[Fe(CN)6]³⁻ is the ferricyanide ion, an octahedral coordination compound with Fe in the +3 oxidation state bound to six cyanide ligands. In Inorganic Chemistry I, it shows up in naming, ligand field, and redox work.
[Fe(CN)6]³⁻ is the ferricyanide ion, a coordination complex with one iron center surrounded by six cyanide ligands. In Inorganic Chemistry I, you usually meet it as an octahedral complex, where the six ligands sit around Fe at the corners of an octahedron.
The key thing to notice is that the cyanide ligands are attached through their carbon end and act as strong field ligands. That means they cause a large splitting of the iron d orbitals. For Fe(III), that splitting changes how electrons arrange themselves, which affects color, magnetism, and reactivity. This is one of the clearest examples of how ligand identity changes the behavior of a metal ion.
The iron oxidation state in [Fe(CN)6]³⁻ is +3. You find that by treating each CN− as a -1 ligand and using the overall -3 charge on the complex. That oxidation state matters because it connects the formula to electron count, naming, and the related reduced form, [Fe(CN)6]⁴⁻, which is ferrocyanide.
This ion is also a good example of a coordination sphere. The metal plus the six directly bonded ligands make up the inner complex, while the overall charge tells you the species is an anion. In problem sets, that means you may be asked to name it, find the oxidation state, identify the geometry, or predict whether it is high spin or low spin. With cyanide, the low-spin description is the one that fits.
You will also see ferricyanide in redox chemistry, because it can gain an electron and convert to ferrocyanide. That makes the pair useful for electrochemistry and for spotting how oxidation state changes show up in formulas. The name, structure, and electron behavior are all tied together here, which is why this complex shows up so often in coordination chemistry units.
[Fe(CN)6]³⁻ matters because it pulls together several coordination chemistry ideas in one formula. You can use it to practice oxidation state, ligand field theory, octahedral geometry, and the difference between a metal ion by itself and a full coordination complex.
It is also a clean example of how a strong-field ligand changes electron placement at a d-block element. That lets you connect structure to observable properties like color and magnetic behavior instead of treating formulas like memorization only.
Ferricyanide also helps you see why naming and formula reading matter in inorganic chemistry. The difference between ferricyanide and ferrocyanide is just one electron, but that small change flips the oxidation state and changes the complex’s redox behavior. In lab or homework, that kind of detail can decide whether you identify a species correctly or miss the chemistry entirely.
When you understand this ion, you are better prepared for coordination compound questions that ask you to move from formula to geometry to electron count to reactivity in a single step.
Keep studying Inorganic Chemistry I Unit 8
Visual cheatsheet
view galleryCoordination Compound
[Fe(CN)6]³⁻ is a coordination compound because the iron center is bonded to ligands through coordinate interactions. That lets you practice the basic idea of a metal complex, not just a lone ion. When you see this formula, you should separate the central metal, the ligands, and the overall charge before doing any deeper analysis.
Ligand
Cyanide is the ligand in ferricyanide, and it is a strong field ligand. That matters because ligand strength affects d-orbital splitting, spin state, and magnetic properties. In this complex, the ligand identity is not just a label, it explains why the ion behaves differently from a weak-field octahedral complex.
Oxidation State
The iron oxidation state in [Fe(CN)6]³⁻ is +3, which you determine from the ligand charges and the total charge on the complex. This is a standard coordination chemistry skill because oxidation state helps you name the compound, compare it to related species, and track redox changes.
Coordination Number 6
This ion has coordination number 6 because six cyanide ligands bind directly to the metal center. A coordination number of 6 usually gives octahedral geometry, so this term connects the formula to the 3D arrangement you are expected to recognize in problems and models.
A quiz or problem set might ask you to name [Fe(CN)6]³⁻, find the oxidation state of Fe, or identify the geometry and spin state. The move is to count the ligands, use the CN− charge to solve for iron, and then connect six ligands to octahedral geometry. If the question includes redox, you may also compare it to [Fe(CN)6]⁴⁻ and explain which ion is oxidized or reduced. In a lab or discussion, ferricyanide often shows up in electrochemistry or as a visible example of color change tied to ligand field effects. The best answers show the chain from formula to structure to electron behavior, not just the name.
[Fe(CN)6]³⁻ is ferricyanide, while [Fe(CN)6]⁴⁻ is ferrocyanide. They have the same ligands and the same octahedral shape, but they differ by one electron and one unit of charge. That means the iron oxidation state changes from +3 in ferricyanide to +2 in ferrocyanide, which is why the pair shows up in redox questions.
[Fe(CN)6]³⁻ is the ferricyanide ion, an octahedral coordination complex with six cyanide ligands around iron.
The iron in this complex has oxidation state +3, which you can work out from the -1 charge on each cyanide ligand and the overall -3 charge.
Cyanide is a strong field ligand, so ferricyanide is a good example of how ligand strength affects d-orbital splitting, spin state, and color.
This ion is often studied alongside [Fe(CN)6]⁴⁻, because the two species form a classic redox pair.
In Inorganic Chemistry I, you use ferricyanide to practice naming, geometry, coordination number, and electron-count reasoning together.
It is the ferricyanide ion, a coordination complex with Fe(III) in an octahedral arrangement of six cyanide ligands. In class, it is a standard example for oxidation state, ligand field, and redox questions.
Yes. Six ligands around one metal center usually give octahedral geometry, and that is the arrangement for ferricyanide. This is one of the most common geometry patterns you will see in coordination chemistry.
Each cyanide ligand is -1, so six of them contribute -6 total. Since the whole complex is -3, iron must be +3 to make the charges add up correctly.
Ferricyanide is [Fe(CN)6]³⁻ and ferrocyanide is [Fe(CN)6]⁴⁻. They have the same ligand set and geometry, but different oxidation states and charges, so they behave differently in redox chemistry.