Now that we have acquired some fundamental ideas about atomic structure, we can more thoroughly discuss the concept of chemical elements. All atoms of a particular element have the same atomic number, Z, and, conversely, all atoms with the same number of protons are atoms of the same element. The elements shown on the inside front cover have atomic numbers from Z = 1 to Z = 112. Each element has a name and a distinctive symbol. Chemical symbols are one- or two-letter abbreviations of the name (usually the English name). The first (but never the second) letter of the symbol is capitalized; for example: carbon, C; oxygen, O; neon, Ne; and silicon, Si. Some elements known since ancient times have symbols based on their Latin names, such as Fe for iron (ferrum) and Pb for lead (plumbum). The element sodium has the symbol Na, based on the Latin natrium for sodium carbonate. Potassium has the symbol K, based on the Latin kalium for potassium carbonate. The symbol for tungsten, W, is based on the German wolfram. Elements beyond uranium (Z = 922) do not occur naturally and must be synthesized in particle accelerators. Elements of the very highest atomic numbers have been produced only on a limited number of occasions, a few atoms at a time. Inevitably, controversies have arisen about which research team discovered a new element and, in fact, whether a discovery was made at all. However, international agreement has been reached on the first 112 elements; each one, except element 112, has an official name and symbol.
To represent the composition of any particular atom, we need to specify its number of protons (p), neutrons (n), and electrons (e). We can do this with the symbolism
This symbolism indicates that the atom is element E and that it has atomic number Z and mass number A. For example, an atom of aluminum represented as Al has 13 protons and 14 neutrons in its nucleus and 13 electrons outside the nucleus. (Recall that an atom has the same number of electrons as protons.)
Contrary to what Dalton thought, we now know that atoms of an element do not necessarily all have the same mass. In 1912, J. J. Thomson measured the mass-to-charge ratios of positive ions formed from neon atoms. From these ratios he deduced that about 91% of the atoms had one mass and that the remaining atoms were about 10% heavier. All neon atoms have 10 protons in their nuclei, and most have 10 neutrons as well. A very few neon atoms, however, have 11 neutrons and some have 12. We can represent these three different types of neon atoms as
Atoms that have the same atomic number (Z) but different mass numbers (A) are called isotopes. Of all Ne atoms on Earth, 90.51% are Ne. The percentages of Ne and Ne are 0.27% and 9.22%, respectively. These percentages 90.51%, 0.27%, 9.22% are the percent natural abundances of the three neon isotopes.
Sometimes the mass numbers of isotopes are incorporated into the names of elements, such as neon-20 (neon twenty). Percent natural abundances are always based on numbers, not masses. Thus, 9051 of every 10,000 neon atoms are neon-20 atoms. Some elements, as they exist in nature, consist of just a single type of atom and therefore do not have naturally occurring isotopes.* Aluminum, for example, consists only of aluminum-27 atoms.
When atoms lose or gain electrons, for example, in the course of a chemical reaction, the species formed are called ions and carry net charges. Because an electron is negatively charged, adding electrons to an electrically neutral atom produces a negatively charged ion. Removing electrons results in a positively charged ion. The number of protons does not change when an atom becomes an ion. For example, 22Ne+ and 22Ne2+ are ions. The first one has 10 protons, 10 neutrons, and 9 electrons. The second one also has 10 protons, but 12 neutrons and 8 electrons. The charge on an ion is equal to the number of protons minus the number of electrons. That is
Another example is the 16O2- ion. In this ion, there are 8 protons (atomic number 8), 8 neutrons (mass number - atomic number), and 10 electrons (18 - 10 = -22).
We cannot determine the mass of an individual atom just by adding up the masses of its fundamental particles. When protons and neutrons combine to form a nucleus, a very small portion of their original mass is converted to energy and released. However, we cannot predict exactly how much this so called nuclear binding energy will be. Determining the masses of individual atoms, then, is something that must be done by experiment, in the following way. By international agreement, one type of atom has been chosen and assigned a specific mass. This standard is an atom of the isotope carbon-12, which is assigned a mass of exactly 12 atomic mass units, that is, 12 u. Next, the masses of other atoms relative to carbon-12 are determined with a mass spectrometer. In this device, a beam of gaseous ions passing through electric and magnetic fields separates into components of differing masses. The separated ions are focused on a measuring instrument, which records their presence and amounts. Figure below illustrates mass spectrometry and a typical mass spectrum.
Although mass numbers are whole numbers, the actual masses of individual atoms (in atomic mass units, u) are never whole numbers, except for carbon-12. However, they are very close in value to the corresponding mass numbers, as we can see for the isotope oxygen-16. From mass spectral data the ratio of the mass of 16O to 12C is found to be 1.33291. Thus, the mass of the oxygen-16 atom is
1.33291 X 12 u = 15.9949 u
Which is very nearly equal to the mass number of 16.