Which Of The Following Atoms Is Diamagnetic

Understanding magnetic properties at the atomic level is a fundamental concept in chemistry and physics that bridges the gap between abstract quantum mechanics and observable macroscopic behavior. That's why when faced with a question asking which of the following atoms is diamagnetic, the answer lies entirely in the electron configuration of the species in question. So diamagnetism is not a rare or exotic property; it is a universal phenomenon present in all matter, though it is often overshadowed by stronger magnetic effects like paramagnetism or ferromagnetism. To correctly identify a diamagnetic atom from a list of options, one must analyze the occupancy of its electron orbitals, specifically looking for the complete pairing of electrons.

The Quantum Mechanical Basis of Magnetism

At the heart of atomic magnetism lies the electron. Think about it: electrons possess two types of angular momentum: orbital angular momentum (from their motion around the nucleus) and spin angular momentum (an intrinsic property often visualized as the electron spinning on its axis). Think about it: both types of motion generate tiny magnetic fields, creating what we call a magnetic moment. The direction of this magnetic moment is quantized.

Short version: it depends. Long version — keep reading.

According to the Pauli Exclusion Principle, no two electrons in an atom can have the same set of four quantum numbers. Also, this means an orbital can hold a maximum of two electrons, and they must have opposite spins (spin-up, $m_s = +1/2$, and spin-down, $m_s = -1/2$). Still, when two electrons occupy the same orbital with opposite spins, their magnetic moments cancel each other out perfectly. This cancellation is the essence of diamagnetism.

Conversely, Hund’s Rule dictates that electrons occupy degenerate orbitals (orbitals of the same energy, like the three $p$ orbitals or five $d$ orbitals) singly before pairing up. These unpaired electrons have parallel spins, meaning their magnetic moments add up rather than cancel out. The result is a net magnetic moment for the atom, leading to paramagnetism—a weak attraction to an external magnetic field.

So, the diagnostic test for diamagnetism is simple in theory but requires careful execution in practice: A species is diamagnetic if and only if it has zero unpaired electrons. All electrons must be paired.

Step-by-Step Guide to Identifying Diamagnetic Atoms

When presented with a multiple-choice question listing several atoms or ions, follow this systematic workflow to isolate the diamagnetic candidate.

1. Determine the Number of Electrons

First, identify the total number of electrons for each species.

• For a neutral atom, the number of electrons equals the atomic number ($Z$).
• For a cation (positive ion), subtract the charge magnitude from $Z$.
• For an anion (negative ion), add the charge magnitude to $Z$.

Example: If the options are $\text{Fe}$, $\text{Fe}^{2+}$, $\text{Fe}^{3+}$, and $\text{Zn}^{2+}$:

• $\text{Fe}$ ($Z=26$): 26 electrons.
• $\text{Fe}^{2+}$: 24 electrons.
• $\text{Fe}^{3+}$: 23 electrons.
• $\text{Zn}^{2+}$ ($Z=30$): 28 electrons.

2. Write the Ground-State Electron Configuration

Write the full or condensed electron configuration using the Aufbau principle (filling order: $1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p...$) Surprisingly effective..

• Critical Exception Alert: For transition metals, the $(n-1)d$ orbitals are lower in energy than the $ns$ orbital after the atom is ionized or in the neutral ground state for Cr and Cu. Electrons are always removed from the highest principal quantum number ($n$) first (the $ns$ orbital) before the $(n-1)d$ orbitals.

Continuing the Iron/Zinc example:

• $\text{Fe}$: $[\text{Ar}] 4s^2 3d^6$
• $\text{Fe}^{2+}$: Remove two $4s$ electrons $\rightarrow [\text{Ar}] 3d^6$
• $\text{Fe}^{3+}$: Remove two $4s$ and one $3d$ $\rightarrow [\text{Ar}] 3d^5$
• $\text{Zn}^{2+}$: $[\text{Ar}] 4s^2 3d^{10} \rightarrow$ Remove two $4s \rightarrow [\text{Ar}] 3d^{10}$

3. Construct Orbital Diagrams (The "Box Method")

This is the most error-proof step. Draw boxes for the valence orbitals (specifically the $d$ or $f$ block for transition metals/lanthanides, $p$ block for main group). Fill them according to Hund's Rule: one electron per box, spin up, before pairing spin down.

• $\text{Fe}$ ($3d^6$): Five $d$ boxes. Fill: $\uparrow\downarrow, \uparrow, \uparrow, \uparrow, \uparrow$. Result: 4 unpaired electrons. Paramagnetic.
• $\text{Fe}^{2+}$ ($3d^6$): Same as neutral Fe valence. Result: 4 unpaired electrons. Paramagnetic.
• $\text{Fe}^{3+}$ ($3d^5$): Five $d$ boxes. Fill: $\uparrow, \uparrow, \uparrow, \uparrow, \uparrow$. Result: 5 unpaired electrons. Paramagnetic (maximum for d-block).
• $\text{Zn}^{2+}$ ($3d^{10}$): Five $d$ boxes. Fill: $\uparrow\downarrow, \uparrow\downarrow, \uparrow\downarrow, \uparrow\downarrow, \uparrow\downarrow$. Result: 0 unpaired electrons. Diamagnetic.

In this scenario, $\text{Zn}^{2+}$ is the correct answer That's the part that actually makes a difference..

Common Categories of Diamagnetic Atoms and Ions

Recognizing patterns saves time during exams or problem-solving. Here are the most frequent "correct answers" for diamagnetism questions.

1. Noble Gases (Group 18)

Helium ($\text{He}$), Neon ($\text{Ne}$), Argon ($\text{Ar}$), Krypton ($\text{Kr}$), Xenon ($\text{Xe}$), Radon ($\text{Rn}$), Oganesson ($\text{Og}$). Their defining feature is a completely filled valence shell ($ns^2 np^6$, except He $1s^2$). All electrons are paired. They are the classic textbook examples of diamagnetic atoms.

2. Atoms/Ions with $d^{10}$ or $f^{14}$ Configurations

A filled subshell is inherently diamagnetic because the maximum capacity ($2(2l+1)$ electrons) forces perfect pairing The details matter here..

• $d^{10}$: $\text{Zn}$ ($3d^{10}4s^2$), $\text{Cd}$, $\text{Hg}$, $\text{Cu}^+$ ($3d^{10}$), $\text{Ag}^+$, $\text{Au}^+$, $\text{Zn}^{2+}$, $\text{Cd}^{2+}$, $\text{Hg}^{2+}$.
• $f^{14}$: $\text{Yb}^{2+}$ ($4f^{14}$), $\text{Lu}^{3+}$ ($4f^{14}$), and many Actinide/Lanthanide ions in specific oxidation states.

3. Main Group Ions with Noble Gas Configurations

Many common ions achieve stability by mimicking the nearest noble gas Took long enough..

• Cations: $\text{