Higher boranes should have less exhaust velocity due to fewer hydrogen atoms per boron (heating diborane gives off hydrogen and produces all the other boranes; like heating methane gives other hydrocarbons). But if they interest you, go ahead. Decaborane in particular is apparently well studied, so its properties shouldn't be that hard to come by.
On second thought... it is obvious why a smaller throat radius would give less dissociation: It leads to higher chamber pressure, and this, following Le Châtelier's Principle, lessens dissociation, as dissociation increases the number of gas molecules.
It's just that the propellants we typically use either dissociate so easily (hydrocarbons, HD) or so little (hydrogen) that we don't see this effect usually.
I think what we need to add is the bond dissociation energy, for the thermal decomposition to be treated properly. From this paper (need to use sci-hub.io to access it, just copy the DOI), the total bond energy (energy of formation from atoms) of diborane appears to be about 558 to 559 kcal/mol, which is equal to about 2337 kJ/mol - compare this to 2810 kJ/mol for ethane, a molecule of very similar size (six hydrogens and two larger atoms), so I think that number does make sense. So, the line that should be included in the diborane definition (under EnthalpyOfFormation_kJ__mol) is: BondDissociationEnergy_kJ__mol 2337.2
The same should also be done probably for B2O3 and BF3, but I need to sleep now.
Edit: Applying the method from the paper I linked, the bond dissociation energy for Diborane should be 2399.84 kJ/mol, for boron trioxide 2730 kJ/mol and for boron trifluoride 1935.09 kJ/mol.
I have successfully created Pentaborane. The reactions might have characteristics wrong though, I copied the stats for N2O4+Pentaborane from LOX Diborane, and the stats for Fluorine Pentaborane are, currently, copied from Fluorine Methane, although I am experimenting with copying the stats from Fluorine Diborane instead.