Why are fluorides of transition metals unstable in low oxidation states? [duplicate]








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[Stability of 3d metal fluorides and iodides in different oxidation states](/questions/102326/stability-of-3d-metal-fluorides-and-iodides-in-different-oxidation-states)

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Why are fluorides of transition metals unstable in low oxidation states? 


I would think that since fluorine and oxygen are highly electronegative, it's obvious that they be stable at high oxidation states. However, $\ce{Cl}$ is also highly electronegative, but it does exist in low oxidation states with transition metals.


For example $\ce{CuCl}$ exists but not $\ce{CuF}$ and $\ce{TiCl2}$ exists but not $\ce{TiF\_{2}}$. Why is it so? 



> 
> *Fluorides of transition metals unstable in low oxidation states* 
> 
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> 


This is the statement given in my high school chemistry text book but no explanation provided. I am interested to know the reason for it. 




You can take the example of $\ce{CuF}$ which is not known.


Reason: $\ce F$ is such an electronegative element that it will always oxidise $ \ce{Cu}$ to $\ce{Cu^2+}$ and not $\ce{Cu^1+}$. Hence, whenever $\ce{Cu}$ reacts with flourine, copper(II) flouride is formed. The following reaction illustrates this:


$$\ce{Cu + F2 -> CuF2}$$


$\ce{CuF2}$ loses fluorine at temperatures above $\pu{950 °C}$. $\ce{CuF}$ will be formed initially and as your question says, it is highly unstable so it will be transformed to some other stuff and the reactions are:


$$\ce{2CuF2 -> 2CuF (unstable) + F2}$$


$$\ce{2CuF -> CuF2 + Cu}$$


Similarly, you can illustrate this for $\ce{VF2}$.


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