Atomic mass

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Not to be confused with relative atomic mass or atomic weight.
Stylized lithium-7 atom: 3 protons, 4 neutrons, & 3 electrons (total electrons are ~1/4300th of the mass of the nucleus). It has a mass of 7.016 u. Rare lithium-6 (mass of 6.015 u) has only 3 neutrons, reducing the atomic weight (average) of lithium to 6.941.

The atomic mass (ma) is the mass of an atomic particle, sub-atomic particle, or molecule. The protons and neutrons account for almost all of the mass of an atom. It is commonly expressed in unified atomic mass units where by international agreement, 1 unified atomic mass unit is defined as 1/12 of the mass of a single carbon-12 atom (at rest).[1]

When divided by unified atomic mass units or daltons to form a pure number ratio, the atomic mass of an atom becomes a dimensionless number called the relative isotopic mass (see section below). Thus, the atomic mass of a carbon-12 atom is 12 u or 12 daltons (Da), but the relative isotopic mass of a carbon-12 atom is simply 12.

The atomic mass or relative isotopic mass refers to the mass of a single particle, and is fundamentally different from the quantities elemental atomic weight (also called "relative atomic mass") and standard atomic weight, both of which refer to averages (mathematical means) of naturally-occurring atomic mass values for samples of elements. Most elements have more than one stable nuclide; for those elements, such an average depends on the mix of nuclides present, which may vary to some limited extent depending on the source of the sample, as each nuclide has a different mass. By contrast, atomic mass figures refer to an individual particle species: as atoms of the same species are identical, atomic mass values are expected to have no intrinsic variance at all. Atomic mass figures are thus commonly reported to many more significant figures than atomic weights.

The atomic mass of atoms, ions, or atomic nuclei is slightly less than the sum of the masses of their constituent protons, neutrons, and electrons, due to binding energy mass loss (as per E=mc2).[2]

Relative isotopic mass: the same quantity as atomic mass, but with different units[edit]

Relative isotopic mass (a property of a single atom) is not to be confused with the averaged quantity "relative atomic mass," which is the same as atomic weight (see above), and is an average of values for many atoms in a given sample of a chemical element.

Relative isotopic mass is similar to atomic mass and has exactly the same numerical value as atomic mass, whenever atomic mass is expressed in unified atomic mass units. The only difference in that case, is that relative isotopic mass is a pure number with no units. This loss of units results from the use of a scaling ratio with respect to a carbon-12 standard, and the word "relative" in the term "relative isotopic mass" refers to this scaling relative to carbon-12.

The relative isotopic mass, then, is the mass of a given isotope (specifically, any single nuclide), when this value is scaled by the mass of carbon-12, when the latter is set equal to 12. Equivalently, the relative isotopic mass of an isotope or nuclide is the mass of the isotope relative to 1/12 of the mass of a carbon-12 atom.

For example, the relative isotopic mass of a carbon-12 atom is exactly 12. For comparison, the atomic mass of a carbon-12 atom is exactly 12 daltons or 12 unified atomic mass units. Alternately, the atomic mass of a carbon-12 atom may be expressed in any other mass units: for example, the atomic mass of a carbon-12 atom is about 1.998467052 x 10-26 kilogram.

As in the case of atomic mass, no nuclides other than carbon-12 have exactly whole-number values of relative isotopic mass. As is the case for the related atomic mass when expressed in unified atomic mass units or daltons, the relative isotopic mass numbers of nuclides other than carbon-12 are not whole numbers, but are always close to whole numbers. This is discussed more fully below.

Similar terms for different quantities[edit]

The atomic mass and relative isotopic mass are sometimes confused, or incorrectly used, as synonyms of relative atomic mass (also known as atomic weight) and the standard atomic weight (a particular variety of atomic weight). However, as noted in the introduction, atomic weight and standard atomic weight represent terms for (abundance-weighted) averages of atomic masses in elemental samples, not for single nuclides. As such, atomic weight and standard atomic weight often differ numerically from relative isotopic mass and atomic mass, and they can also have different units than atomic mass when this quantity is not expressed in unified atomic mass units (see the linked article for atomic weight).

The atomic mass (relative isotopic mass) is defined as the mass of a single atom, which can only be one isotope (nuclide) at a time, and is not an abundance-weighted average, as in the case of relative atomic mass/atomic weight. The atomic mass or relative isotopic mass of each isotope and nuclide of a chemical element is therefore a number that can in principle be measured to a very great precision, since every specimen of such a nuclide is expected to be exactly identical to every other specimen, as all atoms of a given type in the same energy state, and every specimen of a particular nuclide, are expected to be exactly identical in mass to every other specimen of that nuclide. For example, every atom of oxygen-16 is expected to have exactly the same atomic mass (relative isotopic mass) as every other atom of oxygen-16.

In the case of many elements that have one naturally occurring isotope (mononuclidic elements) or one dominant isotope, the actual numerical similarity/difference between the atomic mass of the most common isotope, and the relative atomic mass or (standard) atomic weight can be small or even nil, and does affect most bulk calculations. However, such an error can exist and even be important when considering individual atoms for elements that are not mononuclidic.

For non-mononuclidic elements that have more than one common isotope, the numerical difference in relative atomic mass (atomic weight) from even the most common relative isotopic mass, can be half a mass unit or more (e.g. see the case of chlorine where atomic weight and standard atomic weight are about 35.45). The atomic mass (relative isotopic mass) of an uncommon isotope can differ from the relative atomic mass, atomic weight, or standard atomic weight, by several mass units.

Atomic masses expressed in unified atomic mass units (i.e. relative isotopic masses) are always close to whole-number values, but never (except in the case of carbon-12) exactly a whole number, for two reasons:

  • protons and neutrons have different masses, and different nuclides have different ratios of protons and neutrons.

The ratio of atomic mass to mass number varies from about 0.99884 for 56Fe to 1.00782505 for 1H.

Any mass defect due to nuclear binding energy is experimentally a small fraction (less than 1%) of the mass of equal number of free nucleons. When compared to the average mass per nucleon in carbon-12, which is moderately strongly-bound compared with other atoms, the mass defect of binding for most atoms is an even smaller fraction of a dalton (unified atomic mass unit, based on carbon-12). Since free protons and neutrons differ from each other in mass by a small fraction of a dalton (about 0.0014 u), rounding the relative isotopic mass, or the atomic mass of any given nuclide given in daltons to the nearest whole number always gives the nucleon count, or mass number. The neutron count (neutron number) may then be derived by subtracting the number of protons (atomic number) from the mass number.

Mass defects in atomic masses[edit]

Binding energy per nucleon of common isotopes. A graph of the ratio of mass number to atomic mass would be similar.

The amount that the ratio of atomic masses to mass number deviates from 1 is as follows: the deviation starts positive at hydrogen-1, then decreases until it reaches a local minimum at helium-4, which is very strongly bound. Isotopes of lithium, beryllium, and boron are less strongly bound than helium, as shown by their increasing mass-to-mass number ratios. It is for this reason that lithium, beryllium, and boron cannot be formed in stars by fusion from hydrogen.

At carbon, the ratio of mass (in daltons) to mass number is defined as 1, and after carbon it becomes less than one until a minimum is reached at iron-56, iron-58 and nickel-62, then increases to positive values in the heavy isotopes, with increasing atomic number. This corresponds to the fact that nuclear fission in an element heavier than zirconium produces energy, and fission in any element lighter than niobium requires energy. On the other hand, nuclear fusion reactions: fusion of two atoms of an element lighter than scandium produces energy, whereas fusion in elements heavier than calcium requires energy. These general rules all have the notable exception of helium, which cannot produce energy in fusion of two atoms of He-4, or He-4 with lighter atoms, but requires fusion of three atoms of He-4 in the so-called triple alpha process to skip over lithium, beryllium, and boron (which do not produce energy from being made by fusion) to produce carbon (which does release energy when made by fusion).

Here are some values of the ratio of atomic mass to mass number:

Nuclide Ratio of atomic mass to mass number
1H 1.00782505
2H 1.0070508885
3H 1.0053497592
3He 1.0053431064
4He 1.0006508135
6Li 1.0025204658
12C 1
14N 1.0002195718
16O 0.9996821637
56Fe 0.9988381696
210Po 0.9999184462
232Th 1.0001640315
238U 1.0002133958

Measurement of atomic masses[edit]

Direct comparison and measurement of the masses of atoms is achieved with mass spectrometry.

Conversion factor between atomic mass units and grams[edit]

The standard scientific unit used to quantify the amount of a substance in macroscopic quantities is the mole (symbol: mol), which is defined arbitrarily as the amount of a substance which has as many atoms or molecules as there are atoms in 12 grams of the carbon isotope C-12. The number of atoms in a mole is called Avogadro's number, the value of which is approximately 6.022 × 1023.

One mole of a substance always contains almost exactly the relative atomic mass or molar mass of that substance; however, this may or may not be true for the atomic mass, depending on whether or not the element exists naturally in more than one isotope. For example, the relative atomic mass of iron is 55.847 g/mol, and therefore one mole of iron as commonly found on earth has a mass of 55.847 grams. The atomic mass of the 56Fe isotope is 55.935 u and one mole of 56Fe atoms would then in theory weigh 55.935 g, but such amounts of pure 56Fe have never been found (or separated out) on Earth. However there are 22 mononuclidic elements of which essentially only a single isotope is found in nature (common examples are fluorine, sodium, aluminum and phosphorus) and for these elements the relative atomic mass and atomic mass are the same. Samples of these elements therefore may serve as reference standards for certain atomic mass values.

The formula for conversion between atomic mass units and SI mass in grams for a single atom is:

1\ {\rm{u}}={M_{\rm{u}} \over N_{\rm A}}\ = {{1\ \rm{g/mol}} \over N_{\rm A}}

where M_{\rm u} is the Molar mass constant and N_{\rm A} is the Avogadro constant.

Relationship between atomic and molecular masses[edit]

Similar definitions apply to molecules. One can compute the molecular mass of a compound by adding the atomic masses of its constituent atoms (nuclides). One can compute the molar mass of a compound by adding the relative atomic masses of the elements given in the chemical formula. In both cases the multiplicity of the atoms (the number of times it occurs) must be taken into account, usually by multiplication of each unique mass by its multiplicity.

History[edit]

The first scientists to determine relative atomic masses were John Dalton and Thomas Thomson between 1803 and 1805 and Jöns Jakob Berzelius between 1808 and 1826. Relative atomic mass (Atomic weight) was originally defined relative to that of the lightest element, hydrogen, which was taken as 1.00, and in the 1820s Prout's hypothesis stated that atomic masses of all elements would prove to be exact multiples of that of hydrogen. Berzelius, however, soon proved that this was not even approximately true, and for some elements, such as chlorine, relative atomic mass, at about 35.5, falls almost exactly halfway between two integral multiples of that of hydrogen. Still later, this was shown to be largely due to a mix of isotopes, and that the atomic masses of pure isotopes, or nuclides, are multiples of the hydrogen mass, to within about 1%.

In the 1860s Stanislao Cannizzaro refined relative atomic masses by applying Avogadro's law (notably at the Karlsruhe Congress of 1860). He formulated a law to determine relative atomic masses of elements: the different quantities of the same element contained in different molecules are all whole multiples of the atomic weight and determined relative atomic masses and molecular masses by comparing the vapor density of a collection of gases with molecules containing one or more of the chemical element in question.[3]

In the 20th century, until the 1960s chemists and physicists used two different atomic-mass scales. The chemists used a "atomic mass unit" (amu) scale such that the natural mixture of oxygen isotopes had an atomic mass 16, while the physicists assigned the same number 16 to only the atomic mass of the most common oxygen isotope (O-16, containing eight protons and eight neutrons). However, because oxygen-17 and oxygen-18 are also present in natural oxygen this led to two different tables of atomic mass. The unified scale based on carbon-12, 12C, met the physicists' need to base the scale on a pure isotope, while being numerically close to the chemists' scale.

The term atomic weight is being phased out slowly and being replaced by relative atomic mass, in most current usage. This shift in nomenclature reaches back to the 1960s and has been the source of much debate in the scientific community, which was triggered by the adoption of the unified atomic mass unit and the realization that weight was in some ways an inappropriate term. The argument for keeping the term "atomic weight" was primarily that it was a well understood term to those in the field, that the term "atomic mass" was already in use (as it is currently defined) and that the term "relative atomic mass" might be easily confused with relative isotopic mass (the mass of a single atom of a given nuclide, expressed dimensionlessly relative to 1/12 of the mass of carbon-12; see section above).

In 1979, as a compromise, the term "relative atomic mass" was introduced as a secondary synonym for atomic weight. Twenty years later the primacy of these synonyms was reversed, and the term "relative atomic mass" is now the preferred term.

However, the term "standard atomic weights" (referring to the standardized expectation atomic weights of differing samples) have maintained the same name.[4] In the case of this latter term, simple replacement of the "atomic weight" term with "relative atomic mass" would have resulted in the term "standard relative atomic mass."

See also[edit]

References[edit]

  1. ^ IUPAC, Compendium of Chemical Terminology, 2nd ed. (the "Gold Book") (1997). Online corrected version:  (2006–) "atomic mass".
  2. ^ Atomic mass, Encyclopædia Britannica on-line
  3. ^ Williams, Andrew (2007). "Origin of the Formulas of Dihydrogen and Other Simple Molecules". J. Chem. Ed. 84 (11): 1779. Bibcode:2007JChEd..84.1779W. doi:10.1021/ed084p1779. 
  4. ^ De Bievre, P.; Peiser, H. S. (1992). "'Atomic weight': The name, its history, definition, and units". Pure&App. Chem. 64 (10): 1535. doi:10.1351/pac199264101535. 

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