Periodic Classification and Periodicity Questions and Answers

Q: The correct order of metallic radius of the elements is

The correct order is \(\mathrm{B \lt Be \lt Al \lt Li \lt Mg}\).The standard metallic radii are B 88 pm, Be 111 pm, Al 143 pm, Li 152 pm and Mg 160 pm. Radius decreases across a period (B is smaller than Be, Al smaller than Mg) and increases down a group (Mg is larger than Be, Al larger than B).Li and Mg are a diagonal pair with similar sizes, Mg being slightly the larger. The other orders put the period-2 elements in the wrong sequence or place the large atoms Li and Mg too early.

Q: Which of the following represents the correct increasing order of first ionization enthalpy for Ca, Ba, S, Se and Ar?

The increasing order is \(\mathrm{Ba \lt Ca \lt Se \lt S \lt Ar}\).Ionisation enthalpy decreases down a group and increases across a period. Ba is below Ca in group 2, so Ba has the lowest value; Se is below S in group 16, so Se is lower than S; and argon, a noble gas, has the highest.Approximate values in kJ/mol are Ba 503, Ca 590, Se 941, S 1000 and Ar 1521. The options placing Ar first or Ca above S reverse these trends.

Q: Which of the following statement is incorrect about the periodic trend?

The incorrect statement is Ionisation enthalpies generally decrease across a period and increase down a group.The trend is the reverse: ionisation enthalpy generally increases across a period, as nuclear charge rises and atoms shrink, and decreases down a group, as the outer electrons are farther from the nucleus and better shielded.The other statements are the standard trends: atomic radius falls across a period and rises down a group; electronegativity rises across and falls down; electron gain enthalpy generally becomes more negative across and less negative down; and for representative elements valence equals the number of outer electrons (or eight minus it).

Q: A measure of an atom’s ability to attract electrons from a neighbouring atom to which it is bonded is called as :

The measure is electronegativity.Electronegativity, as defined by Pauling, is the tendency of an atom in a molecule to attract the shared pair of electrons of a bond towards itself. Fluorine is the most electronegative element.Electron affinity is the energy change when an isolated gaseous atom gains an electron, not an atom in a bond. Ionisation potential is the energy to remove an electron from an isolated atom, and electropositivity is the tendency to lose electrons, the opposite property.

Q: Alkali metals are present in ______ of the periodic table.

Alkali metals are in Group-I of the periodic table.Lithium, sodium, potassium, rubidium, caesium and francium form Group 1 (IA). Each has one electron in its outer shell, which it loses easily, making them highly reactive metals.Group zero (Group 18) holds the noble gases, which are unreactive because their shells are full. Group III (13) holds boron and aluminium, and Group IV (14) carbon and silicon. The alkaline earth metals, such as calcium and magnesium, are the neighbouring Group 2.

Q: The correct order of First Ionization energy of group 13 elements is :

The order is B > Tl > Ga > Al > In.The first ionisation energies of group 13 are approximately B 801, Al 577, Ga 579, In 558 and Tl 589 kJ mol-1. Boron is far higher because of its small size. After that the values do not fall steadily. Gallium is slightly above aluminium because the ten 3d electrons it has added shield the nucleus poorly, and thallium is above both because of the poor shielding of its 4f and 5d electrons (the lanthanide contraction) and relativistic stabilisation of its 6s electrons.The order B > Al > Ga > In > Tl is the simple trend of falling ionisation energy down a group, which group 13 does not follow. The other two options misplace aluminium against gallium or put Al lowest.

Q: Size of the d orbitals for Si, P, S and Cl follow the order :

The order is Si \(\gt\) P \(\gt\) S \(\gt\) Cl.Across the third period the nuclear charge rises from Si (Z = 14) to Cl (Z = 17), while the added electrons enter the same shell and screen one another poorly. The effective nuclear charge felt by the outer orbitals therefore increases, and all of them, including the vacant 3d orbitals, contract. The 3d orbitals are largest and most diffuse at silicon and smallest at chlorine.The orders beginning with Cl reverse this trend, and P \(\gt\) S \(\gt\) Si \(\gt\) Cl misplaces silicon, which has the lowest nuclear charge of the four.

Q: Which of the following statements regarding solubility of LiF and LiI in water at room temperature is correct ?

LiI is more soluble than LiF.Solubility depends on the balance between lattice energy and hydration energy. LiF combines the smallest cation with the smallest anion, so its lattice energy is very high (about 1036 kJ/mol) and hydration cannot overcome it; LiF is only sparingly soluble, about 0.13 g per 100 mL.In LiI the large, polarisable iodide gives a much lower lattice energy, while the small Li+ is still strongly hydrated, so LiI is very soluble (over 150 g per 100 mL). The two are therefore neither equally soluble nor both insoluble, and LiF is the less soluble one.

Q: Based on Pauling's electronegativity scale hydrogen can oxidize :

Hydrogen can oxidize boron.On Pauling's scale hydrogen has an electronegativity of 2.20, boron 2.04, carbon 2.55, nitrogen 3.04 and oxygen 3.44. In a bond between two elements, the more electronegative one is assigned the negative oxidation state. Only boron is less electronegative than hydrogen, so in boranes such as B2H6 hydrogen is formally H− (hydridic) and boron is oxidised.Carbon, nitrogen and oxygen are all more electronegative than hydrogen, so in CH4, NH3 and H2O hydrogen is +1: it is oxidised by them rather than oxidising them.

31.

The correct order of metallic radius of the elements is

A.\(\mathrm{Be \lt B \lt Al \lt Mg \lt Li}\)
B.\(\mathrm{Mg \lt Li \lt Al \lt B \lt Be}\)
C.\(\mathrm{Li \lt Mg \lt B \lt Be \lt Al}\)
D.\(\mathrm{B \lt Be \lt Al \lt Li \lt Mg}\)
E.\(\mathrm{Li \lt Mg \lt Be \lt B \lt Al}\)

Exam Year(s)

KEAM - 2026,

Answer: D

The correct order is \(\mathrm{B \lt Be \lt Al \lt Li \lt Mg}\).

The standard metallic radii are B 88 pm, Be 111 pm, Al 143 pm, Li 152 pm and Mg 160 pm. Radius decreases across a period (B is smaller than Be, Al smaller than Mg) and increases down a group (Mg is larger than Be, Al larger than B).

Li and Mg are a diagonal pair with similar sizes, Mg being slightly the larger. The other orders put the period-2 elements in the wrong sequence or place the large atoms Li and Mg too early.

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32.

Which of the following represents the correct increasing order of first ionization enthalpy for Ca, Ba, S, Se and Ar?

A.\(\mathrm{Ca \lt S \lt Ba \lt Se \lt Ar}\)
B.\(\mathrm{Ar \lt S \lt Ba \lt Se \lt Ca}\)
C.\(\mathrm{Ba \lt Ca \lt Se \lt S \lt Ar}\)
D.\(\mathrm{Ba \lt Ca \lt S \lt Se \lt Ar}\)
E.\(\mathrm{Ca \lt S \lt Ar \lt Se \lt Ba}\)

Exam Year(s)

KEAM - 2026,

Answer: C

The increasing order is \(\mathrm{Ba \lt Ca \lt Se \lt S \lt Ar}\).

Ionisation enthalpy decreases down a group and increases across a period. Ba is below Ca in group 2, so Ba has the lowest value; Se is below S in group 16, so Se is lower than S; and argon, a noble gas, has the highest.

Approximate values in kJ/mol are Ba 503, Ca 590, Se 941, S 1000 and Ar 1521. The options placing Ar first or Ca above S reverse these trends.

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33.

Which of the following statement is incorrect about the periodic trend?

A.Atomic radii decrease while going from left to right in a period and increase with atomic number in a group.
B.Ionisation enthalpies generally decrease across a period and increase down a group.
C.Electronegativity generally increases across a period and decreases down a group.
D.Electron gain enthalpies in general, become more negative across a period and less negative down a group.
E.Among representative elements, the valence is equal to the number of electrons in the outermost orbital.

Exam Year(s)

KEAM - 2026,

Answer: B

The incorrect statement is Ionisation enthalpies generally decrease across a period and increase down a group.

The trend is the reverse: ionisation enthalpy generally increases across a period, as nuclear charge rises and atoms shrink, and decreases down a group, as the outer electrons are farther from the nucleus and better shielded.

The other statements are the standard trends: atomic radius falls across a period and rises down a group; electronegativity rises across and falls down; electron gain enthalpy generally becomes more negative across and less negative down; and for representative elements valence equals the number of outer electrons (or eight minus it).

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34.

A measure of an atom’s ability to attract electrons from a neighbouring atom to which it is bonded is called as :

A.Electron affinity
B.Electropositivity
C.Electronegativity
D.Ionisation potential

Exam Year(s)

MPPSC Forest Service - 2022,

Answer: C

The measure is electronegativity.

Electronegativity, as defined by Pauling, is the tendency of an atom in a molecule to attract the shared pair of electrons of a bond towards itself. Fluorine is the most electronegative element.

Electron affinity is the energy change when an isolated gaseous atom gains an electron, not an atom in a bond. Ionisation potential is the energy to remove an electron from an isolated atom, and electropositivity is the tendency to lose electrons, the opposite property.

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35.

Alkali metals are present in ______ of the periodic table.

A.Zero group
B.Group-III
C.Group-IV
D.Group-I

Exam Year(s)

MPPSC Forest Service - 2024,

Answer: D

Alkali metals are in Group-I of the periodic table.

Lithium, sodium, potassium, rubidium, caesium and francium form Group 1 (IA). Each has one electron in its outer shell, which it loses easily, making them highly reactive metals.

Group zero (Group 18) holds the noble gases, which are unreactive because their shells are full. Group III (13) holds boron and aluminium, and Group IV (14) carbon and silicon. The alkaline earth metals, such as calcium and magnesium, are the neighbouring Group 2.

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36.

The correct order of First Ionization energy of group 13 elements is :

A.B \(\gt\) Tl \(\gt\) Ga \(\gt\) Al \(\gt\) In
B.B \(\gt\) Tl \(\gt\) Al \(\gt\) Ga \(\gt\) In
C.B \(\gt\) Al \(\gt\) Ga \(\gt\) In \(\gt\) Tl
D.B \(\gt\) Ga \(\gt\) Tl \(\gt\) In \(\gt\) Al

Exam Year(s)

MH SET - 2018, SET Exam - 2018,

Answer: A

The order is B > Tl > Ga > Al > In.

The first ionisation energies of group 13 are approximately B 801, Al 577, Ga 579, In 558 and Tl 589 kJ mol^-1. Boron is far higher because of its small size. After that the values do not fall steadily. Gallium is slightly above aluminium because the ten 3d electrons it has added shield the nucleus poorly, and thallium is above both because of the poor shielding of its 4f and 5d electrons (the lanthanide contraction) and relativistic stabilisation of its 6s electrons.

The order B > Al > Ga > In > Tl is the simple trend of falling ionisation energy down a group, which group 13 does not follow. The other two options misplace aluminium against gallium or put Al lowest.

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37.

Size of the d orbitals for Si, P, S and Cl follow the order :

A.Si \(\gt\) P \(\gt\) S \(\gt\) Cl
B.Cl \(\gt\) P \(\gt\) S \(\gt\) Si
C.Cl \(\gt\) S \(\gt\) P \(\gt\) Si
D.P \(\gt\) S \(\gt\) Si \(\gt\) Cl

Exam Year(s)

MH SET - 2019, SET Exam - 2019,

Answer: A

The order is Si \(\gt\) P \(\gt\) S \(\gt\) Cl.

Across the third period the nuclear charge rises from Si (Z = 14) to Cl (Z = 17), while the added electrons enter the same shell and screen one another poorly. The effective nuclear charge felt by the outer orbitals therefore increases, and all of them, including the vacant 3d orbitals, contract. The 3d orbitals are largest and most diffuse at silicon and smallest at chlorine.

The orders beginning with Cl reverse this trend, and P \(\gt\) S \(\gt\) Si \(\gt\) Cl misplaces silicon, which has the lowest nuclear charge of the four.

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38.

Out of As, Bi, Sb and graphite semimetals, which one of the following is considerded as metalloid ?

A.Bi
B.As
C.Sp
D.Graphite

Exam Year(s)

MH SET - 2019, SET Exam - 2019,

Answer: B

The answer is As, the one printed option that is a metalloid.

Arsenic is a standard metalloid: it has a metallic grey allotrope but brittle, semiconducting behaviour and amphoteric oxides. Bismuth is listed in the stem as a semimetal and is usually treated as a post-transition metal, and graphite is a form of carbon, a non-metal that conducts along its layers.

Option C prints 'Sp', which is not the symbol of any element. Antimony is Sb, and it is also commonly called a metalloid, so if the option was meant to be antimony the question would have two answers. As printed, only arsenic fits, which agrees with the official key. The booklet's own list of semimetals in the stem is not a standard classification, so this item rests on the usual textbook grouping.

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39.

Which of the following statements regarding solubility of LiF and LiI in water at room temperature is correct ?

A.Both are equally soluble
B.Both are insoluble
C.LiF is more soluble than LiI
D.LiI is more soluble than LiF

Exam Year(s)

MH SET - 2020, SET Exam - 2020,

Answer: D

LiI is more soluble than LiF.

Solubility depends on the balance between lattice energy and hydration energy. LiF combines the smallest cation with the smallest anion, so its lattice energy is very high (about 1036 kJ/mol) and hydration cannot overcome it; LiF is only sparingly soluble, about 0.13 g per 100 mL.

In LiI the large, polarisable iodide gives a much lower lattice energy, while the small Li⁺ is still strongly hydrated, so LiI is very soluble (over 150 g per 100 mL). The two are therefore neither equally soluble nor both insoluble, and LiF is the less soluble one.

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40.

Based on Pauling's electronegativity scale hydrogen can oxidize :

A.Boron
B.Carbon
C.Nitrogen
D.Oxygen

Exam Year(s)

MH SET - 2023, SET Exam - 2023,

Answer: A

Hydrogen can oxidize boron.

On Pauling's scale hydrogen has an electronegativity of 2.20, boron 2.04, carbon 2.55, nitrogen 3.04 and oxygen 3.44. In a bond between two elements, the more electronegative one is assigned the negative oxidation state. Only boron is less electronegative than hydrogen, so in boranes such as B₂H₆ hydrogen is formally H⁻ (hydridic) and boron is oxidised.

Carbon, nitrogen and oxygen are all more electronegative than hydrogen, so in CH₄, NH₃ and H₂O hydrogen is +1: it is oxidised by them rather than oxidising them.

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