Atomic mass

The atomic mass (m_a) is the mass of an atom at rest, most often expressed in chlorine). The atomic mass of an uncommon isotope can differ from the relative atomic mass or standard atomic weight by several mass units.

The relative atomic mass (A_r) (also known as atomic weight and average atomic mass) is the average of the atomic masses of all the chemical element's isotopes as found in a particular environment, weighted by isotopic abundance.^[2] This is frequently used as a synonym for the standard atomic weight and it is not incorrect to do so since the standard atomic weights are relative atomic masses, although it is less specific to do so. Relative atomic mass also refers to non-terrestrial environments and highly specific terrestrial environments that deviate from the average or have different certainties (number of significant figures) than the standard atomic weights.

The standard atomic weight refers to the mean relative atomic mass of an element in the local environment of the Earth's crust and Lithium represents a unique case where the natural abundances of the isotopes have been perturbed by human activities to the point of affecting the uncertainty in its standard atomic weight, even in samples obtained from natural sources such as rivers.

The relative isotopic mass is the relative mass of the isotope, scaled with atomic number.

Mass defects in atomic masses

The pattern in the amounts the atomic masses deviate from their mass numbers is as follows: the deviation starts positive at nuclear fusion reactions: fusion in elements lighter than iron produces energy, and fusion in elements heavier than iron requires energy.

Measurement of atomic masses

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

Conversion factor between atomic mass units and grams

The standard scientific unit for dealing with atoms in macroscopic quantities is the 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 an ⁵⁶Fe isotope is 55.935 u and one mole of ⁵⁶Fe will in theory weigh 55.935g, but such amounts of pure ⁵⁶Fe have never existed.

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

$m_{\rm{grams}}={m_{\rm{u}} \over N_{A}}$

where $u$ is the Avogadro's number.

Relationship between atomic and molecular masses

Similar definitions apply to 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

In the Avogadro's law. He formulated a law to determine atomic weights of elements: the different quantities of the same element contained in different molecules are all whole multiples of the atomic weight and determined atomic weights and molecular weights by comparing the vapor density of a collection of gases with molecules containing one or more of the chemical element in question ^[5].

In the early twentieth century, up until the 1960's 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 old chemists' scale.

The term atomic weight is being phased out slowly and being replaced by relative atomic mass, in most current usage. The history of this shift in nomenclature reaches back to the 1960's and has been the source of much debate in the scientific community. The debate was largely created 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" was in some ways redundant. In 1979, in a compromise move, the definition was refined and the term "relative atomic mass" was introduced as a secondary synonym. Twenty years later the primacy of these synonyms was reversed and the term "relative atomic mass" is now the preferred term; however the "standard atomic weights" have maintained the same name. ^[6]

Table of standard atomic weights

Group →	1	2	3	4	5	6	7	8	9	10	11	12	13	14	15	16	17	18
↓ Period
1	H 1.008																	He 4.003
2	Li 6.941	Be 9.012											B 10.81	C 12.01	N 14.01	O 16.00	F 19.00	Ne 20.18
3	Na 22.99	Mg 24.31											Al 26.98	Si 28.09	P 30.97	S 32.07	Cl 35.45	Ar 39.95
4	K 39.10	Ca 40.08	Sc 44.96	Ti 47.87	V 50.94	Cr 52.00	Mn 54.94	Fe 55.84	Co 58.93	Ni 58.69	Cu 63.55	Zn 65.39	Ga 69.72	Ge 72.61	As 74.92	Se 78.96	Br 79.90	Kr 83.80
5	Rb 85.47	Sr 87.62	Y 88.91	Zr 91.22	Nb 92.91	Mo 95.94	Tc [99]	Ru 101.07	Rh 102.91	Pd 106.42	Ag 107.87	Cd 112.41	In 114.82	Sn 118.71	Sb 121.76	Te 127.60	I 126.90	Xe 131.29
6	Cs 132.91	Ba 137.33	*	Hf 178.49	Ta 180.95	W 183.84	Re 186.21	Os 190.23	Ir 192.22	Pt 195.08	Au 196.97	Hg 200.59	Tl 204.38	Pb 207.2	Bi 208.98	Po [209]	At [210]	Rn [222]
7	Fr [223]	Ra [226]	**	Rf [263]	Db [262]	Sg [266]	Bh [264]	Hs [269]	Mt [268]	Ds [272]	Rg [272]	Uub [277]	Uut [284]	Uuq [289]	Uup [288]	Uuh [292]	Uus [291]‡	Uuo [293]‡

* Lanthanides				La 138.91	Ce 140.12	Pr 140.91	Nd 144.24	Pm [145]	Sm 150.36	Eu 151.96	Gd 157.25	Tb 158.93	Dy 162.50	Ho 164.93	Er 167.26	Tm 168.93	Yb 173.04	Lu 174.97
Actinides**				Ac [227]	Th 232.04	Pa 231.04	U 238.03	Np [237]	Pu [244]	Am [243]	Cm [247]	Bk [247]	Cf [251]	Es [252]	Fm [257]	Md [258]	No [259]	Lr [262]

periodic table

Metals						Metalloids	Nonmetals
Alkali metals	Alkaline earth metals	Inner transition elements		Transition elements	Other metals		Other nonmetals	Halogens	Noble gases
		Lanthanides	Actinides

**standard temperature and pressure (0 °C and 1 atm)**
Solids	Liquids	Gases

**Borders show natural occurrence**
Primordial	From decay	Synthetic	Undiscovered

References

^ IUPAC Definition of Atomic Mass
^ IUPAC Definition of Relative Atomic Mass
^ IUPAC Definition of Standard Atomic Weight
^ ATOMIC WEIGHTS OF THE ELEMENTS 2005 (IUPAC TECHNICAL REPORT), M. E. WIESER Pure Appl. Chem., V.78, pp. 2051, 2006
^ Origin of the Formulas of Dihydrogen and Other Simple Molecules Andrew Williams Vol. 84 No. 11 November 2007 • Journal of Chemical Education 1779
^ 'ATOMIC WEIGHT' -THE NAME, ITS HISTORY, DEFINITION, AND UNITS, P. DE BIEVRE and H. S. PEISER Pure&App. Chem., 64, 1535, 1992

be-x-old:Атамная маса

Categories: Stoichiometry

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