atomic mass unit

(noun)

The standard unit that is used for indicating mass of an atom.

Related Terms

Examples of atomic mass unit in the following topics:

  • Average Atomic Mass

    • The average atomic mass of an element is the sum of the masses of its isotopes, each multiplied by its natural abundance.
    • These different types of helium atoms have different masses (3 or 4 atomic mass units), and they are called isotopes.
    • This is because each proton and each neutron weigh one atomic mass unit (amu).
    • To calculate the average atomic mass, multiply the fraction by the mass number for each isotope, then add them together.
    • Whenever we do mass calculations involving elements or compounds (combinations of elements), we always use average atomic masses.

  • Avogadro's Number and the Mole

    • Avogadro's number is a proportion that relates molar mass on an atomic scale to physical mass on a human scale.
    • For example, since one atom of oxygen will combine with two atoms of hydrogen to create one molecule of water (H2O), one mole of oxygen (6.022×1023 of O atoms) will combine with two moles of hydrogen (2 × 6.022×1023 of H atoms) to make one mole of H2O.
    • For example, the mean molecular weight of water is 18.015 atomic mass units (amu), so one mole of water weight 18.015 grams.
    • This video introduces counting by mass, the mole, and how it relates to atomic mass units (AMU) and Avogadro's number.
    • Amedeo Avogadro is credited with the idea that the number of entities (usually atoms or molecules) in a substance is proportional to its physical mass.

  • Overview of Atomic Structure

    • An atom is the smallest unit of matter that retains all of the chemical properties of an element.
    • This can be determined using the atomic number and the mass number of the element (see the concept on atomic numbers and mass numbers).
    • Scientists define this amount of mass as one atomic mass unit (amu) or one Dalton.
    • Electrons are much smaller in mass than protons, weighing only 9.11 × 10-28 grams, or about 1/1800 of an atomic mass unit.
    • When considering atomic mass, it is customary to ignore the mass of any electrons and calculate the atom's mass based on the number of protons and neutrons alone.

  • Molar Mass of Compounds

    • Even the smallest quantity of a substance will contain billions of atoms, so chemists generally use the mole as the unit for the amount of substance.
    • The mass of one mole of atoms of a pure element in grams is equivalent to the atomic mass of that element in atomic mass units (amu) or in grams per mole (g/mol).
    • Although mass can be expressed as both amu and g/mol, g/mol is the most useful system of units for laboratory chemistry.
    • The characteristic molar mass of an element is simply the atomic mass in g/mol.
    • To calculate the molar mass of a compound with multiple atoms, sum all the atomic mass of the constituent atoms.

  • Atomic Number and Mass Number

    • Protons and neutrons both weigh about one atomic mass unit or amu.
    • Scientists determine the atomic mass by calculating the mean of the mass numbers for its naturally-occurring isotopes.
    • For example, the atomic mass of chlorine (Cl) is 35.45 amu because chlorine is composed of several isotopes, some (the majority) with an atomic mass of 35 amu (17 protons and 18 neutrons) and some with an atomic mass of 37 amu (17 protons and 20 neutrons).
    • Its average atomic mass is 12.11.
    • Determine the relationship between the mass number of an atom, its atomic number, its atomic mass, and its number of subatomic particles

  • John Dalton and Atomic Theory

    • Dalton introduced a theory that proposed that elements differed due to the mass of their atoms.
    • Dalton found an atomic theory of matter could elegantly explain this common pattern in chemistry - in the case of Proust's tin oxides, one tin atom will combine with either one or two oxygen atoms.
    • Dalton hypothesized this was due to the differences in the mass and complexity of the gases' respective particles.
    • Atomic theory has been revised over the years to incorporate the existence of atomic isotopes and the interconversion of mass and energy.
    • However, Dalton's importance in the development of modern atomic theory has been recognized by the designation of the atomic mass unit as a Dalton.

  • Characteristics of Mass Spectra

    • Modern mass spectrometers easily distinguish (resolve) ions differing by only a single atomic mass unit (amu), and thus provide completely accurate values for the molecular mass of a compound.
    • The unit mass resolution is readily apparent in these spectra (note the separation of ions having m/z=39, 40, 41 and 42 in the cyclopropane spectrum).
    • Since a molecule of carbon dioxide is composed of only three atoms, its mass spectrum is very simple.
    • The masses of molecular and fragment ions also reflect the electron count, depending on the number of nitrogen atoms in the species.
    • The weak even -electron ions at m/z=15 and 29 are due to methyl and ethyl cations (no nitrogen atoms).

  • Converting from One Unit to Another

    • Sometimes, it is necessary to deal with measurements that are very small (as in the size of an atom) or very large (as in numbers of atoms).
    • Then find a ratio that will help you convert the units of grams to atoms.
    • $3.41\:g \cdot \frac {1\:mole}{4.002\:g} \cdot \frac {6.022\cdot 10^{23} atoms}{1\:mole} = 5.13\:\cdot \:10^{23} atoms$
    • If you had a sample of a substance with a mass of 0.0034 grams, and you wanted to express that mass in mg, you could use the following dimensional analysis.
    • The given quantity is the mass of 0.0034 grams.

  • Converting between Mass and Number of Moles

    • Chemists generally use the mole as the unit for the number of atoms or molecules of a material.
    • The molar mass of any element can be determined by finding the atomic mass of the element on the periodic table.
    • For example, if the atomic mass of sulfer (S) is 32.066 amu, then its molar mass is 32.066 g/mol.
    • If the equation is arranged correctly, the mass units (g) cancel out and leave moles as the unit.
    • These relationships can be used to convert between units.

  • Mass-to-Mole Conversions

    • However, the measurements that researchers take every day provide answers not in moles but in more physically concrete units, such as grams or milliliters.
    • The relative atomic mass is a ratio between the average mass of an element and 1/12 of the mass of an atom of carbon-12.
    • From the relative atomic mass of each element, it is possible to determine each element's molar mass by multiplying the molar mass constant (1 g/mol) by the atomic weight of that particular element.
    • For a single element, the molar mass is equivalent to its atomic weight multiplied by the molar mass constant (1 g/mol).
    • For a compound, the molar mass is the sum of the atomic weights of each element in the compound multiplied by the molar mass constant.