what is the physical significance of ∆H=∆U for a chemical reaction?
This equation holds for a reaction involving ideal gases at constant T. Under these conditions, $\Delta (pV) = \Delta n RT$. Assume now that a gas phase reaction occurs at constant pressure and temperature. Then  $\Delta (pV) = p\Delta V$. This is equal to the negative of the pressure-volume work done by the system during the reaction since $w = -p\Delta V$. Therefore, at constant p and T, $\Delta H$ is equal to the change in energy of the system minus the pV work it did. This is consistent with the fact that we equate the enthalpy change for such a process at constant p and T with the heat exchanged, that is, $\Delta H = \Delta U -w = q_p$. Now if $\Delta n =0$ no expansion occurs at constant p and T, and so $w=0$ and $\Delta H = \Delta U = q$.