Atomic Structure & Periodicity
About Chemistry MCQs
Chemistry is the scientific study of the properties and behavior of matter. It is a fundamental subject for medical and engineering entry tests, as well as a standard component of science exams. Understanding chemical reactions, the periodic table, and molecular structures is key to success in this subject.
Topics Covered in This Category
- Physical Chemistry: Atomic structure, chemical bonding, thermodynamics, and equilibrium.
- Organic Chemistry: Hydrocarbons, functional groups, and reaction mechanisms.
- Inorganic Chemistry: Transition metals, coordination compounds, and the periodic table.
- Analytical Chemistry: Titrations, spectroscopy, and qualitative analysis.
How to Master Chemistry MCQs
Chemistry requires a balance of understanding and memorization. Keep a notebook of chemical reactions and their conditions. Practice numerical problems daily. For MCQs, pay close attention to units, signs, and exceptions to general rules, as examiners often use these to create trick questions.
All Questions for: Atomic Structure & Periodicity
Question 1: Which quantum number determines the shape of an atomic orbital?
- A Azimuthal quantum number (l)
- B Principal quantum number (n)
- C Magnetic quantum number (ml)
- D Spin quantum number (ms)
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.
Question 2: The maximum number of electrons that can occupy a d-subshell is:
- A 10
- B 6
- C 2
- D 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.
Question 3: Which of the following sets of quantum numbers (n, l, ml, ms) is NOT permissible?
- A (2, 0, 1, +1/2)
- B (2, 1, 0, +1/2)
- C (3, 2, -1, -1/2)
- D (1, 0, 0, -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.
Question 4: The electron configuration of a neutral chlorine atom (Z = 17) is:
- A 1s² 2s² 2p⁶ 3s² 3p⁵
- B 1s² 2s² 2p⁶ 3s² 3p⁶
- C 1s² 2s² 2p⁵ 3s² 3p⁶
- D 1s² 2s² 2p⁶ 3s¹ 3p⁶
Explanation: Chlorine has 17 electrons. Following the Aufbau principle: 1s² 2s² 2p⁶ 3s² 3p⁵.
Question 5: Hund's rule states that:
- A Electrons fill degenerate orbitals singly first, with parallel spins, before pairing
- B Electrons fill orbitals in order of increasing energy
- C No two electrons can have the same four quantum numbers
- D Electrons pair up in the same orbital before occupying empty orbitals
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.
Question 6: The Pauli exclusion principle states that:
- A No two electrons in an atom can have the same set of four quantum numbers
- B Electrons occupy the lowest available energy level
- C Electrons fill degenerate orbitals singly before pairing
- D The position and momentum of an electron cannot both be known precisely
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.
Question 7: The Aufbau principle states that:
- A Electrons fill the lowest available energy orbitals first
- B Electrons fill degenerate orbitals singly before pairing
- C No two electrons can have the same four quantum numbers
- D Orbitals of higher energy are filled before lower energy ones
Explanation: The Aufbau principle dictates that electrons fill atomic orbitals of the lowest available energy levels before occupying higher levels.
Question 8: Which of the following is the correct order of orbital energy levels for a multi-electron atom?
- A 1s < 2s < 2p < 3s < 3p < 4s < 3d
- B 1s < 2s < 2p < 3s < 3p < 3d < 4s
- C 1s < 2p < 2s < 3s < 3p < 4s < 3d
- D 1s < 2s < 3s < 2p < 3p < 3d < 4s
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.
Question 9: The ground-state electron configuration of Cr (Z = 24) is:
- A [Ar] 3d⁵ 4s¹
- B [Ar] 3d⁴ 4s²
- C [Ar] 3d⁶ 4s⁰
- D [Ar] 3d³ 4s² 4p¹
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.
Question 10: The ground-state electron configuration of Cu (Z = 29) is:
- A [Ar] 3d¹⁰ 4s¹
- B [Ar] 3d⁹ 4s²
- C [Ar] 3d¹⁰ 4s²
- D [Ar] 3d⁸ 4s² 4p¹
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.
Question 11: Moving from left to right across a period, atomic radius generally:
- A Decreases
- B Increases
- C Remains constant
- D Increases then decreases
Explanation: Across a period, effective nuclear charge increases while electrons enter the same principal shell. The nucleus pulls electrons closer, decreasing atomic radius.
Question 12: Moving down a group, atomic radius generally:
- A Increases
- B Decreases
- C Remains constant
- D Decreases then 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.
Question 13: Which of the following is the correct order of increasing electronegativity?
- A Na < Si < P < Cl
- B Cl < P < Si < Na
- C Na < P < Si < Cl
- D Si < Na < P < Cl
Explanation: Electronegativity increases across a period and decreases down a group. The trend is: Na < Si < P < Cl.
Question 14: The most electronegative element in the periodic table is:
- A Fluorine
- B Oxygen
- C Chlorine
- D 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.
Question 15: The first ionization energy generally increases across a period because:
- A Effective nuclear charge increases and atomic radius decreases
- B Atomic radius increases and shielding increases
- C New electron shells are added
- D Electron-electron repulsion decreases
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.
Question 16: Which element has the highest first ionization energy?
- A Neon
- B Sodium
- C Magnesium
- D Argon
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.
Question 17: An exception to the general trend of ionization energy occurs between:
- A Beryllium and Boron
- B Lithium and Beryllium
- C Carbon and Nitrogen
- D Oxygen and Fluorine
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.
Question 18: Electron affinity is generally defined as:
- A The energy change when an electron is added to a gaseous atom
- B The energy required to remove an electron from a gaseous atom
- C The energy required to break a chemical bond
- D The energy released when an ionic bond forms
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.
Question 19: Which of the following elements has the most negative (highest) electron affinity?
- A Chlorine
- B Fluorine
- C Bromine
- D Iodine
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.
Question 20: The elements of Group 18 (VIIIA) are called noble gases because:
- A They have stable, filled valence electron shells and are largely unreactive
- B They readily donate electrons to form cations
- C They readily accept electrons to form anions
- D 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.
Question 21: Which group of the periodic table contains the alkali metals?
- A Group 1
- B Group 2
- C Group 17
- D Group 18
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.
Question 22: Which group of the periodic table contains the halogens?
- A Group 17
- B Group 1
- C Group 2
- D Group 16
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.
Question 23: The transition metals are found in which block of the periodic table?
- A d-block
- B s-block
- C p-block
- D f-block
Explanation: Transition metals correspond to the d-block elements, located between Groups 2 and 13. Their differentiating electrons enter d-orbitals.
Question 24: Lanthanides and actinides belong to which block of the periodic table?
- A f-block
- B d-block
- C s-block
- D p-block
Explanation: Lanthanides (4f series) and actinides (5f series) are f-block elements. Their differentiating electrons enter f-orbitals.
Question 25: Which of the following species is isoelectronic with Ar (Z = 18)?
- A K⁺
- B Ca²⁺
- C Cl⁻
- D Na⁺
Explanation: Isoelectronic species have the same number of electrons. Ar has 18 electrons. K⁺ (Z=19) loses one electron to also have 18 electrons.