Beryllium chloride - Highly hydrolysable? Beryllium chloride's geometric structure is linear, with a bond angle of 180 degrees. Drawing the Lewis dot structure for this covalent compound, one can see that the compound is electron deficient and requires two lone pairs to achieve stability. As a result it is a strong Lewis acid. Since water has two lone pairs and is a strong Lewis base, mixing beryllium chloride with water results in stability for beryllium chloride and the evolution of energy.
My question is, why don't two lone pairs from either chlorine atom or even just a single chlorine atom make a coordinate bond with beryllium, as can be seen in the case of carbon monoxide where one of oxygen's lone pairs is donated to carbon to achieve nearest noble gas configuration of neon. If such a compound does not naturally exist, why does it not? Is it a theoretical defect or is there some explanation?
The structures of $\ce{BeX_2}$ where $\ce{X = F}$ or $\ce{Cl}$ can be found in A.F. Wells, Structural Inorganic Chemistry 5th edition, 1984 pages 412 and 1047.  There you will see that while the predicted structures for $\ce{BeX_2}$ are linear (sometimes rationalized with Lewis dot structures displaying double bonds and ignoring formal charge considerations of the halogens).  However, the actual structure of $\ce{BeCl_2}$ in the vapor-phase is a 3-coordinate dimer that converts to the linear structure at higher temperatures.  This behavior supports your 'gut feeling' that the electron deficiency of $\ce{Be}$ should be compensated for by another $\ce{Cl}$ atom.  The dimer network formed in the solid state allows lone pairs from $\ce{Cl}$ to be donated to the $\ce{Be}$ in an attempt to fill the valence shell of beryllium.

The use of a double bond to describe the nature of the $\ce{Be-Cl}$ bond is supported experimental data which shows that the bond length is shorter than expected for a single bond.  

The image below is taken from Inorganic Chemistry by Miessler, Fischer and Tarr, 5th edition.