What's the difference between relative and percentage abundance?
What's the difference between realtive and average atomic mass? Is there a difference?
Relative abundance is the frequency with which "something" exists compared to a different "something". For example, you might find 1 blue bird for every 5 black birds (but you might want to compare it with with some other species).
Percent abundance is the frequency compared to the whole. For example, 1 bluebird sighted in a day versus 100 birds total seen that day is (1/100)*100 = 1%.
I don't believe there is any difference between relative and average atomic mass. All atomic masses that you find in a table are relative to 6-C-12 as well as averages of the masses of each isotope (relative to 6-C-12). You CAN calculate the atomic mass of a single isotope and that is relative to 6-C-12 but with no other isotopes added in.
Thank you!
The difference between relative and percentage abundance lies in how the data is expressed.
Relative abundance refers to the proportion of a specific isotope in a given sample relative to the total abundance of all isotopes of that element. It is usually expressed as a decimal or a fraction. For example, if an element has two isotopes, A and B, and the relative abundance of A is 0.6, it means that 60% of the element's isotopes are A.
Percentage abundance, on the other hand, is the relative abundance expressed as a percentage. It represents the proportion of a specific isotope relative to the total abundance of isotopes and is expressed as a whole number between 0% and 100%. Using the example above, the percentage abundance of isotope A would be 60%.
Now, moving on to the difference between relative and average atomic mass:
Relative atomic mass is the weighted average mass of all the naturally occurring isotopes of an element relative to the mass of carbon-12. It takes into account both the relative abundance and the mass of each isotope. The relative atomic mass of an element is usually given as a decimal number, rounded to the nearest whole number. For example, the relative atomic mass of hydrogen is approximately 1.008.
Average atomic mass, also known as atomic weight, is essentially the same concept as relative atomic mass. It represents the average mass of all the naturally occurring isotopes of an element, taking into account their respective abundance and mass. The average atomic mass is typically expressed with a specified unit (such as atomic mass units, amu) and can be a decimal number. For instance, the average atomic mass of carbon is approximately 12.01 amu.
So, to summarize, relative and percentage abundance highlight the proportion of a specific isotope in a given sample, while relative and average atomic mass represent the average mass of all isotopes in an element, accounting for their respective abundance.