[Fe(CN)6]^{3-} is the hexacyanoferrate(III) complex ion, an octahedral Fe(III) coordination compound with six cyanide ligands. In Inorganic Chemistry I, it is a classic low-spin example for magnetic behavior.
[Fe(CN)6]^{3-} is an octahedral coordination complex in Inorganic Chemistry I with one iron center surrounded by six cyanide ligands. The usual name is hexacyanoferrate(III), which tells you the metal is iron in the +3 oxidation state and that there are six CN- ligands around it.
The structure matters because six ligands arranged symmetrically around a metal ion produce an octahedral crystal field. In that environment, the metal's d orbitals split into two energy groups, and the size of that splitting depends on the ligand. Cyanide is a strong-field ligand, so it creates a large gap between the lower-energy and higher-energy d orbitals.
That large splitting pushes electrons to pair in the lower set before they occupy the higher set. For Fe(III), that usually gives a low-spin arrangement with no unpaired electrons, which makes the complex diamagnetic. If you are working a coordination problem, this is the kind of complex where the ligand field strength directly changes the electron configuration you predict.
A common mistake is to look only at the metal and forget the ligand. Iron(III) by itself does not tell you the magnetic behavior, because the same metal ion can be high-spin or low-spin depending on the ligand environment. With CN-, the strong field effect wins, so [Fe(CN)6]^{3-} becomes a standard example of electron pairing in an octahedral complex.
You will also see this complex used as a comparison point. Once you know its electron count, geometry, and spin state, it becomes easier to compare it with weaker-field complexes that keep unpaired electrons and show paramagnetism instead.
[Fe(CN)6]^{3-} is one of the cleanest examples for connecting structure to magnetic behavior in coordination chemistry. If you can explain why this ion is low-spin and usually diamagnetic, you can do the same reasoning for many other octahedral complexes.
It also gives you a concrete way to use crystal field theory instead of just memorizing the ligand list. You identify the oxidation state, count the d electrons, decide whether the ligand is strong-field or weak-field, and then predict whether electrons pair or remain unpaired. That chain of reasoning shows up again and again in problem sets on magnetic moments, electron configurations, and octahedral geometry.
This complex is especially useful when comparing magnetic data to predicted structures. If a class asks why one iron complex is attracted to a magnetic field while another is not, [Fe(CN)6]^{3-} is a good anchor example because its behavior is tied directly to the large crystal field splitting caused by cyanide.
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Visual cheatsheet
view galleryParamagnetism
Paramagnetism is the opposite magnetic outcome from what you get with [Fe(CN)6]^{3-}. If a coordination complex has unpaired electrons, it is attracted to a magnetic field and shows paramagnetic behavior. This term helps you compare strong-field complexes like this one with weaker-field complexes that keep electrons unpaired.
Diamagnetism
[Fe(CN)6]^{3-} is typically diamagnetic because its electrons pair up in the lower-energy d orbitals. Diamagnetism is the label you use when all electrons are paired and the compound is weakly repelled by a magnetic field. This is the property most often tied to this ion in coordination chemistry problems.
Crystal Field Theory
Crystal Field Theory explains why cyanide causes a large splitting of the d orbitals in [Fe(CN)6]^{3-}. The theory gives you the framework for predicting whether electrons will pair in a low-spin pattern or spread out in a high-spin pattern. That prediction is what links the formula to magnetic behavior.
d-orbitals
The electron arrangement in [Fe(CN)6]^{3-} is all about how the metal's d orbitals split in an octahedral field. You use the d-orbital picture to decide where each electron goes before you talk about spin state or magnetic moment. Without that orbital splitting, the diamagnetic result would not make sense.
A quiz problem might give you [Fe(CN)6]^{3-} and ask for the oxidation state, geometry, electron count, or magnetic behavior. The move is to identify Fe as +3, recognize cyanide as a strong-field ligand, and then predict low spin with paired electrons. If the question gives a comparison set, you use this complex as the diamagnetic benchmark and explain why a weaker-field complex would behave differently.
In a problem set, you may also be asked to sketch the octahedral d-orbital splitting pattern and place the electrons in the lower-energy orbitals first. If magnetic data are included, you match the observed behavior to the predicted spin state instead of guessing from the metal alone. That is the main skill this term supports: turning a formula into a structure and a magnetic prediction.
[Fe(CN)6]^{3-} and [Fe(CN)6]^{4-} look almost identical, but they differ in the iron oxidation state and electron count. That changes the magnetic prediction you make. The +3 complex is the ferric form, while the +2 complex is the ferrous form, so you should always check the charge before assigning spin state.
[Fe(CN)6]^{3-} is the octahedral hexacyanoferrate(III) complex, with Fe3+ surrounded by six cyanide ligands.
Cyanide is a strong-field ligand, so it creates a large octahedral splitting of the metal d orbitals.
Because the splitting is large, electrons pair in the lower-energy orbitals and the complex is usually low spin.
Low-spin behavior makes [Fe(CN)6]^{3-} typically diamagnetic, which is why it is a standard example in magnetic property questions.
When you see this ion, think oxidation state, ligand strength, d-orbital splitting, and unpaired electrons in that order.
[Fe(CN)6]^{3-} is hexacyanoferrate(III), an octahedral coordination complex with Fe3+ at the center and six cyanide ligands around it. In Inorganic Chemistry I, it is used as a classic example of a strong-field, low-spin complex.
It is typically diamagnetic because cyanide causes a large enough splitting that the d electrons pair up. Since there are no unpaired electrons in the usual low-spin arrangement, the complex is weakly repelled by a magnetic field.
Cyanide is a strong-field ligand, so it produces a large octahedral crystal field splitting. That makes it energetically better for electrons to pair in the lower-energy d orbitals than to occupy the higher set.
Check the charge and the iron oxidation state. The 3- complex has Fe(III), while the 4- complex has Fe(II), so they do not have the same electron count or the same magnetic behavior.