The relationship between "relative to carbon-12 isotope" and "the natural abundance of that isotope" is as follows:
When calculating the relative atomic mass of an element, it is expressed relative to the mass of an atom of carbon-12 (denoted as 12C), which is assigned a mass of exactly 12 atomic mass units (u). Carbon-12 is chosen as the reference isotope because it is abundant and has a mass close to the average atomic mass of all the naturally occurring isotopes of carbon.
The relative atomic mass of an element takes into account the masses and abundances of its isotopes. Isotopes are atoms of the same element that have different numbers of neutrons. Different isotopes of an element exist in nature in varying proportions, meaning they have different abundances.
The relative atomic mass is calculated by considering the average mass of all the isotopes of an element, weighted by their natural abundances. The natural abundance of an isotope is the percentage or fraction of that isotope present in a naturally occurring sample of the element.
For example, carbon has two stable isotopes: carbon-12 (with a mass of 12 u) and carbon-13 (with a mass of 13 u). The natural abundance of carbon-12 is approximately 98.9%, and the natural abundance of carbon-13 is about 1.1%. By taking into account these natural abundances and their respective masses, the relative atomic mass of carbon is calculated to be approximately 12.01 u.
In summary, the "relative to carbon-12 isotope" in the context of relative atomic mass means that the mass of an element is expressed relative to the mass of a carbon-12 atom. The "natural abundance of that isotope" refers to the proportion or percentage of a specific isotope in the naturally occurring sample of the element, which is used to calculate the weighted average of the isotopic masses for the relative atomic mass.