ORIGINS OF THE PERIODIC TABLE
By the early 1800s, study of the elements had produced a large collection of observations about their physical properties and chemical behavior. It was known that certain elements had similar properties. When the elements were listed in order of increasing atomic mass, the similarities recurred at definite intervals. In other words, the similarities recurred periodically and were called regularities.
An early periodic law proposed by the Russian chemist Dimitri Mendeleev expressed the regularities as a periodic function of atomic mass. Using this principle, Mendeleev developed the first version of the periodic table in 1869. Although this was a brilliant contribution to chemistry, it left many questions unanswered.
As a result of the work of Henry Moseley, a British physicist, we now know that it is more accurate to describe the regularities in physical and chemical properties as periodic functions, of atomic number. This revised version of Mendeleev’s law is known as the modern periodic law. Scientists also recognize that the chemical similarities among certain elements are due to similarities in the configurations {number and arrangement) of their valence electrons.
STRUCTURE OF THE PERIODIC TABLE
In the modem periodic table, the elements are placed in order of increasing atomic number. They are arranged in vertical columns and horizontal rows.
Groups. Elements with similar valence electron configurations fall into vertical columns known as groups, or families. For example, chlorine and fluorine, both with seven valence electrons, are in the same group called the halogens. These elements and other members of the group show similar properties because the arrangements of their valence electrons are similar; the characteristics of different groups are discussed later in the chapter.
Periods. Elements whose valence electrons have the same principal quantum number {the outermost principal energy level that contains electrons) fall in horizontal rows known as periods, or series. Many properties of the elements change systematically through a period. For example, in Period 2 {second horizontal row), note
the progression from active metal to metalloid (Li, Be, B) followed by a similar progression from least active nonmetal to most active nonmetal (C, N, 0, F). N, 0, and F are diatomic molecules (N2, 02, and F2). There are seven periods in the modern periodic table. For each element, the number of the period in which it is found corresponds to the principal quantum number of its valence electrons.
Trends in Properties. Although the chemical behavior of elements within a period varies because of variations in the electron configurations of the elements, certain generalizations about trends in properties within 3, period are significant. These trends are a function of increasing atomic number and changing valence electron configuration.
Covalent Atomic Radius. Atoms have no specific boundaries. It is therefore convenient to describe the sizes of atoms as they approach one another. The covalent atomic radius, or atomic radius, is one-half the distance between the nuclei of atoms that are joined together by a covalent bond in the solid phase. The relation between covalent atomic radius and atomic number can be interpreted in terms of nuclear charge and in terms of the arrangement of electrons in the orbitals of atoms. Atomic radii are usually measured in Angstrom units (Å). One Angstrom is equal to 10-10 meter.
For all elements within a period, the valence electrons (the electrons in the outermost energy level) are arranged around a kernel that contains the same number of energy levels. However, the number of valence electrons and the total number of protons in the nucleus are different for each element.
As elements are considered from left to right, the atomic number (protons or electrons) increases. Thus, within a period, nuclear charge increases because of the increasing number of protons in the atoms. This increase in positive charge causes the electrons to be attracted more closely to the nucleus. That is, the nucleus is not shielded sufficiently from the added electron repulsions. The increased attraction between the negative electrons and the positive nucleus is greater than any repulsion between the added electron and the other valence electrons. Thus, within a period, as atomic number increases, the covalent atomic radius decreases.
For all elements within a group, the atoms of each successive member have a larger kernel containing more filled energy levels, and the valence electrons are located at successively greater distances from the nucleus. The charge of the nucleus is shielded more and more as energy levels are added to successive members of the group. Thus, within a group, as atomic number increases, the covalent atomic radius increases.
The values for the atomic radii of metals are based on the distance between two bonded atoms in the solid state. For nonmetals, for the distance between two nonbonded atoms, the values given are the van der Waals radii.
Ionic Radius. When atoms form ions and ac- quire a charge, they gain or lose one or more electrons. The change in the number of electrons produces a corresponding change in the size of the electron cloud-the ionic radius. Since metal atoms become ions by losing electrons, the radii of these positive metal ions are smaller than the radii of the corresponding atoms in the ground state. This is because the loss of an electron reduces the repulsive forces, and the electron cloud shrinks. Nonmetal atoms become ions by gaining electrons. Therefore, the radii of negative non-metal ions are larger than the radii of the corresponding atoms in the ground state. This is because more electrons increase the repulsive forces, and the electron cloud expands.
Trends in Properties Within Periods and Groups
Within a period, as atomic number increases:
Covalent atomic radius decreases
Ionization energy increases
Electronegativity increases
Metallic character decreases
Within a group, as atomic number increases:
Covalent atomic radius increases
Ionization energy decreases
Electronegativity decreases
Metallic character increases
Classes of Elements. Elements are classified as metals, nonmetals, or metalloids. More than three-fourths of the elements are metals, approximately twenty are nonmetals, and the remaining few are metalloids.
In the periodic table, metallic properties are most pronounced in the elements found in the lower left comer. Thus, the Group I (alkali metal) family represents the most active metals. Nonmetallic properties are most pronounced in the elements in the upper right comer. Thus, the Group 17 (halogen) family represents the most active nonmetals. Metalloids are found just to the right of the heavy black dividing line between metals and nonmetals, which begins between boron and aluminum and proceeds stepwise down and to the right across groups 13 [IlIA] through 16 [VIA].
Metals. All metals are solid at room temperature except mercury, which is a liquid. Metals generally have a metallic, silvery luster. Metals are malleable {they can be hammered into thin sheets) and ductile {they can be drawn into thin wires). They form chemical bonds in such away that their valence electrons are relatively mobile {the electrons do not belong to a particular atom). This type of bonding, called metallic bonding, makes metals good conductors of heat and electricity. Metals have low ionization energies and low electronegativities; thus, they often act as electron donors in chemical reactions.
Nonmetals. Nonmetals are quite different from metals. Nonmetals tend to be gases, molecular solids, or network solids. In the solid phase, non- metals are brittle, they lack a luster, and they are poor conductors of heat and electricity. When nonmetals bond together, the bonds tend to be predominantly covalent in character. Because of their high ionization energies, nonmetals tend to act as electron acceptors in ionic compounds. They are highly electronegative.
Metalloids. A few elements have properties that are intermediate between those of metals and nonmetals. These elements are known as metalloids, or semimetals. Metalloids tend to form bonds that are partially ionic and partially covalent in character. The best-known metalloid is probably silicon, a semiconductor that is widely used in computers. Boron, arsenic, and tellurium are also metalloids.
CHEMICAL PROPERTIES WITHIN GROUPS
The elements in the periodic table are divided into 18 (or 8) groups on the basis of their electron configurations. Each group is assigned a group number. Under the recently adopted system used in this book, group numbers are Arabic numerals from I to 18. The group at the extreme left of the periodic table is Group 1; the group to its right is Group 2, and so on through Group 18.
In the system commonly used in the past, groups are designated by a combination of Roman numerals and letters.
Similarities Within Groups. Within each group, the elements exhibit related chemical and physical properties. As stated previously, the similarities in properties within a group are associated with similarities in the number of valence electrons.
The similarity in chemical properties within a group is reflected in the types of compounds formed by members of the group and is illustrated by their formulas. For example, the elements in Group I form chlorides with the general formula MCl and oxides with the general formula M2O, where M represents any member of the group. Elements in Group 2 form chlorides with the general formula MCl2 and oxides with the general formula MO.
The properties of elements in a group generally change progressively with increasing atomic number. Properties of the members of a group can be compared in terms of bonding, electronegativity, atomic size, and electron configuration.
Anomalies (exceptions) in the properties of elements within a group do occur. For example, in Group 13 (IlIA), boron does not form an ion as do other members of the group. The anomalies occur most frequently among the elements in Period 2 because in these atoms the valence electrons are relatively close to the nucleus and the two kernel electrons (the two electrons in the Is sublevel) provide a relatively small shielding effect; that is, the added repulsions provided by the two electrons are small.
In general, as the atomic number increases within a group, the radius of the atoms increases and the ionization energy of the elements de- creases. There is a corresponding decrease in electronegativity with increasing atomic number. The decreases in ionization energy and electro- negativity are due to: (1) the increased distance of the valence electrons from the nucleus, an9 (2) the increased shielding effect produced as a newly occupied energy level is added with each successive member of the group. Consistent with these changes in ionization energy and electro- negativity, each successive element within a group has increasingly metallic properties.
Groups 1 [IA] and 2 [IIA]. Groups 1 and 2 include the most reactive metals. Recall that members of Group I are called the alkali metals. (Hydrogen is not considered an alkali metal.) Members of Group 2 are called the alkaline earth metals. The Group I and Group 2 metals react with water to form bases (alkalies), which will be discussed in Unit 7. Because of their reactivity, these elements do not occur in nature in the un- combined state; that is, they occur only in chemical compounds. The elements in both groups undergo reduction to form the uncombined (free) metal by electrolysis of their fused compounds (see page 62).
In the elements of Groups I and 2, the valence shells are nearly empty. In the ground state, each element in Group I has one electron in its valence s orbital, while each element in Group 2 has two electrons in its valence s orbital. The elements in these groups form only ionic compounds because of their low ionization energies and electronegativities. They lose electrons readily, forming positive ions (cations) and relatively stable ionic compounds.
For metals, high reactivity is related to low ion-