Based on experimental data, it is known that chemical reactions involve only the electrons in atoms. In fact, only some of the electrons are involved. Because chemical properties are periodic, there must also be a periodic characteristic about electrons. This characteristic is the manner in which electrons are arranged in the atom. Electrons are in constant motion around the nucleus. They have both kinetic and potential energy, and their total energy is the sum of the two. The total energy is quantized; that is, there are definite, discrete values of total energy that atomic electrons can possess. These energy states can be visualized as spherical shells around the nucleus separated by forbidden areas where electrons cannot exist in a stable state. This sort of arrangement is illustrated in Figure 5.

Figure 5 - Electron Shells of Atoms

It is customary to speak of electron shells around the nucleus, and the shells are referred to by number. The first, or No. 1, shell is the one nearest the nucleus; the second, or No. 2, shell is next; then the third, or No. 3, shell; and so on in numerical order. In general, electrons closer to the nucleus have a lower energy state. Atomic electrons always seek the lowest energy state available.

The electron shells represent major energy states of electrons. Each shell contains one or more subshells called orbitals, each with a slightly different energy. In order of increasing energy, the orbitals are designated by the small letters s, p, d, f, g, h.

No two shells consist of the same number of orbitals. The first shell contains only one orbital, an s orbital. The second shell contains s and p orbitals. In general, each higher shell contains a new type of orbital:

the first shell contains an s orbital,

the second shell contains s and p orbitals, 

the third shell contains s, p, and d orbitals, 

the fourth shell contains s, p, d, and f orbitals,

and so on. Each orbital can hold a definite maximum number of electrons. There is also a limit to the number of electrons in each shell and the limit increases as one goes to higher shells. The numbers of electrons that can occupy the different orbitals and shells are shown in Table 4.

A more specific statement can now be made about which electrons are involved in chemical reactions. Chemical reactions involve primarily the electrons in the outermost shell of an atom. The term outermost shell refers to the shell farthest from the nucleus that has some or all of its allotted number of electrons. Some atoms have more than one partially-filled shell. All of the partially-filled shells have some effect on chemical behavior, but the outermost one has the greatest effect. The outermost shell is called the valence shell, and the electrons in that shell are called valence electrons. The term valence (of an atom) is defined as the number of electrons an element gains or loses, or the number of pairs of electrons it shares when it interacts with other elements.

The periodic chart is arranged so that the valence of an atom can be easily determined. For the elements in the A groups of the periodic chart, the number of valence electrons is the same as the group number; that is, carbon (C) is in Group IVA and has four valence electrons. The noble gases (Group 0) have eight in their valence shell with the exception of helium, which has two.

The arrangement in which the outermost shell is either completely filled (as with He and Ne) or contains eight electrons (as with Ne, Ar, Kr, Xe, Rn) is called the inert gas configuration. The inert gas configuration is exceptionally stable energetically because these inert gases are the least reactive of all the elements.

The first element in the periodic table, hydrogen, does not have properties that satisfactorily place it in any group. Hydrogen has two unique features: (a) the highest energy shell of a hydrogen atom can hold only two electrons, in contrast to all others (except helium) that can hold eight or more; and (b) when hydrogen loses its electron, the ion formed, H⁺, is a bare nucleus. The hydrogen ion is very small in comparison with a positive ion of any other element, which must still have some electrons surrounding the nucleus. Hydrogen can either gain or lose an electron. It has some properties similar to Group IA elements, and some similar to Group VIIA elements.

The number of electrons in the outer, or valence, shell determines the relative activity of the element. The elements are arranged in the periodic table so that elements of the same group have the same number of electrons in the outer shell (except for the Transition Groups). The arrangement of electrons in the outer shell explains why some elements are chemically very active, some are not very active, and others are inert. In general, the fewer electrons an element must lose, gain, or share to reach a stable shell structure, the more chemically active the element is. The likelihood of elements forming compounds is strongly influenced by this valence shell and on the stability of the resulting molecule. The more stable the molecules are, the more likely these molecules are to form.

Summary

The important information from this chapter is summarized below.

Periodic Table Summary

The subdivisions of the periodic table are periods, groups, and classes. The horizontal rows are called periods. The vertical columns are called groups. The entire table consists of three classes: metals, non-metals, and semimetals.

The subdivisions of the periodic chart have been explained such that the student should be able to identify them if given a periodic table.

Elements of the same group share certain physical and chemical characteristics. Examples of the characteristics of several groups are listed below.

Group 0 contains elements that are unreactive gases.

Group IA contains elements that are chemically active soft metals. Group VIIA contains elements that are chemically active nonmetals.

Groups 1B through VIIB and VIII are called transition groups and elements found in them display properties of metals.

The valence of an atom is defined as the number of electrons an element gains or loses, or the number of pairs of electrons it shares when it interacts with other elements.

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CHEMICAL BONDING

The development of matter, no matter what the form, is the result of the practical application of the assumptions, hypotheses, theories, and laws that chemists have formulated from their research into the nature of matter, energy, and change. This chapter will address some of these theories and laws. Chemical bonds and how atoms bond to form molecules will be discussed. In addition, an introduction to organic chemistry is provided.

EO 2.1 DEFINE the following terms:

a.Ionic bonds c.Covalent bonds

b.Van der Waals forcesd.Metallic bonds

EO 2.2 DESCRIBE the physical arrangement and bonding of a polar molecule.

EO 2.3 DESCRIBE the three basic laws of chemical reactions.

EO 2.4 STATE how elements combine to form chemical compounds.

EO 2.5 EXPLAIN the probability of any two elements combining to form a compound.

EO 2.6 DEFINE the following terms: