Which Of The Following Atoms Is Diamagnetic

Which ofthe following atoms is diamagnetic?

Understanding magnetism at the atomic level helps clarify why certain atoms exhibit diamagnetic behavior while others are paramagnetic. This article explains the underlying principles, walks through electron‑configuration rules, and identifies the atoms that are diamagnetic from a typical set of choices.

Understanding Magnetism at the Atomic Level

All matter exhibits magnetic properties, but the type of magnetism depends on how electrons occupy atomic orbitals. Two primary categories are:

• Paramagnetism – occurs when one or more electrons occupy degenerate orbitals with parallel spins, creating a net magnetic moment.
• Diamagnetism – arises when all electrons are paired in orbitals, resulting in a weak repulsion from an external magnetic field.

The key distinction lies in electron pairing. If every electron has a partner with opposite spin, the atom is diamagnetic; if any unpaired electrons remain, the atom shows paramagnetism. ## How Electron Configuration Determines Magnetism The electron configuration of an atom dictates its magnetic behavior. The order in which orbitals fill follows the Aufbau principle, Pauli exclusion principle, and Hund’s rule: 1. Aufbau principle – electrons fill lower‑energy orbitals before higher‑energy ones.
2. Pauli exclusion principle – no two electrons in the same orbital can share the same set of quantum numbers; they must have opposite spins.
3. Hund’s rule – electrons occupy separate orbitals of the same energy with parallel spins before pairing. When applying these rules, the presence or absence of unpaired electrons can be predicted directly from the configuration. For example:

• Carbon (C): 1s² 2s² 2p² → two unpaired electrons in the 2p subshell → paramagnetic.
• Neon (Ne): 1s² 2s² 2p⁶ → all p orbitals fully paired → diamagnetic.

Quick Checklist

• Step 1: Write the full electron configuration.
• Step 2: Identify the subshells that are partially filled.
• Step 3: Count the number of unpaired electrons.
• Step 4: If the count is zero → diamagnetic; otherwise → paramagnetic.

Common Atoms and Their Magnetic Properties

Below is a concise overview of several familiar atoms and whether they are diamagnetic or paramagnetic. | Atom | Electron Configuration | Unpaired Electrons | Magnetic Type | |------|------------------------|--------------------|---------------| | Hydrogen (H) | 1s¹ | 1 | Paramagnetic | | Helium (He) | 1s² | 0 | Diamagnetic | | Lithium (Li) | 1s² 2s¹ | 1 | Paramagnetic | | Beryllium (Be) | 1s² 2s² | 0 | Diamagnetic | | Boron (B) | 1s² 2s² 2p¹ | 1 | Paramagnetic | | Carbon (C) | 1s² 2s² 2p² | 2 | Paramagnetic | | Nitrogen (N) | 1s² 2s² 2p³ | 3 | Paramagnetic | | Oxygen (O) | 1s² 2s² 2p⁴ | 2 | Paramagnetic | | Fluorine (F) | 1s² 2s² 2p⁵ | 1 | Paramagnetic | | Neon (Ne) | 1s² 2s² 2p⁶ | 0 | Diamagnetic | | Sodium (Na) | [Ne] 3s¹ | 1 | Paramagnetic | | Magnesium (Mg) | [Ne] 3s² | 0 | Diamagnetic | | Aluminum (Al) | [Ne] 3s² 3p¹ | 1 | Paramagnetic | | Silicon (Si) | [Ne] 3s² 3p² | 2 | Paramagnetic | | Phosphorus (P) | [Ne] 3s² 3p³ | 3 | Paramagnetic | | Sulfur (S) | [Ne] 3s² 3p⁴ | 2 | Paramagnetic | | Chlorine (Cl) | [Ne] 3s² 3p⁵ | 1 | Paramagnetic | | Argon (Ar) | [Ne] 3s² 3p⁶ | 0 | Diamagnetic |

From the table, the diamagnetic atoms among the first 20 elements are He, Be, Ne, and Mg. Noble gases (He, Ne, Ar, Kr, Xe, Rn) are universally diamagnetic because their valence shells are completely filled.

Identifying Diamagnetic Atoms from a Typical Multiple‑Choice Question

Often textbooks pose the question: “Which of the following atoms is diamagnetic?” and provide a set of options such as O, N, C, Ne. Using the checklist above:

1. O (Oxygen) – 2p⁴ → two unpaired electrons → paramagnetic. 2. N (Nitrogen) – 2p³ → three unpaired electrons → paramagnetic.
2. C (Carbon) – 2p² → two unpaired electrons → paramagnetic.
3. Ne (Neon) – 2p⁶ → all electrons paired → diamagnetic.

Thus, Neon (Ne) is the correct answer. The same logic applies to any list: locate the atom whose electron configuration ends in a full subshell (e.g., s², p⁶, d¹⁰, f¹⁴).

Example Multiple‑Choice Set

Example Multiple‑Choice Set Consider the following question that might appear on an introductory chemistry exam:

Which of the following atoms is diamagnetic?

A. Fe (iron)  B. Cu (copper)  C. Zn (zinc)  D. Ar (argon)

Solution using the four‑step checklist

1. Write the full electron configuration

   • Fe: [Ar] 3d⁶ 4s²
   • Cu: [Ar] 3d¹⁰ 4s¹
   • Zn: [Ar] 3d¹⁰ 4s²
   • Ar: [Ne] 3s² 3p⁶ 2. Identify partially filled subshells
   • Fe: 3d subshell has six electrons → not completely filled (d⁶).
   • Cu: 3d subshell is full (d¹⁰) but the 4s subshell contains a single electron → partially filled.
   • Zn: Both 3d (d¹⁰) and 4s (s²) subshells are completely filled.
   • Ar: All subshells up to 3p are completely filled.

2. Count unpaired electrons - Fe: In a high‑spin d⁶ configuration there are four unpaired electrons.

   • Cu: The lone 4s electron is unpaired → one unpaired electron.
   • Zn: All electrons are paired → zero unpaired electrons. - Ar: All electrons are paired → zero unpaired electrons.

3. Determine magnetic type

   • Fe and Cu possess unpaired electrons → paramagnetic.
   • Zn and Ar have no unpaired electrons → diamagnetic.

Since the question asks for a single diamagnetic atom, both Zn and Ar satisfy the criterion. If only one answer is allowed, the noble‑gas option (Ar) is typically chosen because it is the classic example of a completely filled valence shell.

Conclusion

Recognizing whether an atom is diamagnetic or paramagnetic hinges on a simple electron‑counting procedure: write the configuration, locate any partially filled subshells, and tally unpaired electrons. When the count is zero, all electrons are paired and the substance exhibits diamagnetism—a weak, negative response to an external magnetic field. This principle not only clarifies the magnetic behavior of isolated atoms but also extends to ions, molecules, and solid‑state materials, where pairing or unpairing of electrons dictates properties ranging from magnetic resonance imaging contrast agents to the design of molecular magnets. Mastery of the four‑step checklist equips students to tackle multiple‑choice questions confidently and lays the groundwork for deeper exploration of magnetism in chemistry and physics.