text: In all cases, the bond, whether dative or "normal" electron-sharing, is a covalent bond.  In common usage, the prefix dipolar, dative or coordinate merely serves to indicate the origin of the electrons used in creating the bond.  For example, FB ← O(CH) ("boron trifluoride (diethyl) etherate") is prepared from BF and :O(CH), as opposed to the radical species [•BF] and [•O(CH)].  The dative bond is also a convenience in terms of notation, as formal charges are avoided: we can write D: + []A ⇌ D → A rather than D–A (here : and [] represent the lone-pair and empty orbital on the electron-pair donor D and acceptor A, respectively).  The notation is sometimes used even when the Lewis acid-base reaction involved is only notional (e.g., the sulfoxide RS → O is rarely if ever made by reacting the sulfide RS with atomic oxygen O).  Thus, most chemists do not make any claim with respect to the properties of the bond when choosing one notation over the other (formal charges vs. arrow bond).
  It is generally true, however, that bonds depicted this way are polar covalent, sometimes strongly so, and some authors claim that there are genuine differences in the properties of a dative bond and electron-sharing bond and suggest that showing a dative bond is more appropriate in particular situations.  As far back as 1989, Haaland characterized dative bonds as bonds that are (i) weak and long; (ii) with only a small degree of charge-transfer taking place during bond formation; and (iii) whose preferred mode of dissociation in the gas phase (or low ε inert solvent) is heterolytic rather than homolytic.  The ammonia-borane adduct (HN → BH) is given as a classic example: the bond is weak, with a dissociation energy of 31 kcal/mol (cf. 90 kcal/mol for ethane), and long, at 166 pm (cf. 153 pm for ethane), and the molecule possesses a dipole moment of 5.2 D that implies a transfer of only 0.2 e from nitrogen to boron.  The heterolytic dissociation of HN → BH is estimated to require 27 kcal/mol, confirming that heterolysis into ammonia and borane is more favorable than homolysis into radical cation and radical anion.  However, aside from clear-cut examples, there is considerable dispute as to when a particular compound qualifies and, thus, the overall prevalence of dative bonding (with respect to an author's preferred definition).  Computational chemists have suggested quantitative criteria to distinguish between the two "types" of bonding.
  Some non-obvious examples where dative bonding is claimed to be important include carbon suboxide (O≡C → C ← C≡O), tetraaminoallenes (described using dative bond language as "carbodicarbenes"; (RN)C → C ← C(NR)), the Ramirez carbodiphosphorane (PhP → C ← PPh), and bis(triphenylphosphine)iminium cation (PhP → N ← PPh), all of which exhibit considerably bent equilibrium geometries, though with a shallow barrier to bending.  Simple application of the normal rules for drawing Lewis structures by maximizing bonding (using electron-sharing bonds) and minimizing formal charges would predict heterocumulene structures, and therefore linear geometries, for each of these compounds.  Thus, these molecules are claimed to be better modeled as coordination complexes of :C: (carbon(0) or carbone) or :N: (mononitrogen cation) with CO, PPh, or N-heterocycliccarbenes as ligands, the lone-pairs on the central atom accounting for the bent geometry.  However, the usefulness of this view is disputed.