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Atomic Structure & Periodicity Questions
Q.
Which quantum number determines the shape of an atomic orbital?
Magnetic quantum number (ml)
Azimuthal quantum number (l)
Spin quantum number (ms)
Principal quantum number (n)
Explanation:
The azimuthal (angular momentum) quantum number, l, determines the shape of the orbital (s, p, d, f). The principal quantum number, n, determines size/energy level.
Q.
The maximum number of electrons that can occupy a d-subshell is:
10
2
6
14
Explanation:
A d-subshell has 5 orbitals (ml = -2, -1, 0, +1, +2). Each orbital holds 2 electrons (Pauli exclusion principle), giving a maximum of 10 electrons.
Q.
Which of the following sets of quantum numbers (n, l, ml, ms) is NOT permissible?
(2, 0, 1, +1/2)
(2, 1, 0, +1/2)
(1, 0, 0, -1/2)
(3, 2, -1, -1/2)
Explanation:
The magnetic quantum number ml can range from -l to +l. For l = 0, ml must be 0. The set (2, 0, 1, +1/2) is impossible because ml cannot be 1 when l = 0.
Q.
The electron configuration of a neutral chlorine atom (Z = 17) is:
1s² 2s² 2p⁵ 3s² 3p⁶
1s² 2s² 2p⁶ 3s² 3p⁶
1s² 2s² 2p⁶ 3s² 3p⁵
1s² 2s² 2p⁶ 3s¹ 3p⁶
Explanation:
Chlorine has 17 electrons. Following the Aufbau principle: 1s² 2s² 2p⁶ 3s² 3p⁵.
Q.
Hund's rule states that:
Electrons pair up in the same orbital before occupying empty orbitals
No two electrons can have the same four quantum numbers
Electrons fill orbitals in order of increasing energy
Electrons fill degenerate orbitals singly first, with parallel spins, before pairing
Explanation:
Hund's rule of maximum multiplicity states that for a given electron configuration, the lowest energy term is the one with the greatest value of spin multiplicity. This means electrons fill degenerate orbitals singly first, with parallel spins, before pairing up.
Q.
The Pauli exclusion principle states that:
Electrons fill degenerate orbitals singly before pairing
The position and momentum of an electron cannot both be known precisely
No two electrons in an atom can have the same set of four quantum numbers
Electrons occupy the lowest available energy level
Explanation:
The Pauli exclusion principle states that no two electrons in the same atom can have the same set of all four quantum numbers. This limits each orbital to two electrons with opposite spins.
Q.
The Aufbau principle states that:
No two electrons can have the same four quantum numbers
Electrons fill degenerate orbitals singly before pairing
Orbitals of higher energy are filled before lower energy ones
Electrons fill the lowest available energy orbitals first
Explanation:
The Aufbau principle dictates that electrons fill atomic orbitals of the lowest available energy levels before occupying higher levels.
Q.
Which of the following is the correct order of orbital energy levels for a multi-electron atom?
1s < 2s < 3s < 2p < 3p < 3d < 4s
1s < 2s < 2p < 3s < 3p < 3d < 4s
1s < 2p < 2s < 3s < 3p < 4s < 3d
1s < 2s < 2p < 3s < 3p < 4s < 3d
Explanation:
For multi-electron atoms, orbital energies are not strictly determined by n alone. The (n+l) rule predicts the filling order: 1s < 2s < 2p < 3s < 3p < 4s < 3d.
Q.
The ground-state electron configuration of Cr (Z = 24) is:
[Ar] 3d⁵ 4s¹
[Ar] 3d⁴ 4s²
[Ar] 3d³ 4s² 4p¹
[Ar] 3d⁶ 4s⁰
Explanation:
Chromium is an exception to the Aufbau principle. A half-filled d-subshell provides extra stability, so one 4s electron moves to the 3d subshell: [Ar] 3d⁵ 4s¹.
Q.
The ground-state electron configuration of Cu (Z = 29) is:
[Ar] 3d⁸ 4s² 4p¹
[Ar] 3d¹⁰ 4s²
[Ar] 3d⁹ 4s²
[Ar] 3d¹⁰ 4s¹
Explanation:
Copper is another exception. A fully filled d-subshell is more stable, so the configuration is [Ar] 3d¹⁰ 4s¹ rather than [Ar] 3d⁹ 4s².
Q.
Moving from left to right across a period, atomic radius generally:
Increases then decreases
Decreases
Remains constant
Increases
Explanation:
Across a period, effective nuclear charge increases while electrons enter the same principal shell. The nucleus pulls electrons closer, decreasing atomic radius.
Q.
Moving down a group, atomic radius generally:
Decreases
Decreases then increases
Remains constant
Increases
Explanation:
Down a group, the principal quantum number (n) increases, adding new electron shells. The increased distance from the nucleus outweighs the increased nuclear charge, so atomic radius increases.
Q.
Which of the following is the correct order of increasing electronegativity?
Cl < P < Si < Na
Na < P < Si < Cl
Si < Na < P < Cl
Na < Si < P < Cl
Explanation:
Electronegativity increases across a period and decreases down a group. The trend is: Na < Si < P < Cl.
Q.
The most electronegative element in the periodic table is:
Oxygen
Chlorine
Fluorine
Nitrogen
Explanation:
Fluorine is the most electronegative element, with a Pauling electronegativity of 3.98. It has a small atomic radius and high effective nuclear charge.
Q.
The first ionization energy generally increases across a period because:
Effective nuclear charge increases and atomic radius decreases
Atomic radius increases and shielding increases
Electron-electron repulsion decreases
New electron shells are added
Explanation:
Across a period, effective nuclear charge increases while atomic radius decreases, holding electrons more tightly and making them harder to remove. Hence, ionization energy increases.
Q.
Which element has the highest first ionization energy?
Sodium
Magnesium
Argon
Neon
Explanation:
Noble gases have the highest ionization energies within their periods due to stable, filled electron configurations. Among the choices, Neon (a noble gas) has the highest first ionization energy.
Q.
An exception to the general trend of ionization energy occurs between:
Oxygen and Fluorine
Beryllium and Boron
Carbon and Nitrogen
Lithium and Beryllium
Explanation:
Be (1s² 2s²) has a filled 2s subshell. B (1s² 2s² 2p¹) has an electron in a higher-energy 2p orbital that is more shielded. Thus, B has a lower ionization energy than Be, breaking the general increase across the period.
Q.
Electron affinity is generally defined as:
The energy required to break a chemical bond
The energy change when an electron is added to a gaseous atom
The energy released when an ionic bond forms
The energy required to remove an electron from a gaseous atom
Explanation:
Electron affinity is the energy change when an electron is added to a neutral gaseous atom. A more negative value indicates a greater tendency to accept an electron.
Q.
Which of the following elements has the most negative (highest) electron affinity?
Iodine
Bromine
Chlorine
Fluorine
Explanation:
Chlorine has the most negative electron affinity. Adding an electron to Cl gives a stable [Ne]3s²3p⁶ configuration. Fluorine, despite higher electronegativity, has a less negative electron affinity due to small size and electron-electron repulsion in the compact 2p subshell.
Q.
The elements of Group 18 (VIIIA) are called noble gases because:
They readily accept electrons to form anions
They readily donate electrons to form cations
They have stable, filled valence electron shells and are largely unreactive
They form strong covalent bonds with each other
Explanation:
Noble gases have a complete octet (or duet for He) of valence electrons, giving them exceptional stability and very low reactivity.
Q.
Which group of the periodic table contains the alkali metals?
Group 2
Group 1
Group 18
Group 17
Explanation:
Alkali metals are Group 1 elements (Li, Na, K, Rb, Cs, Fr). They have a single valence electron in an s-orbital and are highly reactive.
Q.
Which group of the periodic table contains the halogens?
Group 1
Group 16
Group 2
Group 17
Explanation:
Halogens are Group 17 elements (F, Cl, Br, I, At). They have seven valence electrons and are one electron short of a noble gas configuration.
Q.
The transition metals are found in which block of the periodic table?
d-block
s-block
f-block
p-block
Explanation:
Transition metals correspond to the d-block elements, located between Groups 2 and 13. Their differentiating electrons enter d-orbitals.
Q.
Lanthanides and actinides belong to which block of the periodic table?
s-block
d-block
f-block
p-block
Explanation:
Lanthanides (4f series) and actinides (5f series) are f-block elements. Their differentiating electrons enter f-orbitals.
Q.
Which of the following species is isoelectronic with Ar (Z = 18)?
Ca²⁺
Na⁺
K⁺
Cl⁻
Explanation:
Isoelectronic species have the same number of electrons. Ar has 18 electrons. K⁺ (Z=19) loses one electron to also have 18 electrons.