The atomic mass unit (amu) is defined as "one-twelfth of the mass of one atom of carbon-12" because carbon-12 is a commonly occurring isotope of carbon and has a relative atomic mass of exactly 12 amu.
The choice of carbon-12 as the reference isotope is primarily due to historical reasons and practical considerations. When the concept of atomic mass was first developed, scientists needed a reference point against which they could compare the masses of other atoms. Carbon-12 was chosen as a reference because it is abundant, stable, and has a whole number atomic mass, making calculations easier.
The decision to use carbon-12 as the reference also allows for consistent comparisons of atomic masses across different elements. Since atomic masses are relative to the atomic mass of carbon-12, it provides a standard scale for expressing the masses of other atoms. Using the mass of one proton as a reference would not be as convenient because it does not provide a whole number atomic mass, and it would complicate the comparison of atomic masses between different elements.
Furthermore, the atomic mass unit is not solely based on the mass of one proton because it also takes into account the masses of neutrons and electrons. Neutrons and electrons contribute to the overall mass of an atom, and using carbon-12 as the reference provides a more comprehensive measure that considers the mass of the entire atom.