Paramagnetic vs diamagnetic behavior is determined by the presence or absence of unpaired electrons. A species with at least one unpaired electron is paramagnetic (net magnetic moment, attracted to a field), while a species with all electrons paired is diamagnetic (no permanent moment, weakly repelled).
Electronic origin of magnetic response
Electron spin produces a magnetic moment. In an atom, ion, or molecule, paired electrons occupy the same orbital with opposite spins, so their spin moments cancel. Unpaired electrons leave a nonzero spin contribution, producing a net moment that can align partially with an external magnetic field.
Core criterionelectron-count test
At least one unpaired electron implies paramagnetism; complete pairing implies diamagnetism.
Prediction from orbital filling and electron configuration
Orbital diagrams and electron configurations provide the fastest prediction method. Hund’s rule places electrons singly (with parallel spins) in degenerate orbitals before pairing begins. As a result, partially filled subshells often generate unpaired electrons and therefore paramagnetism.
For many main-group species, the conclusion follows directly from whether the highest-energy subshell is fully filled. For transition-metal ions, the d-electron count and the ligand field can alter pairing patterns, so both oxidation state and coordination environment can matter.
Magnetic moment and the role of unpaired electrons
The magnitude of the magnetic moment increases with the number of unpaired electrons. A commonly used estimate (spin-only) is \[ \mu_{\text{so}} \approx \sqrt{n(n+2)} \;\; \mu_B, \] where \(n\) is the number of unpaired electrons and \(\mu_B\) is the Bohr magneton. The qualitative classification (paramagnetic vs diamagnetic) only requires deciding whether \(n\) is zero or nonzero.
Representative examples in general chemistry
| Species | Paramagnetic or diamagnetic? | Electron-based justification |
|---|---|---|
| O2 | Paramagnetic | Two unpaired electrons in the highest-occupied molecular orbitals (MO picture); the bond order is 2, but the spins do not fully pair. |
| N2 | Diamagnetic | All electrons paired in the occupied molecular orbitals (closed-shell molecule). |
| Ne | Diamagnetic | Closed-shell electron configuration; all orbitals filled with paired electrons. |
| Al atom | Paramagnetic | Ground-state configuration ends with 3p1, giving one unpaired electron. |
| Fe3+ (typical high-spin) | Paramagnetic | d5 often contains multiple unpaired electrons; many complexes show strong paramagnetism. |
| Zn2+ | Diamagnetic | d10 configuration is fully filled; no unpaired electrons. |
Substance-level observations and terminology
Paramagnetic materials are attracted into stronger-field regions and exhibit a small positive magnetic susceptibility. Diamagnetic materials are weakly repelled and have a small negative susceptibility. In ordinary laboratory fields, diamagnetism is subtle; paramagnetism is typically more noticeable when multiple unpaired electrons are present.
Common pitfalls
A common mistake is equating “odd number of electrons” with paramagnetism. The decisive quantity is not odd or even electron count but the presence of unpaired electrons after the correct electron configuration or MO filling is established. Another mistake is assuming all transition-metal ions are paramagnetic; d0 and d10 cases are diamagnetic, and intermediate cases can become low-spin (more paired) depending on ligand field strength.