If the atomic weight of some element happens to be very close to a whole number to begin with, it may consist of a single kind of atom. For instance, the gas fluorine (chemical symbol F) has an atomic weight of nearly 19, while that of the metal sodium (chemical symbol Na) is nearly 23. As it turns out, all the atoms of fluorine are of the single variety ¹⁹F, while all the atoms of sodium are ²³Na. Sometimes the atomic weight of an element, as it occurs in nature, is nearly a whole number and yet it is made up of more than 1 isotope. In that case, one of the isotopes makes up very nearly all of it, while the others are present in such minor quantities that the average is hardly affected. Helium, for instance (atomic symbol He) has an atomic weight of just about 4 and, indeed, almost all the atoms making it up are ⁴He. However, 0.0001% of the atoms, or one out of a million, are ³He. Again, 99.6% of all the nitrogen atoms (atomic symbol N) are ¹⁴N, but 0.4% are ¹⁵N. Then, 98.9% of all carbon atoms (atomic symbol C) are ¹²C, but 1.1% are ¹³C. It is not surprising that the atomic weights of nitrogen and carbon are just about 14 and 12, respectively. [Illustration: _Harold Urey_] Even hydrogen does not escape. Its atomic weight is just about 1 and most of its atoms are ¹H. The American chemist Harold Clayton Urey (1893- ) detected the existence of a more massive isotope, ²H. This isotope has almost twice the mass of the lighter one. No other isotopes of a particular atom differ in mass by so large a factor. For that reason ²H and ¹H differ in ordinary chemical properties more than isotopes usually do and Urey therefore gave ²H the special name of “deuterium” from a Greek word meaning “second”. [Illustration: _W. F. Giauque_] In 1929 the American chemist William Francis Giauque (1895- ) found that oxygen was composed of more than 1 isotope. Its atomic weight had been set arbitrarily at 16.0000 so it was a relief that 99.76% of its atoms were ¹⁶O. However, 0.20% were ¹⁸O, and 0.04% were ¹⁷O. As you see, ¹⁶O must have a mass number of slightly less than 16.0000 and it must be the more massive isotopes ¹⁷O and ¹⁸O that pull the average up to 16.0000. Disregarding this, chemists clung to a standard atomic weight of 16.000 for oxygen as it appeared in nature, preferring not to concern themselves with the separate isotopes. Physicists, however, felt uneasy at using an average as standard for they were more interested in working with individual isotopes. They preferred to set ¹⁶O at 16.0000 so that the average atomic weight of oxygen was 16.0044 and all other atomic weights rose in proportion. Atomic weights determined by this system were “physical atomic weights”. Finally, in 1961, a compromise was struck. Chemists and physicists alike decided to consider the atomic weight of ¹²C as exactly 12 and to use that as a standard. By this system, the atomic weight of oxygen became 15.9994, which is only very slightly less than 16.