CBSE Class 10 Chemistry Periodic Classification of Elements MCQs Set 26

Answer: (c) 7 and 18

Explanation:
1. The long form of the periodic table has 7 horizontal rows called periods.
2. It has 18 vertical columns called groups.
3. So the pair is 7 and 18.

Teacher's Note:
1. Remember 7 periods and 18 groups together.
2. The 7th period is incomplete only in the sense that new elements keep being added.

Answer: (c) greater

Explanation:
1. After one electron is removed, the positive ion holds its remaining electrons more strongly.
2. More energy is needed to remove the next electron, so IE₂ > IE₁.

Teacher's Note:
1. Successive ionisation energies always increase.
2. A big jump shows that a stable noble-gas core is being broken.

Answer: (d) increase in atomic size and also shielding effect

Explanation:
1. Down a group a new shell is added, so the atomic size increases.
2. Inner shells shield the outer electron more, so it is held less tightly and is easier to remove.

Teacher's Note:
1. The nuclear charge does increase, but its effect is cancelled by the larger size and shielding.
2. Shielding means inner electrons screen the outer ones from the nucleus.

Answer: (c) half the distance between the centres of the nuclei of two non-bonded atoms of adjacent molecules

Explanation:
1. The van der Waals radius is measured between atoms that are close but not bonded.
2. It is half the distance between the nuclei of two such non-bonded atoms of neighbouring molecules.

Teacher's Note:
1. Covalent radius is half the bond length between two bonded atoms.
2. The van der Waals radius is always larger than the covalent radius because there is no bond pulling the atoms together.

Answer: (b) Be²⁺, Mg²⁺, Na⁺

Explanation:
1. Be²⁺ has only the 1s shell, so it is the smallest.
2. Mg²⁺ and Na⁺ are isoelectronic (2 shells), and Na⁺ has fewer protons, so it is slightly larger than Mg²⁺.
3. So Be²⁺ < Mg²⁺ < Na⁺ is increasing.
4. In (a) K⁺ is larger than Ca²⁺; in (c) and (d) the order is not increasing (O^2– > F^–, and S^2– > O^2–).

Teacher's Note:
1. For isoelectronic ions, the larger the nuclear charge, the smaller the radius.
2. Na⁺ and Mg²⁺ have the same number of electrons, but Mg has 12 protons against 11.

Answer: (d) Ca²⁺ < Cl^– < S^2– < P^3–

Explanation:
1. All four ions have 18 electrons, so they are isoelectronic.
2. The nuclear charge is P (15) < S (16) < Cl (17) < Ca (20).
3. A higher nuclear charge gives a smaller radius, so the radii increase as Ca²⁺ < Cl^– < S^2– < P^3–.

Teacher's Note:
1. In an isoelectronic series, the ion with more protons is smaller.
2. Anions with higher negative charge are the largest.

Answer: (d) phosphorus

Explanation:
1. The valence shell is 5s²5p³, with 5 valence electrons (this element is antimony).
2. Five valence electrons put it in group 15, the same group as phosphorus.

Teacher's Note:
1. Chemical behaviour depends mainly on the valence electron configuration.
2. The filled 4d¹⁰ subshell is inside the valence shell and does not change the group.

Answer: (b) Ne

Explanation:
1. Ne has a completely filled 2s²2p⁶ shell, so it has no tendency to accept an extra electron.
2. Na and F readily accept electrons and release energy.

Teacher's Note:
1. Noble gases have the most stable octets (duplet for He).
2. Their electron gain enthalpy is taken as zero at this level.

Answer: (b) 2s²2p⁵

Explanation:
1. The most electronegative element is fluorine (Z = 9).
2. Its configuration is 1s²2s²2p⁵, so the outermost configuration is 2s²2p⁵.

Teacher's Note:
1. Electronegativity increases across a period and decreases down a group.
2. The top-right element (excluding noble gases) is fluorine.

Answer: (c) some of the elements of the 2nd and 3rd periods

Explanation:
1. Diagonal relationship appears between pairs such as Li and Mg, Be and Al, B and Si.
2. These pairs are found in the 2nd and 3rd periods only.

Teacher's Note:
1. Similar size (charge to radius ratio) causes the similarity.
2. It does not hold for every diagonal pair.

Answer: (d) just enough to cause the electron to escape from the isolated gaseous atom

Explanation:
1. Ionisation energy is the minimum energy needed to remove the most loosely held electron from an isolated gaseous atom.
2. Adding an electron relates to electron affinity, not ionisation energy.

Teacher's Note:
1. The atom must be isolated and in the gaseous state.
2. Unit used: kJ mol^–1 or eV per atom.

Answer: (b) 1s²2s²2p⁶3s¹

Explanation:
1. The tendency to form a unipositive ion is highest for the element with the lowest first ionisation energy.
2. The configuration with one 3s electron is sodium, an alkali metal, which has the lowest first ionisation energy of the four.

Teacher's Note:
1. Alkali metals easily lose their single outer electron.
2. Mg, Si and P have higher ionisation energies than Na.

Answer: (d) 1s²2s²2p⁵

Explanation:
1. The configurations belong to H, Li, O and F respectively.
2. Fluorine needs one electron to complete its octave and releases the most energy per atom (about 328 kJ mol^–1) among these four.

Teacher's Note:
1. The nearer an atom is to a complete shell, the larger its electron affinity.
2. Alkali metals have low electron affinity values.

Answer: (a) 1s²2s²2p⁶3s¹

Explanation:
1. The four configurations belong to Na, F, Ne and Si.
2. Na has a single electron in a new shell, far from the nucleus, so it loses it most easily.
3. Its first ionisation energy is the lowest of the four.

Teacher's Note:
1. Alkali metals have the lowest ionisation energies in their periods.
2. Noble gases have the highest.

Answer: (a) zero

Explanation:
1. Noble gases have completely filled outermost shells and do not normally gain, lose or share electrons.
2. So their valency is zero.

Teacher's Note:
1. Chemical inertness is the result of a stable octet (duplet for helium).
2. A few heavier members (Xe, Kr) form compounds only in special conditions.

Answer: (c) Ni²⁺

Explanation:
1. Zn²⁺ (3d¹⁰), Cu⁺ (3d¹⁰) and Ag⁺ (4d¹⁰) have all electrons paired.
2. Ni²⁺ has 3d⁸, which contains two unpaired electrons, so it is paramagnetic.

Teacher's Note:
1. Paramagnetic species contain unpaired electrons.
2. Ions with a completely filled d subshell are diamagnetic.

Answer: (b) Mg < Na < K < Rb

Explanation:
1. Mg is in the same period as Na but has a higher nuclear charge, so Mg is smaller than Na.
2. Down group 1 the size increases: Na < K < Rb.
3. So the order is Mg < Na < K < Rb.

Teacher's Note:
1. Size decreases across a period and increases down a group.
2. Compare elements in the same period first, then in the same group.

Answer: (d) O^2– > F^– > O > F

Explanation:
1. O^2– and F^– are isoelectronic; O^2– has fewer protons, so it is larger.
2. Each anion is larger than its own atom, and O is larger than F because the nuclear charge is lower.
3. Order: O^2– > F^– > O > F.

Teacher's Note:
1. Anions are larger than the neutral atoms of the same element.
2. Among atoms of one period, the one with the higher Z is smaller.

Answer: (b) Cl

Explanation:
1. Fluorine is small, so the 2p electrons repel an incoming electron strongly.
2. Chlorine therefore has the highest electron gain enthalpy of all elements.

Teacher's Note:
1. The decreasing trend down the group starts from Cl, not F.
2. F is an exception because of the small size.

Answer: (b) high ionisation energy

Explanation:
1. Highly electronegative atoms are small and hold electrons strongly.
2. Hence they have high ionisation energy and high electron affinity.

Teacher's Note:
1. Fluorine is the standard example of high electronegativity.
2. Large size and low ionisation energy are properties of metals, not of strongly electronegative atoms.

Answer: (c) additional electrons are accommodated in a new electron shell

Explanation:
1. On moving down a group, each element has one more shell than the one above.
2. The extra shell puts the outer electrons farther from the nucleus, so the atomic radius increases.

Teacher's Note:
1. Atomic number does increase down a group, but it is the new shell that determines size.
2. Shielding by inner shells offsets the greater nuclear charge.

Answer: (b) Both A and R are true, but R is not the correct explanation of A.

Explanation:
1. Nitrogen has a half-filled 2p³ subshell, which is extra stable, so its first ionisation energy (1402 kJ mol^–1) is higher than that of oxygen (1314 kJ mol^–1).
2. The reason is also true, but it would predict a higher value for oxygen, so it does not explain the assertion.

Teacher's Note:
1. The real explanation is the stability of the half-filled subshell.
2. Always check whether the reason actually explains the assertion.