Hydrogen is not rejected as a standard element for relative atomic mass, but it does present some challenges when it comes to assigning it a precise atomic mass value. The relative atomic mass (also known as atomic weight) is a measure of the average mass of atoms of an element, taking into account the various isotopes and their abundances.
The main reason hydrogen poses challenges for atomic mass determination is its isotopic variability. Hydrogen has three isotopes: protium (¹H), deuterium (²H), and tritium (³H). Protium is by far the most abundant isotope and is commonly considered as "normal" hydrogen. However, deuterium and tritium exist in smaller amounts in nature and can significantly influence the average atomic mass.
The challenge arises because the atomic mass of protium is very close to 1 atomic mass unit (amu), while deuterium has a mass of approximately 2 amu and tritium has a mass of around 3 amu. To complicate matters further, the isotopic composition of hydrogen can vary across different sources or samples. Therefore, it is difficult to assign a single precise atomic mass value to hydrogen that accurately represents its isotopic diversity.
To address this issue, the International Union of Pure and Applied Chemistry (IUPAC) and the International Union of Pure and Applied Physics (IUPAP) have established a standard atomic mass scale based on the carbon-12 isotope. They assigned a value of exactly 12 amu to the carbon-12 isotope and derived the atomic masses of other elements relative to carbon-12 using precise measurements and isotopic abundances.
However, it is worth noting that hydrogen is still used as a reference point for expressing atomic masses on a scale known as the relative atomic mass scale. The atomic mass of hydrogen is often listed as approximately 1.008 amu, which reflects the average atomic mass of naturally occurring hydrogen isotopes weighted by their abundance.