General Chemistry: Atomic Mass vs. Atomic Weight
Atomic Mass vs. Atomic Weight
In chemistry, the terms atomic mass and atomic weight are often used to describe the heaviness of an element, but
they have distinct meanings that are important to understand.
ATOMIC MASS Atomic mass refers to the mass of a specific isotope of an element, measured in atomic mass
units (amu). The atomic mass unit is defined as one-twelfth the mass of a carbon-12 atom, which is approximately
1.66 × 10^-24 grams. Since a carbon-12 nucleus contains six protons and six neutrons, the atomic mass unit is
roughly equivalent to the mass of a proton or neutron. It is important to note that while protons and neutrons have
similar masses, there is a negligible difference between them, which is approximately equal to the mass of an
electron.
ATOMIC WEIGHT In contrast, atomic weight is a weighted average of the masses of all naturally occurring
isotopes of an element. This value takes into account both the mass and relative abundance of each isotope and is
reported as a constant for each element on the periodic table. Unlike atomic mass, which can vary depending on the
specific isotope being considered, atomic weight remains consistent for a given element.
In summary, while atomic mass pertains to individual isotopes and can vary among them, atomic weight provides an
average value that reflects all isotopes’ contributions based on their natural abundances. Understanding these
differences is crucial for accurately interpreting data in chemistry.
Atomic Mass and Mass Number
The atomic mass of an atom, measured in atomic mass units (amu), is nearly equivalent to its mass number, which
is the total count of protons and neutrons in the nucleus. However, it is important to note that some mass is lost
due to binding energy, a concept elaborated in Chapter 9 of the MCAT Physics and Math Review.
Isotopes
Atoms that belong to the same element but have different mass numbers are known as isotopes. The term
“isotope” derives from Greek, meaning “same place,” indicating that these atoms occupy the same position on the
periodic table. Isotopes vary in their neutron count while maintaining the same number of protons and electrons.
They are typically named by combining the element’s name with its mass number; for instance, carbon-12 (C-12) or
iodine-131 (I-131).
Hydrogen Isotopes
Hydrogen has three notable isotopes, each with distinct characteristics:
1. Protium: This isotope has one proton and an atomic mass of 1 amu.
2. Deuterium: This isotope contains one proton and one neutron, giving it an atomic mass of 2 amu.
3. Tritium: This isotope consists of one proton and two neutrons, resulting in an atomic mass of 3 amu.
Despite differences in their neutron counts, isotopes share identical chemical properties because they possess the
same number of protons and electrons.
Atomic Mass vs. Atomic Weight
In chemistry, the terms atomic mass and atomic weight are often used to describe the heaviness of an element, but
they have distinct meanings that are important to understand.
ATOMIC MASS Atomic mass refers to the mass of a specific isotope of an element, measured in atomic mass
units (amu). The atomic mass unit is defined as one-twelfth the mass of a carbon-12 atom, which is approximately
1.66 × 10^-24 grams. Since a carbon-12 nucleus contains six protons and six neutrons, the atomic mass unit is
roughly equivalent to the mass of a proton or neutron. It is important to note that while protons and neutrons have
similar masses, there is a negligible difference between them, which is approximately equal to the mass of an
electron.
ATOMIC WEIGHT In contrast, atomic weight is a weighted average of the masses of all naturally occurring
isotopes of an element. This value takes into account both the mass and relative abundance of each isotope and is
reported as a constant for each element on the periodic table. Unlike atomic mass, which can vary depending on the
specific isotope being considered, atomic weight remains consistent for a given element.
In summary, while atomic mass pertains to individual isotopes and can vary among them, atomic weight provides an
average value that reflects all isotopes’ contributions based on their natural abundances. Understanding these
differences is crucial for accurately interpreting data in chemistry.
Atomic Mass and Mass Number
The atomic mass of an atom, measured in atomic mass units (amu), is nearly equivalent to its mass number, which
is the total count of protons and neutrons in the nucleus. However, it is important to note that some mass is lost
due to binding energy, a concept elaborated in Chapter 9 of the MCAT Physics and Math Review.
Isotopes
Atoms that belong to the same element but have different mass numbers are known as isotopes. The term
“isotope” derives from Greek, meaning “same place,” indicating that these atoms occupy the same position on the
periodic table. Isotopes vary in their neutron count while maintaining the same number of protons and electrons.
They are typically named by combining the element’s name with its mass number; for instance, carbon-12 (C-12) or
iodine-131 (I-131).
Hydrogen Isotopes
Hydrogen has three notable isotopes, each with distinct characteristics:
1. Protium: This isotope has one proton and an atomic mass of 1 amu.
2. Deuterium: This isotope contains one proton and one neutron, giving it an atomic mass of 2 amu.
3. Tritium: This isotope consists of one proton and two neutrons, resulting in an atomic mass of 3 amu.
Despite differences in their neutron counts, isotopes share identical chemical properties because they possess the
same number of protons and electrons.
