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mass number

mass number, often represented by the symbol A, the total number of nucleons (neutrons and protons) in the nucleus of an atom. All atoms of a chemical element have the same atomic number (number of protons in the nucleus) but may have different mass numbers (from having different numbers of neutrons in the nucleus). Atoms of an element with the same mass number make up an isotope of the element. Different isotopes of the same element cannot have the same mass number, but isotopes of different elements often do have the same mass number, e.g., carbon-14 (6 protons and 8 neutrons) and nitrogen-14 (7 protons and 7 neutrons).

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mass number

mass num·ber • n. Physics the total number of protons and neutrons in a nucleus.

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mass number

mass number (nucleon number) The total number of protons and neutrons found in the nucleus of an atom.

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mass number

mass number The sum of the protons and neutrons in the nucleus of an atom.

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"mass number." A Dictionary of Earth Sciences. . Encyclopedia.com. 17 Nov. 2018 <https://www.encyclopedia.com>.

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Mass Number

Mass Number

The mass number (A) of an atom, within physics and chemistry, is the total number of nucleons (protons and neutrons) in its nucleus. It is also called atomic mass number or nucleon number.

Different isotopes of the same element have different mass numbers because their nuclei contain different numbers of neutrons. In the written symbol for a particular isotope, the mass number is written at the upper left of the symbol for the element, as in23892 U, where 92 is the atomic number (Z) of uranium (U) and 238 is the mass number (A) of this particular isotope. The symbol is read uranium-238. (The difference between an atoms mass number and its atomic number provides the number of neutrons (n) that the atom possesses: n = A - Z. In the above example, isotope uranium-238 has 146 neutrons: 238 - 92.)

The mass number is always a whole number; because it is a count of the particles. It differs from the exact mass of the atom in atomic mass units, amu, which is often known and expressed to six decimal places. (One amu is exactly one-twelfth of the mass of an atom of carbon-12,12C, and is equal to approximately 1.66× 10-24 g.) There are two reasons why the mass number of an atom is different from its exact mass. First, neutrons and protons do not happen to weigh exactly one amu apiece; the proton actually weighs 1.0072765 amu and the neutron weighs 1.0086650 amu. Second, when neutrons and protons are bound together as an atomic nucleus, the nucleus has less mass than the sum of the masses of the neutrons and protons. The difference in mass, when expressed in energy (E) units according to GermanAmerican physicist Albert Einsteins (18791955) formula E = mc2 (where m refers to mass and c to speed of light), is called the binding energy of the nucleus.

To understand this situation, one can think of the binding energy as the strength of the glue that holds the protons and neutrons together as a nucleus. It is, therefore, the amount of energy required to break the glue and pull the nucleus apart into its individual neutrons and protons. However, if energy must be added to an object in order to pull it apart, and if energy and mass are equivalent, then one could say that mass had to be added to pull it apart. The separated particles will, therefore, have more mass than when they were bound together as a nucleus.

See also Periodic table.

Robert L. Wolke

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Mass Number

Mass number

The mass number of an atom is the total number of protons plus neutrons in its nucleus.

Different isotopes of the same element have different mass numbers because their nuclei contain different numbers of neutrons. In the written symbol for a particular isotope , the mass number is written at the upper left of the symbol for the element, as in 23892U, where 92 is the atomic number of uranium (U) and 238 is the mass number of this particular isotope. The symbol is read "uranium-238." The mass number is always a whole number; it is a count of the particles. It differs from the exact mass of the atom in atomic mass units, amu, which is often known and expressed to six decimal places. (One amu is exactly one-twelfth of the mass of an atom of carbon-12, 12C, and is equal to approximately 1.66 × 10-24 g.) There are two reasons why the mass number of an atom is different from its exact mass. First, neutrons and protons don't happen to weigh exactly one amu apiece; the proton actually weighs 1.0072765 amu and the neutron weighs 1.0086650 amu. Second, when neutrons and protons are bound together as an atomic nucleus, the nucleus has less mass than the sum of the masses of the neutrons and protons. The difference in mass, when expressed in energy units according to Einstein's formula E=mc2, is called the binding energy of the nucleus.

To understand this situation, think of the binding energy as the strength of the "glue" that holds the protons and neutrons together as a nucleus. It is, therefore, the amount of energy that would be required to break the "glue" and pull the nucleus apart into its individual neutrons and protons. But if energy must be added to an object in order to pull it apart, and if energy and mass are equivalent, then we could say that mass had to be added to pull it apart. The separated particles will therefore have more mass than when they were bound together as a nucleus.

See also Periodic table.

Robert L. Wolke

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