Naturally occurring hydrogen on Earth has an atomic mass of \(1.0079\) amu. Suppose you were on another planet and found the atomic mass of hydrogen to be \(1.2000\) amu. How would you explain this? (The atomic mass of \({ }_{1}^{1} \mathrm{H}\) is \(1.0078252\) amu; the atomic mass of \({ }_{1}^{2} \mathrm{H}\) is \(2.1041022\) amu.)

Hydrogen, the simplest and most abundant element in the universe, possesses multiple forms called isotopes. These isotopes of hydrogen are atoms with the same number of protons but different numbers of neutrons.

The most common isotope, known as protium, has one proton and no neutrons, hence its atomic mass is approximately 1 amu. The second isotope, deuterium, contains one neutron in addition to the proton, giving it an atomic mass close to 2 amu. Lastly, there is tritium, with two neutrons and one proton, which has an atomic mass of about 3 amu but is radioactive and less commonly found.

These isotopic variations are integral to understanding variations in hydrogen's atomic mass across different environments, such as Earth versus another planet.

The average atomic mass of an element is the weighted average of the masses of its isotopes, considering their relative abundance in a given sample or environment. For hydrogen on Earth, this value is around 1.0079 amu, a figure that closely resembles the mass of the most abundant isotope, protium.

To calculate this value, one multiplies the mass of each isotope by its relative abundance (expressed as a decimal fraction) and then sums the results. If the relative abundance changes, so does the average atomic mass, which explains differences you might find on other planets or when comparing any two samples of the element.

The term relative abundance refers to the proportion of each isotope present in a mixture of isotopes. For example, on Earth, protium is far more abundant compared to deuterium. The variation in relative abundance can significantly impact the average atomic mass of an element.

Should you encounter a higher atomic mass of hydrogen, such as the 1.2000 amu on another planet, this would indicate a significantly greater proportion of the heavier deuterium isotope in that particular environment. This concept is essential for scientists in fields such as geochemistry and astrophysics, where isotopic compositions can reveal important clues about the origins and histories of substances.

The atomic mass unit (amu) is a standard unit of mass used to express atomic and molecular weights. Precisely defined, one amu is one twelfth the mass of a carbon-12 atom. Since the masses of electrons are negligible compared to those of neutrons and protons, and both have masses close to 1 amu, the atomic mass of an isotope is approximately the sum of its protons and neutrons.

Understanding the concept of amu is critical when discussing isotopes and their average atomic mass, as it provides a common scale for comparison. For hydrogen, the difference of small fractions of an amu can have a significant impact on its physical and chemical properties, just as the atomic mass difference has on the hypothetical planet.