[Fe(CN)6]4- is a coordination complex with an iron(II) ion bonded to six cyanide ligands. In Intro to Chemistry, it is a classic example of octahedral bonding, oxidation state, and low-spin behavior.
[Fe(CN)6]4- is the hexacyanoferrate(II) ion, a coordination complex in Intro to Chemistry where one Fe2+ center is surrounded by six CN- ligands. It is a good example of how a metal ion and ligands can form a stable, charged species with its own structure and properties.
The iron has an oxidation state of +2, and each cyanide ligand is monodentate, meaning it attaches through one bonding site. Since there are six ligands around the metal, the coordination number is 6. That arrangement gives the complex an octahedral geometry, not a simple ionic cluster or a random collection of atoms.
What makes this term worth knowing is that the bond between the metal and each ligand is a coordinate covalent bond. The ligand donates both electrons in the bond to the metal center. In cyanide, the lone pair on carbon is the donor site, so CN- acts as a strong-field ligand in many general chemistry treatments.
Because CN- is a strong-field ligand, [Fe(CN)6]4- is usually described as a low-spin complex. For iron(II), that means the d electrons pair up in the lower-energy d orbitals instead of spreading out as much as they would in a weak-field case. In plain terms, the ligand field is strong enough to change how the electrons arrange themselves.
That electron arrangement matters because it changes magnetic behavior and sometimes color and stability. So when you see [Fe(CN)6]4-, do not treat it like a formula to memorize by sight only. It is a compact example of the whole coordination chemistry idea: metal ion, ligands, oxidation state, geometry, coordination number, and electron pairing all working together.
[Fe(CN)6]4- shows up right where Intro to Chemistry moves from basic bonding into transition metal chemistry. If you can read this ion correctly, you can handle a lot of coordination-compound questions without getting lost in the notation.
It connects several ideas at once. You identify the central metal, count ligands, find the oxidation state, and connect that to geometry and electron arrangement. That same chain of reasoning shows up again and again with other complexes, especially when you compare one ligand set to another.
It also gives you a clean example of why ligands are not all equal. Cyanide is not just a placeholder letter group. It changes the splitting of the d orbitals, which changes whether the complex is low spin or high spin. That is the kind of detail that turns a formula into chemistry.
In labs or problem sets, this kind of ion often appears in naming questions, structure questions, or electron-counting practice. If you can explain why the complex is octahedral and how the +2 oxidation state is found, you are doing more than memorizing a formula. You are showing that you can read a coordination compound as a system.
Keep studying Intro to Chemistry Unit 19
Visual cheatsheet
view galleryCoordination Complex
This ion is a coordination complex because it has a central metal ion bonded to surrounding ligands. The whole formula is one charged unit, not separate iron and cyanide ions floating independently. When you identify [Fe(CN)6]4-, you are really practicing how to read and analyze the full coordination sphere.
Ligand
CN- is the ligand in [Fe(CN)6]4-. It donates a lone pair to the iron center through one atom, so it counts as monodentate. This matters because the type of ligand changes the geometry, electron arrangement, and sometimes the magnetic behavior of the complex.
Oxidation State
The iron in [Fe(CN)6]4- is +2, and you find that by combining the overall charge of the complex with the charge on each cyanide ligand. This is a standard chemistry move for coordination compounds. Once you know the oxidation state, you can make sense of the d electron count and spin behavior.
Coordination Number
The coordination number here is 6 because six cyanide ligands are attached to the iron center. That number helps determine the shape of the complex, and for six ligands the most common arrangement is octahedral. Counting coordination number is one of the first steps in analyzing any coordination ion.
A quiz or problem set might give you [Fe(CN)6]4- and ask you to find the oxidation state of iron, identify the ligand, or predict the geometry. You might also be asked to explain why it is low spin or to count the coordination number from the formula.
If you get a naming question, you need to read the brackets as one complex ion and not split the formula apart. If you get a comparison question, you may need to tell the difference between this strong-field complex and a similar one with a weaker ligand set. On a lab or discussion question, you could be asked to connect the electron arrangement to color or magnetic behavior.
These are both coordination complexes, but they differ in the metal, the ligand, and the geometry. [Fe(CN)6]4- has six cyanide ligands around iron(II), while [Cu(NH3)4]2+ has four ammonia ligands around copper(II). When you compare them, look at the charge, coordination number, and whether the ligands produce octahedral or square planar-type arrangements.
[Fe(CN)6]4- is the hexacyanoferrate(II) ion, a coordination complex built from Fe2+ and six CN- ligands.
The coordination number is 6, so the complex is octahedral in shape.
CN- is a monodentate ligand, meaning each cyanide binds through one donor atom.
The iron oxidation state is +2, which helps you determine the d electron count and spin behavior.
This ion is a classic low-spin example in Intro to Chemistry because cyanide is a strong-field ligand.
It is a coordination complex made of one iron(II) ion surrounded by six cyanide ligands. You use it to study oxidation state, coordination number, and octahedral geometry. It is also a common example of a low-spin metal complex.
It is octahedral because there are six ligands arranged around the iron center. A tetrahedral complex usually has four ligands, so the coordination number tells you this one is not tetrahedral. That is a common place to make a mistake if you only look at the charge and not the ligand count.
Each cyanide ligand has a -1 charge, so six of them contribute -6 total. Since the whole complex is -4, the iron must be +2 to make the charges add up correctly. That oxidation state is the starting point for d electron counting.
Cyanide is a strong-field ligand, so it creates a large splitting between the d orbitals. That encourages electrons to pair in the lower-energy orbitals instead of staying unpaired in higher ones. In Intro to Chemistry, this is one of the clearest examples of ligand field effects.