Which One Of The Following Is Paramagnetic

Paramagnetism describes the weak attraction of a material to an external magnetic field, a property tied directly to the presence of unpaired electrons in its atomic or molecular structure, which makes the question which one of the following is paramagnetic a common fixture in chemistry exams and coursework. The following sections break down the underlying science of paramagnetic behavior, outline step-by-step methods to identify paramagnetic species, and provide practice examples to help you master this concept for any assessment.

Introduction

Paramagnetism is one of three primary types of magnetic behavior in matter, alongside diamagnetism and ferromagnetism. Unlike diamagnetic materials, which are repelled by magnetic fields and have all electrons paired, paramagnetic materials are weakly attracted to magnetic fields because they contain at least one unpaired electron. Ferromagnetic materials, by contrast, are strongly attracted and retain magnetization, but this requires aligned unpaired electrons in domains, a separate mechanism from paramagnetism.

The question which one of the following is paramagnetic typically presents a list of chemical species – atoms, ions, or molecules – and asks test-takers to identify the one with unpaired electrons. Plus, common examples include diatomic molecules like O2 vs N2, transition metal complexes like [Fe(CN)6]4- vs [Fe(CN)6]3-, or main group ions like Na+ vs Cu2+. Mastering this question requires a solid understanding of electron configuration, Hund’s rule, and molecular orbital theory for covalent species.

Steps to Identify Paramagnetic Species

To answer which one of the following is paramagnetic correctly, follow this systematic process for any given chemical species:

Determine the total number of electrons in the atom, ion, or molecule. For ions, add electrons for negative charges, subtract electrons for positive charges. For molecules, sum the total electrons of all atoms in the species, accounting for any ionic bonding adjustments.
Write the electron configuration for the species, following the Aufbau principle, Hund’s rule, and the Pauli exclusion principle. For transition metals, remember that 4s electrons are lost before 3d electrons when forming cations.
Check for unpaired electrons in the highest energy orbitals. Hund’s rule states that electrons will fill degenerate orbitals (orbitals with the same energy, like the three p orbitals or five d orbitals) singly first, with parallel spins, before pairing. Any orbital with a single electron counts as unpaired.
Confirm with magnetic behavior (if applicable): Paramagnetic species are attracted to magnetic fields, while diamagnetic species are repelled. A quick rule of thumb for main group species: all electrons paired = diamagnetic; at least one unpaired electron = paramagnetic.

Special Cases: Covalent Molecules

For diatomic or polyatomic molecules, electron configuration alone is not enough – you must use molecular orbital (MO) theory to determine bonding and antibonding electrons. Take this: O2 has 16 total electrons (8 from each oxygen atom). Its MO diagram fills core 1s and 2s orbitals first, then valence orbitals: σ2s, σ2s, σ2pz, π2px, π2py, then π2px, π2py. The last two electrons go into separate π orbitals (per Hund’s rule), giving two unpaired electrons – so O2 is paramagnetic, which explains why liquid oxygen is visibly attracted to a magnet. N2, by contrast, has 14 total electrons, fills all bonding valence orbitals with paired electrons, and has no unpaired electrons – making it diamagnetic.

Special Cases: Transition Metal Complexes

Transition metal complexes require accounting for ligand field splitting, which changes the relative energy of d orbitals. Strong field ligands (like CN-, CO) cause large splitting, forcing electrons to pair in lower energy d orbitals (low spin complexes). Weak field ligands (like H2O, Cl-) cause small splitting, so electrons fill all d orbitals singly first (high spin complexes). For example:

[Fe(CN)6]4-: Fe2+ has 6 d electrons. CN- is a strong field ligand, so d electrons pair in t2g orbitals: t2g6 eg0 – all paired, diamagnetic.
[Fe(CN)6]3-: Fe3+ has 5 d electrons. Strong field CN- still forces pairing: t2g5 eg0 – one unpaired electron, paramagnetic.
[Fe(H2O)6]2+: Fe2+ with 6 d electrons, weak field H2O: t2g4 eg2 – four unpaired electrons, paramagnetic.

Scientific Explanation

Paramagnetism arises from the magnetic moment generated by unpaired electrons. Every electron has a spin magnetic moment, a quantum property that acts like a tiny bar magnet. When electrons are paired in an orbital, their spins are opposite (per the Pauli exclusion principle), so their magnetic moments cancel each other out. In a species with unpaired electrons, the unpaired electrons’ magnetic moments do not cancel, leaving a net magnetic moment for the atom or molecule.

When an external magnetic field is applied, these net magnetic moments align with the field, causing the material to be attracted to the field. This alignment is temporary – unlike ferromagnets, paramagnets lose their magnetization when the external field is removed, because thermal energy randomizes electron spin orientations at room temperature.

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The strength of paramagnetism is measured by the effective magnetic moment (μeff), calculated using the formula μeff = √[n(n+2)] Bohr magnetons, where n is the number of unpaired electrons. But 83 BM

n=3: μeff ≈ 3. Which means 87 BM
n=4: μeff ≈ 4. Think about it: 73 BM
n=2: μeff ≈ 2. On top of that, for example:
n=1: μeff ≈ 1. 90 BM
n=5: μeff ≈ 5.

This formula is a useful tool to confirm paramagnetism: if experimental μeff matches the calculated value for n unpaired electrons, the species is paramagnetic.

Common Practice Examples

Let’s apply the steps above to common variations of the which one of the following is paramagnetic question:

Example 1: Main Group Atoms and Ions

Question: Which one of the following is paramagnetic? A) Na+ B) Mg2+ C) Al3+ D) O2- E) B2

Step 1: Count total electrons: A) Na+ = 11 - 1 = 10 electrons (Ne configuration: 1s2 2s2 2p6 – all paired, diamagnetic) B) Mg2+ = 12 - 2 = 10 electrons (identical to Na+, diamagnetic) C) Al3+ = 13 - 3 = 10 electrons (identical to Na+, diamagnetic) D) O2- = 8 + 2 = 10 electrons (1s2 2s2 2p6 – all paired, diamagnetic) E) B2 = 5 + 5 = 10 total electrons (6 valence electrons). MO diagram for B2 fills σ2s2, σ*2s2, π2px1, π2py1 – two unpaired electrons, paramagnetic. Answer: E

Example 2: Transition Metal Ions

Question: Which one of the following is paramagnetic? A) Zn2+ B) Cu+ C) Ni2+ D) Ca2+ E) Sc3+

Step 1: Electron configurations: A) Zn2+: [Ar] 3d10 – full d shell, all paired, diamagnetic B) Cu+: [Ar] 3d10 – full d shell, diamagnetic C) Ni2+: [Ar] 3d8 – two unpaired electrons (5 d orbitals filled singly first, then 3 paired, leaving 2 unpaired), paramagnetic D) Ca2+: [Ar] – no d electrons, diamagnetic E) Sc3+: [Ar] – no d electrons, diamagnetic Answer: C

Example 3: Covalent Molecules

Question: Which one of the following is paramagnetic? A) N2 B) O2 C) F2 D) CO E) NO+

Step 1: MO diagrams: A) N2: 14 total electrons, 10 valence electrons, all paired, diamagnetic B) O2: 16 total electrons, 12 valence electrons, two unpaired electrons in π* orbitals, paramagnetic C) F2: 18 total electrons, 14 valence electrons, all paired, diamagnetic D) CO: 14 total electrons, 10 valence electrons, identical to N2, diamagnetic E) NO+: 14 total electrons (5 from N, 6 from O, minus 1 for positive charge), 10 valence electrons, identical to N2, diamagnetic Answer: B

FAQ

Is paramagnetic the same as ferromagnetic? No. Paramagnetic materials are weakly attracted to magnetic fields and lose magnetization when the field is removed. Ferromagnetic materials are strongly attracted, retain magnetization, and have aligned unpaired electrons in domains.
Can a species with no unpaired electrons be paramagnetic? No. Paramagnetism requires at least one unpaired electron. All electrons paired = diamagnetic, repelled by magnetic fields.
Why is O2 paramagnetic but N2 is not? O2 has 12 valence electrons, leaving two unpaired electrons in antibonding π* orbitals. N2 has 10 valence electrons, all paired in bonding orbitals.
How do ligands affect paramagnetism in transition metal complexes? Strong field ligands cause low spin complexes (electrons pair early, fewer unpaired electrons), weak field ligands cause high spin complexes (more unpaired electrons). Some complexes can be diamagnetic or paramagnetic depending on the ligand.
Is the paramagnetic question only for chemistry students? No. It appears in physics, materials science, and engineering assessments, as magnetic properties are relevant to semiconductor design, catalysis, and magnetic material development.

Conclusion

Answering which one of the following is paramagnetic relies on a systematic check for unpaired electrons, whether in isolated atoms, ions, or covalent molecules. Remember that paramagnetism is defined by at least one unpaired electron, driven by the net magnetic moment of unpaired electron spins. For main group species, use electron configuration and Hund’s rule. For molecules, apply molecular orbital theory. For transition metal complexes, account for ligand field splitting and high vs low spin states.

With practice, you can quickly eliminate diamagnetic options (full shells, all paired electrons) and identify the paramagnetic species in any list. Always verify with the number of unpaired electrons, and use the effective magnetic moment formula if you need to confirm the strength of paramagnetism.