The Chemical Nature of Atoms

Although we now know that Dalton's original idea that atoms were small, hard particles of various types is to simple a 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 20th 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 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. (There are other subatomic particles and nuclear reactions, but these will be postponed till the senior course in chemistry). 

Characteristics of the Fundamental Particles

Electrical Charge
Rest Mass
Molar Mass
-1.60217733(49) x 10-19
0.91093897(54) x 10-30
+1.60217733(49) x 10-19
1672.6231(10)  x 10-30
1674.9543(86)  x 10-30 

The electrons are the portion of the atom which engage 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 neglegible effect upon the properties of the electrons which are of chemical significance. The neutrons play are far more important role in the nuclear reactions. In normal chemical reactions they are only important for their mass. The chemical properties of an element, are 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 upon the sum of the number of protons and 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.

Symbols of the Elements: A Chemist's Shorthand

The elements each have their own unique symbol. The names of all of the other elements were assigned by their discoverers or are traditional names which have become universally accepted. Five of the elements have been known from antiquity and have symbols derived from their Latin names: Fe, iron, ferrum; Cu, copper, cuprum; Ag, silver, argentum; Au, gold, aurum: and Hg, mercury, hydrargyrum. Most of the other elements have symbols which are related to their English names, but three do not: Na, sodium (natrium); K, potassium (kalium); and W, tungsten (wolfram) have symbols related to their names in other languages. Different authors write these symbols in different ways.

For example:  1123Na   or 2311Na    The smaller number is always the Atomic number. The larger number is called the Mass number. The second version will be used in these pages, with the mass number on top and the atomic number on the bottom. The mass number is also never found on a periodic table as a nice whole number. It usually has some decimal places after it for reasons which we will cover shortly.

The atomic number tells the number of protons that are found in the nucleus. It also tells you the number of electrons that the element has in its outside shells. If the atomic number of nickel is 28 then every atom of nickel has 28 protons in its nucleus and 28 electrons outside the nucleus.

The mass number is a sum. It is the number of protons + the number of neutrons together in the nucleus.  Why not the electrons? The electrons are so tiny that it would take 1836 of them to equal the mass of a proton. Since the largest element known has only 103 electrons the mass given to the atom by its electrons is to tiny the be considered and can be safely ignored.

You can calculate the number of neutrons in a nucleus very easily.

Number of neutrons = mass number - atomic number

In nickel,   5928Ni, the 28 is the atomic number and 59 is the mass number.

Therefore the number of protons = 28, the number of electrons is also 28, and the number of neutrons is 59-28=31.

Isotopes and Atomic Masses

For many of the chemical elements there are several known isotopes.  Isotopes are atoms with different atomic mass which have the same atomic number. The atoms of different isotopes can still be atoms of the same element. They differ only 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 or in a sentence with the atomic mass number following the symbol.

Example. The fissionable isotope of uranium is U-235. The nonfissionable isotope U-238 makes up most of naturally occurring uranium. Since uranium has the atomic number 92, a nucleus of U-235 contains 92 protons and 143 neutrons (235-92) while a nucleus of U-238 contains 92 protons and 146 neutrons(238-92). When used in a chemical equation the symbols are 23592U  and 23892U.

Although the atomic number of an atom is automatically understood in the element symbol, nuclear physicists and some chemists choose to write the atomic number deliberately in nuclear reactions. The atomic number can be written as a subscript preceding the element symbol, as 92U, but most regular chemists prefer to omit it when used in sentence form.

Look at a periodic table. The masses are not nice whole numbers! This is because there are many different isotopes of the same elements. For example there are two isotopes of chlorine; Cl-35 and Cl-37.

Stop here and do the exercise on Elemental Symbols, Fundamental Particles

Both atoms of an isotope of chlorine look the same, act and react the same. Lets say we had a beaker of pure Cl-35 and another beaker of pure Cl-37. They both have a pale yellow-green colour, they act like gases as they try to fill their container and they are both poisonous. But one beaker of atoms has two extra neutrons which makes it slightly heavier. The relative natural amounts of the isotopes and their molar masses for a few selected elements are given in the table below.

Most elements as they occur naturally on earth are mixtures of several isotopes. This is why the periodic table has the mass numbers it does.

Example. Chlorine 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%). It is safe to assume that no matter where you go the percentage abundance will be the same.

Atomic mass =  the sum of the percentages of each isotope of the element

              =  (75.77/100 * 35) +  (24.23/100 * 37)
      =  (0.7577 * 35) + (0.2423 * 37)

      =  26.5195 + 8.9651

      =  35.4846   This is the atomic mass of a sample of chlorine. It is a mixture of the masses of both isotopes.

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.

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        100.000         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:  If you know the relative amounts of each isotope the percentage of each can be calculcated. The relative percentage abundances of the isotopes Li-6 and Li-7 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 Li-6 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 Li-7 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 Li-6 is used in nuclear weapons, naturally occurring lithium can subjected to processing and the percentage of Li-6 is greater than Li-7 than that of the naturally occurring lithium. For this reason the relative abundances of lithium isotopes in lithium and lithium compounds which can be purchased may vary significantly due to the extraction of Li-6 by the governments of the United States and other nations for use in hydrogen fusion weapons.

Stop here and do the Isotopes Exercise