Why don't electrons pair in [Fe(NH3)6]2+?

Why doesn't electron pairing take place in $\ce{[Fe(NH3)6]^2+}$ though $\ce{NH3}$ is a strong ligand field producer?




(If you know the basics, go to last paragraph and skip the introduction part)


There are two types of ligands. Strong field ligands and weak field ligands. Those which only cause a small extent of crystal field splitting is called weak field ligand and those which cause large splitting is called strong field ligand.


***Why does crystal field splitting take place?***


Let's take an octahedral complex. The $d\_{x^2-y^2}$ and $d\_{z^2}$ point along x,y,z axis and the $t\_{2g}$ orbitals point between the axis. When the 6 ligands approach along the three axes, the energy of the $e\_g$ orbitals increase more than the increase in the energy of $t\_{2g}$ orbitals. Hence the d-orbital splits into two. Filling each electron in $t\_{2g}$ orbitals decrease energy by $0.4\Delta\_0$ and filling electrons in $e\_g$ increase energy by $0.6\Delta\_0$


Coming back to your question, you know how electrons are filled in the orbitals (according to hund's rule of course). But how does electron pairing take place in your case?


Let the energy required to pair an electron be P.
Let's say there are 4 electrons. The two arrangements are one electron in each sub-orbital (no pairing) and 4 electrons all in $t\_{2g}$ itself (instead of 4th electron going to $e\_g$, its paired up in $t\_{2g}$).
In the first case, the total energy is $-0.6\Delta\_0$ and in second case, its $-1.6\Delta\_0+P$. So, only strong field ligands can form the second arrangement. $P$ value is constant for a given metal ion.


Finally, ammonia is weak field ligand but in certain cases, it acts as a strong field ligand (example: for Cobalt). For Fe, ammonia is a weak field ligand. Hence no pairing takes place.


 10

The dichotomy between weak- and strong-field ligands that I have often seen constructed here is *false* and physically inaccurate.


Complexes can be in a low-spin state or in a high-spin state. The high-spin state is the one according to Hund’s rule where electrons are filled into the five d orbitals one after the other. A low-spin state is any complex in which the Hund rule is not strictly observed, but some orbitals are filled with two electrons while others remain empty. Classic examples for high-spin complexes are, e.g. $\ce{[FeCl6]^3-, [CoCl4]^2-}$ and others. Examples for low-spin complexes are, e.g. $\ce{[Fe(CN)6]^3-, [Cr(CN)6]^3-}$ and others.


While these examples may seem to exemplify your stated dichotomy (with chloride being on the weak-field side and cyanide being on the high-field side), reality is much more complex and the spin state of a complex depends on ligands, metal, oxidation state, counterions and so on. In fact, for 2-picolylamine ($\ce{2{-}pic}$) as a ligand to iron(II), both the low-spin and the high-spin $\ce{[Fe(2{-}pic)3]^2+}$ complexes have been isolated and characterised by crystal structures. The $\ce{Fe-N}$ bond length differs by $19~\mathrm{pm}$ with the high-spin complex obviously having longer bonds.


Complexes have a tendency to be in the low-spin state if


* the metal is not a 3d metal
* the metal’s oxidation state is higher
* metal to ligand π-backbonding can take place.


Especially that final point is important. Ammonia is very bad at π backbonding (and amide is even π-*donating*). So calling it a high-field ligand is misappreciating physicochemical reality. It is just ammonia’s strong σ bonding ability that causes a slightly larger field split than water. **As such, ammonia is best termed a medium-field ligand.** (The same is true for most neutral nitrogen donors.)


 8

$\ce{NH3}$ is placed in the middle of the spectrochemical series.
It mostly acts as SFL but sometimes may act as WFL.
When a metal ion having higher charge is interacting with ligand NH3 it acts as a SFL.
When the metal ion has lower charge it acts as WFL.


EG: $\ce{Co^2+}$ and $\ce{Co^3+}$
In $\ce{Co^2+}$ $\ce{NH3}$ acts as WFL.


SFL:Strong field ligand


WFL: weak..


 0