To determine the strongest base among the options given (AsH₃, NH₃, PH₃, SbH₃), we should look into the basicity of these hydrides. Basicity refers to the ability of a molecule to donate a lone pair of electrons, which in the case of these hydrides involves the donation of the lone pair from the central atom to a proton. This ability largely depends on the electron density around the central atom which has the lone pair.

In Group 15 of the periodic table (N, P, As, Sb), as we move down the group from nitrogen (N) to antimony (Sb), the size of the central atom increases and the electronegativity decreases. Consequently, the electron density available to be donated decreases because the larger atomic size causes the lone pair to be more diffused and less concentrated.

NH₃ (Ammonia):

Ammonia is known to be a stronger base than the other Group 15 hydrides due to nitrogen's higher electronegativity relative to phosphorus, arsenic, and antimony. This higher electronegativity means the lone pair on nitrogen is held more tightly, making it more available for bonding compared to those lower in the group.

PH₃ (Phosphine), AsH₃ (Arsine), SbH₃ (Stibine):

These molecules, being lower in the group, have central atoms with larger sizes and lower electronegativities than nitrogen. This results in a weaker ability to donate their lone pair of electrons, thus demonstrating weaker basic properties compared to ammonia.

From the given choices, Option B: NH₃ is the strongest base. This is because the ammonia molecule's smaller, more electronegative nitrogen atom can more effectively share its electron density compared to the larger, less electronegative atoms found in the other options.