Which One Of The Following Is Paramagnetic

Which One of the Following Is Paramagnetic? A Plain-English Guide

You probably remember the word "paramagnetic" from a chemistry class that felt like a lifetime ago. The good news: the concept itself is way less scary than it sounded under fluorescent classroom lights. In real terms, once you understand what paramagnetism actually means — and what it doesn't* mean — those multiple-choice questions become almost easy to handle. Let's walk through it.

What "Paramagnetic" Actually Means

Atoms and molecules have electrons, and those electrons behave like tiny magnets because they spin and orbit. Most of the time, electrons pair up, and their tiny magnetic fields cancel each other out. The result? The atom or molecule as a whole isn't magnetic. Easy to understand, harder to ignore.

Paramagnetic substances are the opposite story. They have unpaired electrons — solo electrons that aren't paired with another one. Because of that, those solo electrons each carry a small magnetic moment that doesn't get canceled, so the whole atom or molecule has a net magnetic field. When you put a paramagnetic material near an external magnetic field, it gets weakly pulled in. Take the field away, and the attraction disappears. In practice, no memory. No permanent magnetization.

This is different from ferromagnetism, which is what iron and a few other metals do — they stay magnetized after the external field is gone. Most of the stuff around you (wood, plastic, water, your hand) is diamagnetic. And it's different from diamagnetism, which is the very weak repulsion you get from materials where all electrons are paired. The paramagnetic cases sit in a smaller, more interesting middle ground.

Unpaired Electrons, In Plain Language

Every orbital holds up to two electrons, and those two electrons spin in opposite directions so their magnetic contributions cancel. Which means the moment you have an orbital with just one electron in it, you have a net magnetic moment. That single unpaired electron is the entire reason a substance is paramagnetic.

This is why counting unpaired electrons is the single most useful thing you can do when you see a paramagnetism question. You don't have to memorize a long list. You just count.

Why It Matters (Beyond the Test)

Look, I get it. But paramagnetism isn't just a textbook trick. Now, if you're here, there's a good chance you're studying for an exam. It's a real thing with real consequences.

Oxygen gas, the stuff you're breathing right now, is paramagnetic. That's not trivia — it's a clue about how oxygen behaves in magnetic fields, and it has implications in fields from medical imaging to industrial chemistry. Because of that, transition metal complexes often have unpaired electrons, which is partly why they're so colorful and so reactive as catalysts. The magnetic properties of a material tell you something about its electronic structure, and that electronic structure determines how it behaves chemically. Worth knowing.

So the question "which of the following is paramagnetic?" is really asking: which of these has unpaired electrons?* Once you see it that way, the multiple-choice list is just a puzzle.

How to Solve "Which of the Following Is Paramagnetic" Questions

Here's the step-by-step approach. It's the same every time, and once you've done it two or three times, you'll fly through it.

Step 1: Write Out the Electron Configuration

Take the species in question — it might be a single atom, an ion, or a molecule — and figure out where the electrons go. For main group elements, this is usually quick. For transition metals, you'll need to fill the d-orbitals.

Step 2: Look at the Valence (or d) Orbitals

The unpaired electrons almost always show up either in the outermost s and p orbitals (for main group) or in the d orbitals (for transition metals). That said, count carefully. This is the step where most mistakes happen — people miscount because they forget an electron was removed or added when an ion formed.

Step 3: Count the Unpaired Electrons

Zero unpaired electrons? Diamagnetic. One or more? In real terms, paramagnetic. That's the whole rule.

Step 4: Watch For Ions

Here's the trap that catches a lot of people. Sodium atom (Na) has one unpaired electron in its 3s orbital, so it's paramagnetic. Sodium ion (Na⁺) has lost that electron, so all its remaining electrons are paired, and it's diamagnetic. Here's the thing — a neutral atom and its ion often have different numbers of unpaired electrons. Same element, opposite magnetic behavior, depending on charge.

Common Ion Examples You'll See

A few species show up over and over on these questions. Worth knowing:

• Cu⁺ is diamagnetic (3d¹⁰, all paired). Cu²⁺ is paramagnetic (3d⁹, one unpaired electron). Easy to mix up if you forget the charge.
• Fe²⁺ is paramagnetic (3d⁶, four unpaired electrons in the high-spin configuration). Zn²⁺ is diamagnetic (3d¹⁰). Iron and zinc sit right next to each other on the periodic table, but their ions behave completely differently.
• O₂ is one of the most famous paramagnetic molecules. Both oxygen atoms have unpaired electrons, and when they bond, two of those unpaired electrons end up in separate antibonding orbitals. So molecular oxygen is paramagnetic, which surprises a lot of students who assume "stable diatomic gas" means "boring magnetically." It isn't.
• N₂ is diamagnetic. All electrons paired, no net magnetic moment. So even though O₂ and N₂ look like cousins, they behave very differently in a magnetic field.

What Most People Get Wrong

Honestly? The mistakes are almost always about counting, not concept.

Forgetting about ions. A question might list "Fe" and "Fe²⁺" as separate options, and people treat them as if they had the same electron count. They don't. Always re-derive the configuration for the actual species given.

Confusing the rule for molecules with the rule for atoms. The concept is the same — count unpaired electrons — but with molecules you sometimes have to look at molecular orbital diagrams instead of atomic orbitals. The classic example is O₂ versus N₂. If you only look at Lewis structures, both look fully paired. The molecular orbital picture tells the real story.

Assuming transition metals are always paramagnetic. Most are, but not all. Cu⁺, Zn²⁺, Ag⁺, and a few others have full d-subshells and zero unpaired electrons. Don't pattern-match. Count.

Forgetting that "more paramagnetic" is a thing. Some questions don't ask whether* something is paramagnetic — they ask which option is most* strongly paramagnetic. That depends on how many unpaired electrons there are. A d⁵ high-spin ion like Fe³⁺ has five unpaired electrons and is more strongly paramagnetic than a d¹ ion with only one. Both are paramagnetic, but they're not equally so.

Practical Tips That Actually Help

If you want to get these questions right consistently, a few habits go a long way. Simple, but easy to overlook.

Draw the orbital diagram. Even a quick sketch on scratch paper beats mental counting every time. Boxes for orbitals, half-arrows for electrons, and you can literally see the unpaired ones.

Memorize the d-orbital filling order for the common ions. You don't need all of them, but d⁰ through d¹⁰ for the 3d row is well worth knowing. Once you've got those, the 4d row follows the same pattern.

Watch the question's wording. "Is paramagnetic" and "is most paramagnetic" are different questions. So is "is not paramagnetic," which is really asking about diamagnetic species. Read carefully.

Use the periodic table as a shortcut. Group numbers tell you how many valence electrons an atom has, and that gives you a head start on counting. Chromium and copper are the two 3d exceptions that trip people up — Cr is 3d⁵4s¹, not 3d⁴4s², and Cu is 3d¹⁰4s¹, not 3d⁹4s². Worth memorizing.

When in doubt, oxygen is a giveaway. If O₂ appears as an option, it's paramagnetic. If you're stuck between choices and O₂ is one of them, that's a strong hint.

FAQ

Is every atom with an odd number of electrons paramagnetic?

Yes. If an atom has an odd number of electrons, at least one of them must be unpaired. So nitrogen, hydrogen, sodium atoms, and chlorine atoms are all paramagnetic.

Is every atom with an odd number of electrons paramagnetic?
Yes – an atom that possesses an odd total number of electrons must have at least one unpaired electron, because electrons pair up in orbitals. Nitrogen (7 e⁻), hydrogen (1 e⁻), sodium (11 e⁻) and a chlorine atom (17 e⁻) all fall into this category. The catch is that many atoms appear in compounds or ions where the electron count changes. A sodium ion (Na⁺) has lost its single 3s electron, leaving an even‑electron configuration (10 e⁻) and is therefore diamagnetic, even though the neutral Na atom is paramagnetic. So always evaluate the species you are asked about, not just the neutral element.

Want to learn more? We recommend the class with the greatest relative frequency is and do you eat apples in spanish for further reading.

Frequently asked follow‑ups

Do ions follow the same rule as neutral atoms?
Absolutely. The presence or absence of unpaired electrons depends on the total electron count of the ion. Here's one way to look at it: Fe²⁺ has a 3d⁶ configuration (four unpaired electrons in the high‑spin case) and is paramagnetic, whereas Zn²⁺ (3d¹⁰) has all electrons paired and is diamagnetic. When you move from a neutral atom to its ions, re‑count the electrons from scratch.

How do polyatomic molecules and ions behave?
You must look at the overall electron count of the molecule, not just the sum of the atomic valence electrons. Molecular‑orbital diagrams are the most reliable tool. O₂ is a textbook case: its MO diagram shows two electrons in degenerate π* orbitals, giving it two unpaired electrons (triplet ground state). In contrast, CO has a bond order of three and all electrons are paired, making it diamagnetic.

Can a species be both paramagnetic and diamagnetic?
No. A substance’s magnetic response is dominated by the strongest interaction present. If any unpaired electrons exist, the paramagnetic contribution (positive susceptibility) overwhelms the weak temperature‑independent diamagnetic contribution. In practice you either classify a species as paramagnetic (has unpaired electrons) or diamagnetic (all electrons paired).

Why is paramagnetism measured in Bohr magnetons?
The magnetic moment of an electron spin is approximately one Bohr magneton (µ_B). For a given number of unpaired electrons, the spin‑only magnetic moment can be estimated with the “spin‑only” formula:

[ \mu_{so} = \sqrt{n(n+2)},\mu_B ]

where n is the number of unpaired electrons. This provides a quick way to compare expected paramagnetic strength, which is useful in coordination chemistry and material science.

What about temperature?
Paramagnetic susceptibility follows the Curie or Curie‑Weiss law (χ = C/T) and grows stronger as temperature drops, because thermal agitation less effectively randomizes the magnetic moments. Diamagnetism is essentially temperature‑independent. In practice, this means

In practice, this means that when you are trying to determine the magnetic behavior of a compound you must first identify the exact species you are dealing with and then apply the appropriate temperature regime for the measurement. Now, if a sample shows a χ vs T trend that follows the Curie–Weiss law (χ = C⁄(T – θ)), the dominant paramagnetic contribution will grow as the temperature is lowered, and any diamagnetic background becomes a minor correction. Conversely, for a purely diamagnetic material, χ remains essentially constant over a wide temperature range, so any observed temperature dependence usually signals the presence of trace paramagnetic impurities or instrumental artifacts.

Experimental determination of magnetic susceptibility

When interpreting data, remember to subtract the diamagnetic contribution of the sample holder, ligands, and counter‑ions. For transition‑metal complexes, the spin‑only formula μ_so = √[n(n + 2)] μ_B gives a quick estimate of the expected moment, but crystal‑field splitting, spin‑orbit coupling, and orbital contributions can cause deviations of 10–30 %.

Why the rule matters in real‑world contexts

1. Medical imaging – Gadolinium‑based MRI contrast agents are deliberately paramagnetic (typically 7–8 unpaired electrons) because the large magnetic

moment of Gd³⁺ shortens the T₁ relaxation time of nearby water protons, producing bright signal enhancement. The very high χ of these agents is precisely what makes them effective at the millimolar concentrations used clinically. If a contrast agent were diamagnetic, it would produce no contrast at all, rendering the diagnostic test useless.

2. Materials science – The growing field of spintronics depends on materials with strong, tunable magnetic responses. Ferromagnetic metals such as Fe, Co, and Ni (which are technically a special class of paramagnetism with cooperative ordering) are chosen because they retain a net magnetization even after the external field is removed. In contrast, diamagnetic materials like bismuth or pyrolytic carbon are employed where magnetic repulsion is desired—for example, in levitation experiments or in non-magnetic cryogenic housings where any stray susceptibility would distort sensitive measurements.

3. Geophysics and environmental science – Susceptibility measurements of rock cores provide rapid, non-destructive proxies for mineral composition. Since magnetite (Fe₃O₄) is strongly ferromagnetic, even trace amounts dominate the bulk susceptibility, allowing geologists to map paleomagnetic directions and reconstruct past continental movements. Conversely, diamagnetic minerals such as quartz and calcite serve as reference standards for background corrections.

4. Quality control in chemistry – In coordination chemistry, confirming the spin state of a complex is essential. A measured μ_eff that matches the spin‑only value confirms a high‑spin d⁵ or d⁶ configuration, while a lower value suggests low‑spin pairing or strong ligand-field stabilization. Failing to account for the diamagnetic ligand shell would lead to an underestimation of μ_eff and potentially a misassignment of oxidation state.

Common pitfalls and how to avoid them

• Ignoring temperature dependence: Reporting a single χ value without specifying the temperature can be misleading. Always state T (and the applied field H) when publishing susceptibility data.
• Overlooking impurity contributions: A nominally diamagnetic sample showing a slight positive χ at low T may contain ppm‑level paramagnetic impurities. Use high‑purity starting materials and consider magnetic separation or recrystallization if precise diamagnetic corrections are critical.
• Misidentifying the magnetic regime: Assuming Curie‑Weiss behavior at temperatures near a phase transition (e.g., near T_C for a ferromagnet) will produce erroneous θ values. Map the full χ(T) curve to detect anomalies such as λ‑peaks or divergence.
• Using incorrect units: Magnetic susceptibility is dimensionless in SI (χ_SI) but has units of emu mol⁻¹ in the CGS system. Mixing the two is a frequent source of error; always convert explicitly and report which convention is used.

Future directions

Advances in instrumentation are pushing the limits of sensitivity ever lower. Because of that, nitrogen‑vacancy (NV) center magnetometry now allows spatial mapping of χ with sub‑micrometer resolution, opening the door to imaging magnetic domains in individual nanoparticles. Simultaneously, machine‑learning algorithms are being trained to predict χ from structural descriptors, accelerating the discovery of new paramagnetic materials for quantum information processing and high‑density magnetic storage.

Worth adding, the emerging concept of “magnetic transparency” in topological insulators and Dirac semimetals challenges the simple paramagnetic/diamagnetic dichotomy. In these systems, the magnetic response is governed by Berry curvature rather than localized moments, and the sign of χ can be tuned by carrier density—an effect that cannot be captured by the classical Curie–Weiss framework.

Conclusion

Magnetic susceptibility is a powerful, versatile probe that translates the invisible arrangement of electrons into a measurable quantity. Consider this: whether you are synthesizing a new coordination complex, calibrating an MRI contrast agent, or analyzing the magnetic fabric of a rock, the core principle remains the same: **paramagnetic materials attract, diamagnetic materials repel, and the strength of that response is dictated by both the nature of the electronic structure and the temperature at which it is measured. ** By mastering the underlying theory, selecting the appropriate experimental technique, and diligently correcting for diamagnetic backgrounds, scientists across chemistry, physics, biology, and earth sciences can extract reliable magnetic information that informs everything from fundamental research to everyday technology.