The Chemical Nature of Atoms

James Richard Fromm


Although we now know that John Dalton's original idea that atoms were small, hard particles of various types is an oversimplified picture, it is still convenient to use this concept in dealing with many aspects of chemistry. Dalton was quite aware of the fact that different kinds of atoms had different masses. However, information about the inner structure of atoms has been obtained only in our own twentieth century. It still lies more in the domain of physics than in that of chemistry, although modern chemists make considerable use of it.

In 1911, Ernest Rutherford showed, by bombardment of atoms with the nuclei of helium atoms (alpha particles), that the mass of an atom is concentrated in a very small central portion of the atom which is called the atomic nucleus. The atomic nucleus is made up of nucleons, of which there are two important fundamental types: electrically positive protons and electrically neutral neutrons. Surrounding the atomic nucleus are the electrically negative electrons. The masses and charges of these three fundamental constituents of atoms are given in the Table below. (We now know of other subatomic particles and nuclear reactions, but the more detailed discussions are found elsewhere.)


Table: Characteristics of the Fundamental Particles
Particle Electrical Charge Rest Mass Molar Mass
C kg g/mol
electron -1.60217733(49) x 10-19 0.91093897(54) x 10-30 0.0005486
proton +1.60217733(49) x 10-19 1672.6231(10)  x 10-30 1.0072697
neutron 0.0 1674.9543(86)  x 10-30 1.0086650

The electrons are the portion of the atom which engages in chemical reactions. However, the properties of the electrons of an atom are determined in large part by the number of protons present in the nucleus of the atom. The number of neutrons generally has a negligible effect upon the properties of the electrons which are of chemical significance. Thus the chemical nature of an atom, which is to say the chemical properties of an element, is determined by the number of protons in the nucleus. This number of protons is called the atomic number. The mass of the atom, its atomic mass, depends significantly upon both the number of protons and upon the number of neutrons present in the nucleus.

Since the properties of the electrons depend upon the number of protons in the nucleus of an atom, each different atomic number corresponds to a qualitatively different kind of atom. Each different kind of atom makes up a different chemical element. The atomic numbers of the known elements are all integers, and range from one (hydrogen) to well above 100. Above atomic number 100, atomic nuclei are found to become increasingly unstable; they rapidly break apart or fission into nuclei of elements with lower atomic numbers.

Isotopes and Atomic Masses

For many of the chemical elements there are several known isotopes. Isotopes are atoms with different atomic masses which have the same atomic number. The atoms of different isotopes are atoms of the same chemical element; they differ in the number of neutrons in the nucleus.

Chemists sometimes find it necessary to specify the atomic mass of an isotope. This is done by writing the atomic mass as a superscript preceding the atomic letter symbol.


Example. The fissionable isotope of uranium is 235U; the nonfissionable isotope 238U makes up most of naturally occurring uranium. Since uranium has the atomic number 92, a nucleus of 235U contains 92 protons and 143 neutrons while a nucleus of 238U contains 92 protons and 146 neutrons.


Although the atomic number of an atom is implicit in the element symbol, nuclear physicists and some chemists choose to write the atomic number explicitly in nuclear reactions. The atomic number can be written as a subscript preceding the element symbol, as 92U, but most chemists prefer to omit it since the atomic symbol includes the same information.

The relative abundances of the different isotopes of an element are constant to a very great degree in natural materials. Man-made isotopes, natural radioactive decay, and isotopic separation processes can, however, produce samples of an element with relative abundances different from those which are normally found. Relative natural isotopic abundances and atomic masses for a few selected elements are given in the Table below.

Atoms of the same chemical element all have essentially the same chemical properties and reactivity but they do not always have the same mass because, although the number of protons in the nucleus is the same for all atoms of the same element, the number of neutrons is not. The number of electrons also may vary, but only if the atom ionizes, and in any case the relative mass of the electron is very much less than that of a proton or neutron. The loss or gain of electrons is often ignored. However, the mass of the neutron is large enough that for any element a difference of one neutron is significant. As a consequence the molar masses of the different isotopes of an element are significantly different. Most elements as they occur naturally on earth are mixtures of several isotopes.


Example. Fluorine is one of the few elements that has only one naturally occurring isotope, 19F. Chlorine, however, exists as a mixture of isotopes of approximate atomic mass 35 g/mol (relative abundance 75.77%) and 37 g/mol (relative abundance 24.23%). The mixture is uniform across all naturally occurring samples of chlorine in terrestrial (and, so far, extraterrestrial) samples. Chlorine is typical of most elements in that the distribution of isotopes in naturally occurring samples, the natural abundance, is uniform to within a few tenths or hundredths of a per cent.


The atomic mass of an element used by chemists is always the atomic mass of the naturally occurring mixture of isotopes of that element. For many elements, the atomic masses of the individual isotopes are now known more accurately than is the atomic mass of the element.

Moles of Atoms

The atomic mass of an element is a relative quantity. Originally the atomic mass of hydrogen, the lightest of the elements, was taken to be one and the atomic masses of all other elements were measured in relation to the atomic mass of hydrogen. This later proved to have been a poor choice. Not only does hydrogen naturally consist of more than one isotope, but there was the additional question (particularly among early chemists) as to whether monatomic hydrogen or diatomic hydrogen should be taken as having atomic mass one.

After some effort, and one major false start with oxygen, chemists and physicists agreed on a common scale of relative atomic mass. Carbon of isotopic mass twelve was assigned an atomic mass of exactly twelve, and all other atomic masses whether of isotopes or of elements were specified relative to carbon of atomic mass twelve. This had the effect of making the relative atomic mass of hydrogen 1.0079...rather than exactly 1.0000.... The difference of less than 1% is too small to matter in many approximate chemical calculations, but it is large enough to be significant when accurate work must be done.

Physicists, and some chemists, measure the masses of individual atoms in kg, g, or atomic mass units. For most chemists, however, the mass of a single atom is inconveniently small and the molar mass of a substance is used. The molar mass of an atom is the mass of a very large number of identical atoms-- one mole of atoms. One mole of atoms is by definition that number of atoms which exist in exactly twelve grams of carbon of isotopic mass twelve (12C). This number is called the Avogadro number, NA, and the best current determination of its value is 6.0221367(36) x 10+23. Moles of atoms and molecules are so central to chemistry that they will be used continually throughout all courses in chemistry.

Molar Atomic Masses of Elements

The molar mass of an atom is simply the mass of one mole of identical atoms. However, most of the chemical elements are found on earth not as one isotope but as a mixture of isotopes, so the atoms of one element do not all have the same mass. Chemists therefore distinguish the molar atomic mass of an isotope, which is the mass of one mole of the identical atoms which form that isotope, from the molar atomic mass of an element, which is the mass of one mole of the atoms of the various isotopes of that element having the natural abundances as they are found on earth. For many elements the variation found in the natural abundances limits the accuracy with which the molar atomic mass of that element can be known. Those elements for which this is true are indicated in the table of atomic masses.

Chemists deal with elements as they are naturally found, and so the atomic mass of a particular isotope is of less interest than the weighted mean molar atomic mass of the individual isotopes which is the molar atomic mass of the naturally occurring element. This property has often been called the atomic weight or chemical atomic weight of the element. The weighted mean molar atomic mass of an element as it naturally occurs will be referred to simply as the atomic mass of the element from now on. The natural abundances of the isotopes of a few selected elements are given in the Table below.


Table: Natural Abundances of the Isotopes of Selected Elements
Element Isotope Abundance Molar Mass Half-Life
% Natural g/mol Years
Hydrogen 1H 99.985(1) 1.007825035(12) Stable
Deuterium 2H 0.015(1) 2.014101779(4) Stable
Tritium 3H Trace 3.01604927(4) 12.6
helium 3He Trace 3.01602931(4) Stable
4He 100.000 4.00260324(5) Stable
Carbon 12C 98.90 (3) 12.0 (exactly) Stable
13C 1.10(3) 13.003354826(17) Stable
14C Trace 14.003241982(27) 5760
Oxygen 16O 99.762(15) 15.99491463(5) Stable
17O 0.038(3) 16.9991312(4) Stable
18O 0.200(12) 17.9991603(9) Stable
Fluorine 19F 100.000 18/99840322(15) Stable
Chlorine 35Cl 75.77 (5) 34.968852721(69) Stable
36Cl Trace --- 0.31  x 10+6
37Cl 24.23 (5) 36.96590262(11) Stable
Potassium 39K 93.2581(30) 38.9637074(12) Stable
40K 0.0117(2) 39.963992(12) 1.29 x 10+9
41K 6.7302(30) 40.9618254(12) Stable
Radium 226Ra 100.000 226.025403(3) 1620
Thorium 232Th 100.00 232.0380508(23) 1.41  x 10+9
Uranium 233U Trace 233.0395 0.162 x 10+6
234U 0.0055 (5) 234.0409468(24) 0.247 x 10+6
235U 0.7200(12) 235.0439242(24) 0.71  x 10+9
238U 99.2745(15) 238.0507847(23) 4.51  x 10+9

Example. The atomic mass of chlorine can be computed from the abundance data and molar masses given above as follows: (0.2423 x 36.9659) + (0.7577 x 34.9689) = 35.4528 g/mol


Example. The relative abundances of the isotopes 6Li and 7Li in naturally occurring lithium can be computed as follows. Their atomic masses are 6.0151214(7) and 7.0160030(9) respectively. The atomic mass of naturally occurring lithium given in the table of atomic masses of the elements is 6.941(2) g/mol. If the relative abundance of 6Li is designated as a, then

(a)(6.0151214) + (1.0 - a)(7.0160030) = 6.941

7.0160 - 1.0008816a = 6.941

which gives a value of the abundance as 0.0749 or 7.49%. The remaining 92.51% is the 7Li isotope.


The isotopic abundances of lithium cannot be given with greater accuracy than this due to natural variations, and are usually given only as 7.5% and 92.5%. Since 6Li is used in nuclear weapons, lithium which has been subjected to processing for this purpose may have an isotopic composition which is greater in 7Li than that of the naturally occurring lithium. The relative abundances of lithium isotopes in commercial lithium salts may vary significantly due to the extraction of 6Li by the governments of the United States and other nations for use in hydrogen fusion weapons.


Copyright 1997 James R. Fromm