# Why Cl⁻ can't act as bidentate ligand?

$\ce{Cl-}$ is a monodentate ligand. It has four lone pairs then why it can't donate two lone pairs to central atom and act as a bidentate ligand?

A bidentate ligand is one that forms bonds to the metal center through two different atoms, not through two lone pairs. While chloride does have more than one pair of valence electrons, it only has one atom. Thus it cannot be bidentate.

If the chloride shares a second pair of electrons with the metal center, there would still only be one bond site between the ligand and the metal:

$$\ce{[M+]-Cl <-> [M]=Cl+}$$

A bidentate ligand needs two lone pairs on two separate atoms. Consider ethylenediamine (en), one of the classic bidentate ligands: $\ce{H2NCH2CH2NH2}$. Each nitrogen atom has a lone pair that can be used to form a bond to the metal center, creating two bonds from two different sites on the ligand. See the structure of the tris(ethylenediamine)cobalt(III) ion from tris(ethylenediamine)cobalt(III) chloride $(\ce{Co(en)3Cl3})$.

Update: As Nicolau points out in a comment, chloride can bridge two metal centers. This is not the denticity phenomenon, which is one ligand with two bonding sites to the same metal. This is one ligand with bonds to two metal centers. See the structure of niobium pentachloride (really $\ce{Nb2Cl10}$ or better $\ce{(NbCl5)2}$.

• Also of some relevance is that chlorine can make two bonds with different metal centres, forming a bridged compound. Again, this is not an example of bidenticity. – Nicolau Saker Neto Dec 15 '13 at 15:22
• Thanks! Does that means in some compounds chlorine can form multiple bonds with a metal? – Ayush Pateria Dec 15 '13 at 19:38
• @AyushPateria That may be relatively tricky to find (I know of no such compound in ordinary conditions, at least). Chlorine is one of the most electronegative elements, whilst most metals are on the other end of the scale, so there probably isn't a strong tendency for a chlorine atom to donate electrons to a metal to the point where it acquires a significant partial positive charge. Also, atoms outside the second period tend to form somewhat weaker $\pi$ bonds, and will generally find a more stable structure without multiple bonding, often by dimerization. – Nicolau Saker Neto Dec 15 '13 at 21:17
• It is probably best to think that halogens can, in certain situations, display some multiple bond character. A very simple and interesting example is boron trifluoride, $\ce{BF3}$. Boron is quite a strong Lewis acid, and is able to partially stabilize resonance structures where the fluorine atoms donate an extra bit of electron density to the boron, thus creating bonds which are shorter than expected for a boron-halogen single bond. This may also happen in some structures with metals and halogens, as implied by Ben's first equation displaying canonical structures with double bonding. – Nicolau Saker Neto Dec 15 '13 at 21:28
• Shouldn't chlorine with one bond have a negative formal charge? – Dissenter Mar 9 '15 at 1:16

Chlorine ion is monodentate.For bidentate ligand it should have atleast two bonding sites.In chlorine it has only one donatable lone pair.So it can't act as bidentate ligand.