 was developed and apply it to name elements with Z=113 and Z=120.

The need for systematic nomenclature arose because new elements with very high atomic numbers are synthesised in minute quantities in only a few laboratories worldwide; scientists sometimes claimed discovery before collecting reliable data, causing disputes (e.g., element 104 was named Rutherfordium by Americans and Kurchatovium by Soviets).
IUPAC recommended that until a new element's discovery is officially proven, a systematic name derived from its atomic number should be used, employing numerical roots: nil(0), un(1), bi(2), tri(3), quad(4), pent(5), hex(6), sept(7), oct(8), enn(9).
The roots are placed in order of the digits of the atomic number and 'ium' is added at the end; the symbol consists of three letters representing the first letter of each numerical root.
For Z=113: digits 1, 1, 3 give roots un+un+tri+ium = Ununtrium, symbol Uut; its official IUPAC name is Nihonium (Nh), reflecting Japan (Nihon) where it was discovered.
For Z=120: digits 1, 2, 0 give roots un+bi+nil+ium = Unbinilium, symbol Ubn; it would belong to Group 2 (alkaline earth metals) based on its expected position in the periodic table.
Once officially recognised, a permanent name reflecting the country or state of discovery, or honouring a notable scientist, is given by vote of IUPAC representatives; elements up to Z=118 have all been officially named.

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Question 10

Long Form

Analyse the relationship between periodic trends in physical properties and the chemical reactivity of elements, illustrating with the formation of oxides.

Chemical reactivity of elements is directly linked to ionization enthalpy and electron gain enthalpy: elements with low ionization enthalpy (alkali metals, extreme left) lose electrons easily to form cations, while elements with high negative electron gain enthalpy (halogens, extreme right) gain electrons to form anions.
Chemical reactivity is therefore highest at the two extremes of a period and lowest at the centre where intermediate ionization enthalpies and electron gain enthalpies make neither electron loss nor gain particularly favourable.
Metallic character (tendency to lose electrons) is highest at the left extreme of a period and increases down a group; non-metallic character (tendency to gain electrons) increases toward the right extreme of a period and decreases down a group.
Elements on the left extreme of a period combine with oxygen to form basic oxides (e.g., Na2O + H2O → 2NaOH); elements on the right form acidic oxides (e.g., Cl2O7 + H2O → 2HClO4).
Elements in the centre form amphoteric oxides (e.g., Al2O3 reacts with both acids and bases) or neutral oxides (e.g., CO, NO, N2O that have neither acidic nor basic properties).
Among transition metals (3d series), the change in atomic radii across a period is much smaller than in representative elements; ionization enthalpies are intermediate between s- and p-blocks, making them less electropositive than Group 1 and 2 metals; in a group, metallic character increases as atomic radius and ionic radius increase.

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Chapter 3: Classification of Elements and Periodicity in Properties | Long Questions

Trace the historical development of the periodic table from Dobereiner to Mendeleev, highlighting the key contributions and limitations of each classification attempt.

Compare and contrast the properties, electronic configurations, and chemical behaviour of s-block, p-block, d-block, and f-block elements.

Analyse the periodic trends in ionization enthalpy across a period and down a group, explaining the role of effective nuclear charge and shielding effect, including anomalies observed.

Discuss electron gain enthalpy and electronegativity as periodic properties, explaining their trends across a period and down a group, including exceptions.

Explain the anomalous properties of second period elements compared to their group members, focusing on the reasons for their different behaviour.

Analyse how electronic configuration forms the theoretical basis for the periodic classification, with specific reference to periods, groups, and block division.

Evaluate Mendeleev's Periodic Table in terms of its achievements and limitations compared to the Modern Periodic Table.

Discuss the periodic trends in atomic and ionic radii, explaining why noble gas radii are not compared with covalent radii of other elements.

Describe how the IUPAC systematic nomenclature for elements with atomic numbers above 100 was developed and apply it to name elements with Z=113 and Z=120.

Analyse the relationship between periodic trends in physical properties and the chemical reactivity of elements, illustrating with the formation of oxides.

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