Carbon is the foundational element of all known life on Earth, forming complex molecules through its unique ability to bond with other atoms. To understand carbon’s characteristics, we must look inside the atom, which is composed of a nucleus containing protons and neutrons, surrounded by electrons. The count and arrangement of these subatomic particles dictate an element’s identity and its physical properties.
The Difference Between Mass Number and Atomic Weight
The terms mass number and atomic weight are often used interchangeably, but they represent distinct concepts in chemistry. The mass number is a straightforward count of the total number of protons and neutrons found within the nucleus of an individual atom. Since it is a count of whole particles, the mass number is always a whole, positive integer.
Atomic weight, conversely, reflects the mass of an element as it exists naturally in the world, which is a blend of all its various forms, known as isotopes. This value is a calculated, weighted average of the mass numbers of all the naturally occurring isotopes. Because this calculation factors in the relative abundance of each isotope, the atomic weight is typically a decimal number. For example, the element Bromine has isotopes with mass numbers of 79 and 81, but its atomic weight averages to approximately 79.904.
The mass number is specific to a single atom or isotopic form. The atomic weight, however, is the standard value listed on the periodic table, offering a general mass for the element based on its terrestrial presence.
The Composition of Carbon-12
When the mass number of carbon is requested, the answer almost always refers to the element’s most common and stable form, Carbon-12. All atoms of Carbon are defined by having six protons in their nucleus; this number is called the atomic number and is what makes the element carbon.
The standard Carbon-12 isotope contains exactly six neutrons in addition to its six protons. The mass number is the sum of these particles, six protons plus six neutrons, which equals 12. This specific isotope is so prevalent that it accounts for approximately 98.93% of all naturally occurring carbon on Earth.
The stability and abundance of Carbon-12 led scientists to establish it as the reference standard for measuring the mass of all other atoms. The unified atomic mass unit (u) is defined as precisely one-twelfth the mass of a single, unbound Carbon-12 atom.
How Isotopes Change Carbon’s Mass Number
While Carbon-12 is the dominant form, carbon atoms can also exist with different numbers of neutrons, creating different isotopes with varying mass numbers. Isotopes are atoms of the same element that share the same atomic number (six protons) but differ in their neutron count.
Carbon-13 is the second stable isotope of carbon, making up about 1.1% of natural carbon. This isotope has six protons and seven neutrons, which gives it a mass number of 13. Its stability and nuclear properties make it valuable in scientific research, such as in nuclear magnetic resonance (NMR) spectroscopy and metabolic tracing studies.
The third naturally occurring isotope is Carbon-14, which contains six protons and eight neutrons, resulting in a mass number of 14. Unlike the other two forms, Carbon-14 is unstable and undergoes radioactive decay over time. This characteristic makes Carbon-14 useful for radiocarbon dating, where scientists measure the remaining amount of the isotope in organic materials to determine their age.