11th HSC Chemistry
Periodic Classification
of Elements
0 / 30 answered
Section A — History & Early Classification
1
Who proposed the Law of Triads, grouping elements in sets of three based on atomic mass?
Döbereiner (1829) observed that when three chemically similar elements were arranged in order of increasing atomic mass, the atomic mass of the middle element was approximately the arithmetic mean of the other two — called Döbereiner's Law of Triads.
2
Newlands' Law of Octaves states that every eighth element has properties similar to the first. Up to which element was this law found to be valid?
Calcium (Ca) — Newlands' Law of Octaves worked only up to calcium; beyond that the pattern broke down because noble gases were unknown and heavier elements didn't fit the octave scheme.
3
Mendeleev's Periodic Law states that properties of elements are a periodic function of their:
Mendeleev (1869) arranged elements in order of increasing atomic mass and stated that properties repeat periodically — the original periodic law. Modern periodic law uses atomic number (Moseley).
4
The name given by Mendeleev to the element predicted to fill the gap below aluminium in his table was:
Mendeleev predicted Eka-aluminium (later discovered as Gallium, 1875) below aluminium, and Eka-silicon (later Germanium). 'Eka' means 'one beyond' in Sanskrit.
5
The Modern Periodic Law is based on atomic number. Who experimentally established the concept of atomic number using X-ray spectra?
Henry Moseley (1913) showed that the frequency of X-rays emitted by elements is directly proportional to the square of the atomic number (√ν ∝ Z), establishing atomic number as the fundamental property of elements.
Section B — Modern Periodic Table Structure
6
How many periods and groups are there in the Modern Long Form of the Periodic Table?
The Modern Periodic Table has 7 periods (horizontal rows) and 18 groups (vertical columns), accommodating all 118 known elements.
7
Which period of the periodic table contains the largest number of elements?
Period 6 contains 32 elements (Z = 55 to 86), the most of any period, including the lanthanide series (4f elements).
8
Elements in which the last electron enters the d-subshell are called:
d-block elements (transition metals) have their differentiating electron entering the d-subshell. They belong to groups 3–12 of the periodic table.
9
The elements of group 1 (alkali metals) and group 2 (alkaline earth metals) belong to which block?
Groups 1 and 2 are s-block elements; their valence electrons occupy the s-subshell. Hydrogen and Helium are also technically s-block elements.
10
What is the IUPAC name for Group 17 elements?
Halogens — Group 17 elements (F, Cl, Br, I, At) are called halogens (Greek: 'salt-forming'). Group 16 = Chalcogens; Group 15 = Pnictogens; Group 18 = Noble gases.
Section C — Atomic Radius & Ionic Radius
11
Across a period (left to right), atomic radius generally:
Across a period, nuclear charge increases while electrons are added to the same shell, so effective nuclear charge increases and atomic radius decreases.
12
Which of the following has the largest atomic radius?
Na (sodium, Z=11) has the largest radius among the four, as all are in period 3 and atomic radius decreases left to right: Na > Mg > Al > Si.
13
Cationic radius compared to the parent atom is:
A cation is formed by loss of electrons. Fewer electrons means less electron–electron repulsion and a higher effective nuclear charge per electron, making the cation smaller than the parent atom.
14
Which of the following is an isoelectronic series arranged in increasing order of ionic radius?
Na⁺, Mg²⁺, Al³⁺ all have 10 electrons (isoelectronic). Higher nuclear charge → smaller radius. Z: Al(13) > Mg(12) > Na(11), so radius: Al³⁺ < Mg²⁺ < Na⁺.
Section D — Ionisation Enthalpy
15
Which of the following has the highest first ionisation enthalpy among period 2 elements?
Neon (Ne) has the highest first ionisation enthalpy in period 2 due to its completely filled, highly stable noble gas configuration. Ne > F > N > O (note N > O due to half-filled stability).
16
The ionisation enthalpy of nitrogen is higher than that of oxygen because nitrogen has:
Nitrogen (2p³) has a half-filled 2p subshell, which is exceptionally stable due to symmetrical electron distribution and exchange energy. This makes it harder to remove an electron from N than from O.
17
Down a group, ionisation enthalpy:
Down a group, atomic size increases and more inner shells shield the valence electrons from the nucleus, reducing effective nuclear charge on them. Hence ionisation enthalpy decreases.
Section E — Electron Gain Enthalpy & Electronegativity
18
Which element has the most negative (most exothermic) electron gain enthalpy?
Chlorine has the highest electron affinity (most negative ΔegH) despite F being more electronegative. F has a small, compact atom where electron–electron repulsion in the 2p subshell makes addition of an electron less favourable than in Cl's larger 3p subshell.
19
Electronegativity across a period (left to right) generally:
Electronegativity increases across a period as nuclear charge increases and atomic size decreases, making atoms more capable of attracting shared electrons. Fluorine is the most electronegative element (Pauling scale: 4.0).
20
The electronegativity scale most commonly used in Indian HSC chemistry textbooks is:
The Pauling scale (Linus Pauling, 1932), based on bond energies, is the most widely used. On this scale, F = 4.0 is the highest, and Cs/Fr ≈ 0.7 is the lowest.
Section F — Valence, Oxide & Periodic Properties
21
What is the valency (combining capacity) of elements in group 15?
Group 15 elements (N, P, As…) have 5 valence electrons. They show valency of 3 or 5 (e.g., NH₃ shows valency 3; PF₅ shows valency 5 by using d-orbitals).
22
Oxides of elements on the left side of the periodic table (metals) are generally:
Metal oxides are generally basic in nature (e.g., Na₂O, CaO react with water to form hydroxides). Non-metal oxides are acidic (e.g., SO₃, P₄O₁₀). Middle elements like Al₂O₃ are amphoteric.
23
Which of the following is an amphoteric oxide?
Al₂O₃ (aluminium oxide) is amphoteric — it reacts with both acids (Al₂O₃ + H₂SO₄ → Al₂(SO₄)₃ + H₂O) and bases (Al₂O₃ + NaOH → NaAlO₂ + H₂O).
24
The metallic character of elements along a period (left to right):
Metallic character decreases across a period (left to right) as nuclear charge increases, making it harder for atoms to lose electrons. Non-metallic character simultaneously increases.
Section G — Anomalies & Special Cases
25
Mendeleev placed tellurium (Te, mass≈128) before iodine (I, mass≈127) based on properties. This is an example of:
This is a defect of Mendeleev's table — placing elements by properties violated the order of atomic masses. Moseley's atomic number resolved this: Te(Z=52) correctly precedes I(Z=53), validating the Modern Periodic Law.
26
Hydrogen is placed in group 1, yet it differs from alkali metals. Which property does hydrogen share with halogens (group 17)?
Hydrogen forms hydrides similar to halides (NaH resembles NaF) and can accept one electron to form H⁻ (hydride ion), analogous to halogens. It also lacks one electron from a filled shell, like halogens — hence the anomalous position.
27
In Mendeleev's periodic table, the noble gases (Group 18) were absent because:
Noble gases were not yet discovered when Mendeleev published his periodic table in 1869. Argon was discovered in 1894, and a new Group 0 (now 18) was added to accommodate them without disturbing the existing arrangement.
Section H — Periodic Trends — Mixed
28
Which of the following correctly represents the trend in second ionisation enthalpy (IE₂)?
IE₂ is always greater than IE₁ for the same element, because after removal of the first electron the remaining cation has a higher effective nuclear charge per electron, making the second electron harder to remove.
29
Diagonal relationship in the periodic table is best exhibited between:
The diagonal relationship occurs when an element resembles not its group neighbours but the element diagonally below-right, due to similar charge density. Classic pairs: Li–Mg and Be–Al. Both A and C are correct.
30
Which of the following elements is a metalloid (semi-metal)?
Germanium (Ge) is a metalloid — it lies on the staircase line separating metals and non-metals on the periodic table and displays properties of both. Other metalloids include Si, As, Sb, Te, and Po.
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