A cornerstone of modern chemistry is the arrangement of elements into groups with similar chemical properties. Scientists of the 1700’s and 1800’s collected facts about the elements then known. Soon, similarities among some of these elements were recognized. By the second half of the 1800’s, scientists had enough information to begin classifying all the elements.
, Avogadro’s hypothesis had paved the way for determining relative atomic weights, making it possible to classify all the known elements according to their properties and to list them by weight, from lightest to heaviest. Dmitri Mendeleev, a Russian chemist, began this classification in 1869. By arranging the elements in columns, the forerunner of today’s periodic table, Mendeleev discovered the natural order that exists among elements. Listing the 57 known to him, Mendeleev noticed that certain chemical properties reappeared in every eighth element. This repetition at regular intervals came to be known as the periodic law.
Mendeleev’s table not only listed the elements known in his time—it also left gaps where the periodic law predicted that elements of a certain weight with known properties would appear. He predicted the properties of three such unknown elements. Toward the end of the century, when the “missing” elements—gallium, scandium, and germanium— were discovered, scientists had all the proof they needed of the validity of Mendeleev’s grand concept.
Regularities
When Mendeleev began assembling his chart of the elements, scientists knew nothing about the parts of the atom. It is much easier for today’s scientists to see why elements fall into regular groupings because of their knowledge of atomic structure—how electrons are distributed about the nucleus and subatomic particles are arranged within it.
Unlike Mendeleev’s chart, which ordered elements according to their atomic weight, today’s periodic table arranges the elements by atomic number—the number of protons in the nucleus of the atom. The number of protons determines the number of electrons, and the number and arrangement of electrons are the surest guides to an element’s chemical behaviors. The table lists elements by increasing atomic number, from left to right, in seven horizontal rows called periods. Arranged this way, sodium is the first element to repeat the chemical properties of another element (lithium), so sodium starts the second period. Potassium repeats sodium’s chemical properties, so potassium starts the third period.
The table is like a calendar with eight-day weeks. Instead of a day like Monday or Tuesday repeating itself at regular intervals, a chemical property is repeated. Sodium and potassium, for example, both of which have only one electron in their outer shells, are metals that react violently with water, and both easily lose that valence electron to form ionic bonds with other elements. Despite the similarities, the heavier element, potassium, has one more complete electron shell between its outer, valence electron and its nucleus. As a result, the valence electron is not held as strongly, which makes potassium even more chemically reactive than sodium.
Electrons of an atom are arranged in shells A shell is a group of electrons at the same average distance from the atomic nucleus, and the outermost shell contains the electrons involved in chemical bonding—valence electrons. Having arranged all the elements in horizontal rows according to their atomic numbers, and having started a new row with the first repetition of chemical properties, it is apparent how the similarities in every eighth element give rise to a series of columns (1-8) in which groups of related elements appear.
Elements grouped vertically in the table have the same number of electrons in their outer shells and therefore have similar characteristics. All the elements in the first column (1 A) have a single valence electron in the outer shells of their atoms, and all the elements of the second column (2A) have two valence electrons in their outer shells. This pattern continues until the last column (8A), where all the elements except helium have eight electrons in their outer shells.
While the inner shell of an atom can hold only two electrons, the outer shell can hold up to eight, and when an outer shell holds a full complement, the atom is stable, inert, and un-reactive—just like the six gases in the eighth column of the table. But these six gases are unique among the 109 elements.
All of the other elements are chemically reactive to some degree. Elements with fewer than four valence electrons tend to donate those electrons when forming chemical bonds—the essential characteristic of a metallic element Elements with more than four valence electrons tend to accept electrons from other elements when forming a chemical bond—the essential characteristic of a non-metal element An element like carbon, with just four valence electrons, can go either way— either accepting or giving.
Elements whose properties are intermediate between those of metals and nonmetals— metalloids— include a class called semiconductors: silicon (Si), germanium (Ge), and antimony (Sb). Like nonmetals, these three elements do not conduct electrical current at room temperature. But when heated they acquire a metallic characteristic—they conduct electricity.
Bearing in mind that an element’s identity as a metal or nonmetal depends mainly on whether it gives or receives electrons, look again at the periodic table. The electron givers—the metals—cluster toward the left side of the table, and the electron acceptors cluster toward the right. Running diagonally down the dividing line between the metals and nonmetals are the metalloids: boron (B), silicon (Si), arsenic (As), tellurium (Te), and astatine (At). The most active metals, those with the fewest valence electrons, are on the far left The most active nonmetals, elements just one or two electrons short of a full complement of eight, are on the far right (disregarding the last column, the rare gases).
