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

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

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Question 1

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State the Modern Periodic Law and explain why atomic number is a more fundamental property than atomic mass.

The Modern Periodic Law states: the physical and chemical properties of elements are periodic functions of their atomic numbers.
Henry Moseley showed in 1913 that a plot of the square root of X-ray frequency against atomic number gave a straight line, unlike the plot against atomic mass, proving atomic number is more fundamental.
Atomic number equals the nuclear charge (number of protons), which directly determines the electronic configuration and hence all chemical and physical properties of an element.

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Question 2

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State Dobereiner's Law of Triads and give one example with atomic weights.

Dobereiner's Law of Triads states that in a group of three chemically similar elements, the atomic weight of the middle element is approximately the arithmetic mean of the atomic weights of the other two.
Example: In the triad Li (7), Na (23), K (39), the atomic weight of Na = (7+39)/2 = 23, which matches exactly.
The properties of the middle element in each triad are also intermediate between those of the other two members; however, this law worked only for a few elements and was dismissed as coincidence.

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Question 3

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Explain Newlands' Law of Octaves and state its main limitation.

Newlands' Law of Octaves (1865) states that when elements are arranged in increasing order of atomic weights, every eighth element has properties similar to the first, analogous to the octaves of music.
The law seemed to hold true only for elements up to calcium; beyond calcium, the pattern broke down and the law could not accommodate later-discovered elements.
Newlands' idea was not widely accepted at the time, though he was later awarded the Davy Medal in 1887 by the Royal Society, London.

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Question 4

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How did Mendeleev predict the existence of gallium and germanium before their discovery?

Mendeleev left deliberate gaps in his periodic table where no known element fit, calling unknown elements Eka-Aluminium and Eka-Silicon based on their expected positions under aluminium and silicon.
He predicted their properties such as atomic weight (Eka-Aluminium: 68, found 70 for Ga; Eka-Silicon: 72, found 72.6 for Ge), density, formulas of oxides and chlorides.
When gallium (1875) and germanium (1886) were discovered with properties remarkably close to Mendeleev's predictions, it validated his periodic table and made him famous.

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Question 5

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Explain the IUPAC system for naming elements with atomic numbers above 100. Give the name and symbol for element 108.

IUPAC assigns systematic names using numerical roots for each digit of the atomic number: nil(0), un(1), bi(2), tri(3), quad(4), pent(5), hex(6), sept(7), oct(8), enn(9); 'ium' is added at the end.
For element 108: digits are 1, 0, 8 giving roots un + nil + oct + ium = Unniloctium, symbol Uno.
The official IUPAC name given to element 108 is Hassium (Hs); permanent names may reflect the country of discovery or honour a notable scientist, decided by IUPAC vote.

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Question 6

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How does the number of elements in each period of the periodic table relate to quantum numbers?

The period number corresponds to the highest principal quantum number (n) of elements in that period; each successive period involves filling of the next higher principal energy level.
The number of elements in each period equals twice the number of atomic orbitals available at that energy level: period 1 has 2 (1s), periods 2 and 3 have 8 each, periods 4 and 5 have 18 each, and period 6 has 32.
For the 5th period (n=5), the available orbitals are 5s, 4d, and 5p totalling 9 orbitals which can accommodate 18 electrons, hence 18 elements.

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Question 7

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What are the characteristic properties of s-block elements?

s-Block elements belong to Groups 1 (alkali metals, ns1) and Group 2 (alkaline earth metals, ns2); they are all reactive metals with low ionization enthalpies.
They readily lose their outermost electron(s) to form 1+ ions (alkali metals) or 2+ ions (alkaline earth metals); their metallic character and reactivity increase down the group.
Because of high reactivity, s-block elements are never found in pure form in nature; their compounds (except those of lithium and beryllium) are predominantly ionic.

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Question 8

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Why is helium placed in the p-block (Group 18) despite having an s-block configuration?

Helium has the electronic configuration 1s2, which technically makes it an s-block element, but it is placed in Group 18 of the p-block with noble gases.
This placement is justified because helium has a completely filled valence shell and exhibits properties — very low chemical reactivity and stability — that are characteristic of other noble gases.
The classification is based on chemical behaviour and periodic properties rather than strict orbital type, and helium fits best with the noble gas family in terms of reactivity and valence shell stability.

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Question 9

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Give three characteristic properties of d-block (transition) elements.

d-Block elements (Groups 3-12) have the general outer electronic configuration (n-1)d1-10ns0-2; they are all metals and form coloured ions due to incomplete d-subshells.
They exhibit variable valence (oxidation states) because both the ns and (n-1)d electrons can participate in bonding, and they often show paramagnetism.
Transition metals are frequently used as catalysts (e.g., iron in Haber process); however, Zn, Cd, and Hg with configuration (n-1)d10ns2 do not show most typical transition metal properties.

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Question 10

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Explain why atomic radius decreases across a period but increases down a group.

Across a period, electrons are added to the same valence shell while nuclear charge increases progressively; the increased attraction of the nucleus pulls electrons closer, decreasing atomic radius.
Down a group, each successive element has an additional principal energy level, so valence electrons are farther from the nucleus; inner electrons also shield the outer electrons from nuclear attraction.
The combined effect of increasing principal quantum number and shielding by inner electrons outweighs the increasing nuclear charge down a group, resulting in a regular increase in atomic radius.

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Question 11

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Why is the covalent radius of chlorine taken as 99 pm?

The covalent radius of a non-metallic element is defined as half the distance between two atoms of the same element when bound by a single covalent bond in a molecule.
The bond distance in the chlorine molecule (Cl2) is 198 pm; half of this value, 99 pm, is taken as the covalent (atomic) radius of chlorine.
Atomic radii cannot be directly measured since the electron cloud has no sharp boundary; measuring bond distances between atoms in the combined state provides a practical estimate.

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Question 12

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Why is a cation smaller and an anion larger than its parent atom?

A cation is formed by the removal of one or more electrons from the parent atom; with fewer electrons but the same nuclear charge, the effective nuclear charge per electron increases, contracting the electron cloud.
An anion is formed by the addition of one or more electrons; the extra electron(s) increase electron-electron repulsion and reduce the effective nuclear charge experienced per electron, causing expansion.
For example, the atomic radius of sodium is 186 pm while Na+ has an ionic radius of only 95 pm; conversely, fluorine's atomic radius is 64 pm while F- has an ionic radius of 136 pm.

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Question 13

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Define isoelectronic species and give one example from the chapter.

Isoelectronic species are atoms or ions that contain the same number of electrons, regardless of the nature of elements involved; they also have the same number of atoms and same structural arrangement.
Example: O2-, F-, Na+, and Mg2+ are all isoelectronic species as they each contain 10 electrons.
Although they have the same number of electrons, isoelectronic species have different ionic radii because of their different nuclear charges; a higher nuclear charge means a smaller radius, so Mg2+ is smallest and O2- is largest.

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

State the Modern Periodic Law and explain why atomic number is a more fundamental property than atomic mass.

State Dobereiner's Law of Triads and give one example with atomic weights.

Explain Newlands' Law of Octaves and state its main limitation.

How did Mendeleev predict the existence of gallium and germanium before their discovery?

Explain the IUPAC system for naming elements with atomic numbers above 100. Give the name and symbol for element 108.

How does the number of elements in each period of the periodic table relate to quantum numbers?

What are the characteristic properties of s-block elements?

Why is helium placed in the p-block (Group 18) despite having an s-block configuration?

Give three characteristic properties of d-block (transition) elements.

Explain why atomic radius decreases across a period but increases down a group.

Why is the covalent radius of chlorine taken as 99 pm?

Why is a cation smaller and an anion larger than its parent atom?

Define isoelectronic species and give one example from the chapter.

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