Major groups of elements

Few elements in their pure formoccur free in na­ture. Exceptions are unreac-tive metals such as gold (left) and silver, and the non-metals carbon and sulfur (right). Nearly all other elements—apart from the gases in the atmosphere-occur in combination as compounds.

The basis of modern chemistry is the arrang­ing of elements into groups that show similar behavior. During the eighteenth and nine­teenth centuries, scientists gathered factual in­formation about the different elements known to them at that time. Soon, similarities began to appear among some of these elements. By the second half of the nineteenth century, enough information was available to make possible a serious attempt at classifying all the elements.

A number of attempts were made at classi­fying them. The classifications were based on the empirical method. That is, scientists simply observed the elements and how they behaved and noted the similarities among them. The best such classification was done by the Rus­sian chemist Dmitri Mendeleev in 1869. His periodic table included not only the elements known in his time, but also left blank spaces for unknown elements. Later, using the peri­odic law, he predicted the properties of three unknown elements. It was the discovery of some of these "missing" elements that led sci­entists to accept his basic scheme.

Regularities

Mendeleev developed the periodic table with­out knowing about the parts of the atom. With a modern knowledge of the structure of elec­trons and other subatomic particles, it is easy to see why elements fall into regular group­ings (called periods). The Russian chemist's classification of the elements has proven very accurate.

The periodic table lists the elements hori­zontally in order of increasing atomic number. Thus, each element has one more proton in its nucleus than the element on its left, but one less than the element on its right. For example, copper (Cu) has 29 protons. On its left, nickel

(Ni) has 28. On its right, zinc (Zn) has 30.

Vertically, the periodic table lists elements whose atoms have the same number of outer electrons. The total number of electrons, how­ever, varies considerably.

As the atomic number increases, so does the number of electrons surrounding the nu­cleus. A shell is a group of electrons at the same average distance from the nucleus. At particular intervals—specified by a set of laws called quantum theory—a shell of electrons is completed. It can hold no more electrons. This makes the structure stable, as we have already seen from the formation of molecules. Further­more, the extra electron of the next element starts a new shell. It has a single electron in this shell that is relatively easy to remove.

For example, sodium (Na) and potassium (K) are two elements that have a single electron in their outer shells. Both are metals that react vi­olently with water. Both easily lose that single electron in the outer shell and combine with other elements to form ionic (electrically charged) compounds, such as chlorides (so­dium chloride, potassium chloride). Neverthe­less, they are different elements for potassium has one more complete shell of electrons be­tween its outer electron and its nucleus, which also has eight more protons than sodium. This makes potassium's properties different in other details from those of sodium.

If we consider the elements with one more proton and electron than potassium or so­dium, namely magnesium (Mg) and calcium (Ca), we find that both form ionic compounds by the loss of two electrons. This is also the total number of electrons in the outer shell.

Thus, the elements are grouped vertically to form families with certain common charac­teristics. These characteristics are based on an equal number of electrons in the outer shell.

Major groups of elements 23

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1.5

1.0

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He

~*F

Periodicity,the tendency for certain physical proper­ties to repeat themselves at regular intervals, occurs with the elements. Here, atomic radius is plotted in the vertical rise against the horizontal lay of atomic number. The first 83 ele­ments (to bismuth) are plot­ted. The peaks correspond to the rare gases and alkali metals. The troughs corre­spond to the halogens.

—1_ 20

—j_ 30

—p.40

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75 80 Atomic number

Quantum numbers

The electrons in each element are character­ized by four different quantum numbers. Only the first two will be discussed here. The first number describes the average distance of an electron from the nucleus. All the electrons at a particular distance from the nucleus form a shell. The second quantum number (actually a letter) defines the shape of the atomic orbital (the path followed by an electron around the nucleus). There are pictures of the various or-bitals on pages 12 and 13.

The positioning of electrons around an atom is often referred to in a type of short­hand. The first quantum number is indicated by a number. The second is denoted by a let­ter (s, p, or d). The number of electrons in a particular orbital is shown as a superscript.

What the quantum rules say is that when the first quantum equals 1, only an s orbital is

H

1.0079

Be

9.0128

V

50.9415

Nb

92.9064

Ta

180.948

allowed. When this quantum number equals 2, both s and p orbitals are allowed. And when it is 3, there can be s, p, and d orbitals. The quantum rules also state the maximum num­ber of electrons that can occupy an orbital (2 electrons for an s orbital, 6 for p, and 10 for d).

For example, the element Argon (Ar) has 18 electrons, written out as 1 s 2s 2p 3s 3p. This means that there are two electrons in an s or­bital in the first shell. In the second shell, there are two electrons in an s orbital and six elec­trons in a p orbital. In the third shell, there are also two electrons in an s orbital and, again, six in a p orbital.

In the following sections, the elements are divided into small family groups. Each part be­gins with an illustration of the section of the periodic table that includes the element or el­ements being described.

N

14.0067

10.81

2B

Co

58.9332

Rh

102.906

lr

192.22

Date: 2015-12-11; view: 1639