atomic mass

Examples of atomic mass 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.

• Molar Mass of Compounds

  • 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).
  • The characteristic molar mass of an element is simply the atomic mass in g/mol.
  • However, molar mass can also be calculated by multiplying the atomic mass in amu by the molar mass constant (1 g/mol).
  • To calculate the molar mass of a compound with multiple atoms, sum all the atomic mass of the constituent atoms.
  • For example, the molar mass of NaCl can be calculated for finding the atomic mass of sodium (22.99 g/mol) and the atomic mass of chlorine (35.45 g/mol) and combining them.

• 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

• Overview of Atomic Structure

  • Atoms are made up of particles called protons, neutrons, and electrons, which are responsible for the mass and charge of atoms.
  • 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.

• Mass-to-Mole Conversions

  • 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.
  • Multiplying by the molar mass constant ensures that the calculation is dimensionally correct because atomic weights are dimensionless.
  • 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.

• Converting between Mass and Number of Moles

  • 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.
  • According to the periodic table, the atomic mass of nickel (Ni) is 58.69 amu, which means that the molar mass of nickel is 58.69 g/mol.
  • To determine the number of atoms, convert the moles of Ni to atoms using Avogadro's number:
  • Convert between the mass and the number of moles, and the number of atoms, in a given sample of compound

• Early Ideas about Atoms

  • Dalton's experiments with gases led to some of the earliest measurements of atomic masses and a concept of atomic structure and reactivity.
  • Dalton's atomic theory contained the following ideas:
  • Atoms are indestructible.
  • Like the early philosophers, Dalton's theories were not popularly accepted for much of the 19th century, but his ideas have since been accepted, with amendments addressing subatomic particles and the interconversion of energy and mass.
  • Describe the early developments leading to the modern concept of the atom

• 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.

• High Resolution Spectra

  • In assigning mass values to atoms and molecules, we have assumed integral values for isotopic masses.
  • Because the strong nuclear forces that bind the components of an atomic nucleus together vary, the actual mass of a given isotope deviates from its nominal integer by a small but characteristic amount (remember E = mc2).
  • By designing mass spectrometers that can determine m/z values accurately to four decimal places, it is possible to distinguish different formulas having the same nominal mass.
  • Mass spectrometry therefore not only provides a specific molecular mass value, but it may also establish the molecular formula of an unknown compound.
  • Tables of precise mass values for any molecule or ion are available in libraries; however, the mass calculator provided below serves the same purpose.

• The Law of Conservation of Mass

  • The law of conservation of mass states that mass in an isolated system is neither created nor destroyed.
  • However, Antoine Lavoisier described the law of conservation of mass (or the principle of mass/matter conservation) as a fundamental principle of physics in 1789.
  • In other words, in a chemical reaction, the mass of the products will always be equal to the mass of the reactants.
  • We can therefore visualize chemical reactions as the rearrangement of atoms and bonds, while the number of atoms involved in a reaction remains unchanged.
  • This video explains how atoms are conserved in a chemical reaction.