The atoms of a compound are held together by chemical bonds formed by the interaction of electrons from each atom. According to the octet rule Section 5.7C1, atoms bond together to form molecules in such a way that each atom participating in a chemical bond acquires an electron configuration resembling that of the noble gas nearest it in the periodic table. Thus the outer shell of each bonded atom will contain eight electrons (or two electrons for hydrogen and lithium). The simplest chemical bond is that formed between two hydrogen atoms. Each hydrogen atom has one electron. As the two atoms approach each other, the nucleus of one atom attracts the electron of the other. Eventually the two orbitals overlap, becoming a single orbital containing two electrons (see Figure 7.1). FIGURE 7.1 Two hydrogen atoms, each with one electron, combine to form a hydrogen molecule, in which the two electrons are shared between the atoms and serve to give each atom a filled valence shell. This orbital encompasses space around both nuclei. Although the electrons may be in any part of this orbital, we can predict that they are most likely to be in the space between the nuclei, shielding one nucleus from the other and being attracted by both. In the resulting molecule, both atoms have two electrons and a filled outer (valence) shell. These shared electrons form a bond between the two atoms. This chemical bond is a covalent bond, a pair of electrons shared between two atoms. When this bond forms, energy is released. This release of energy shows that the molecule of hydrogen is more stable than the separate atoms. A. Covalent, Polar covalent, and Ionic Bonds Because the hydrogen molecule contains two identical atoms, it can be assumed that the bonding electrons in this covalent bond are shared equally by these atoms. Most chemical bonds are not between like atoms but form between atoms of different elements. These bonds are slightly different from that in a hydrogen molecule. Consider the bond between hydrogen and chlorine: Again both atoms require one more electron to satisfy the octet rule. As the atoms come together, their orbitals overlap and the two atoms share a pair of electrons. However, the hydrogen-chlorine bond differs from the hydrogen-hydrogen bond because the electrons are not shared equally between hydrogen and chlorine but are more strongly attracted to the chlorine. They are more apt to be found close to the chlorine than close to the hydrogen. Because of this unequal sharing, the chlorine atom assumes a slightly negative character and the hydrogen atom a slightly positive character. We say that the bond is polar covalent, meaning that the bond consists of electrons shared between two atoms (therefore covalent) but shared unequally, thus giving the bond a positive and a negative end, a condition described by the term polar. We can also say that the bond is a dipole or has a dipole moment, meaning that the bond has a positive end (the hydrogen) and a negative end (the chlorine). The more negative atom in a bond is often shown with the symbol and the more positive atom is shown with the symbol. The bond between hydrogen atoms is nonpolar (has no positive and negative ends) covalent (electrons are shared). An ionic bond is the extreme case of a polar covalent bond. In an ionic bond, the bonding atoms differ so markedly in their attraction for electrons that one or more electrons are essentially transferred from one atom to the other. The sodium-chlorine bond is an example of an ionic bond. The attraction of the chlorine atom for electrons is so much greater than that of a sodium atom that the 3s electron of sodium is said to be completely transferred from sodium to chlorine. In summary, then, the three types of bonds are: (1) a covalent bond, in which the electrons are shared equally; (2) a polar covalent bond, in which the electrons are shared unequally; and (3) an ionic bond, in which electrons are transferred from one atom to the other. These bonds are illustrated in Figure 7.2. FIGURE 7.2 Electrons in nonpolar covalent, polar covalent, and ionic bonds: (a) the electrons are shared equally; (b) the electrons are held closer to the more-negative chlorine atom; (c) one electron has been transferred from sodium to chlorine. B. Predicting Bond Type; Electronegativity It is possible to predict the type of bond that will form between two elements. The farther apart (left to right) the two elements are in the periodic table, the more ionic and the less covalent will be the bond between them. Thus, metals react with nonmetals to form ions joined predominantly by ionic bonds. Bonds with the highest degree of ionic character are formed by the reaction of alkali or alkaline earth metals with the halogens, particularly with fluorine or chlorine. Nonmetals react together to form covalent bonds. If the bond is between two neighbors in the table, the bond will be less polar than if the nonmetals are separated by other element. For example, carbon and nitrogen are in neighboring columns, and carbon and fluorine are in Groups 4 and 7, respectively. A carbon-nitrogen bond will be less polar than a carbon-fluorine bond. Finally, if the two atoms are of the same element, as in a hydrogen molecule or a chlorine molecule, the bond will be essentially nonpolar. The concepts in the previous paragraph have been quantified by the concept of electronegativity. The electronegativity (EN) of an element measures its attraction for the electrons in a chemical bond. One scale of electronegativity was developed by the American chemist Linus Pauling (b. 1901). On this scale, fluorine, the most electronegative element, has an electronegativity of 4.0. Carbon has an electronegativity of 2.5, hydrogen, 2.1, and sodium 0.9. Figure 7.3 shows the electronegativities of the elements with which we deal most often.
FIGURE 7.3 Electronegativities of some elements (Pauling scale). Notice that the electronegativity of most metals is close to 1.0 and that the electronegativity of a nonmetal, although dependent on its location in the table, is always greater than 1.0. In general, electronegativity increases from bottom to top in a column and from left to right across a period. Note that the noble gases, Group 8, do not appear in this table. Electronegativity measures the relative attraction of atoms for electrons in chemical bonds. The noble gases react differently from the halogens and other nonmetals.
The concepts of electronegativity do not apply to them. When two atoms combine, the nature of the bond between them is determined by the difference between their electronegativities (denoted EN). If the atoms forming the bond differ in electronegativity by more than 1.7 units, the bond will be at least 50% ionic (referred to as percent ionic character); we treat such a bond as wholly ionic. If the values differ by less than 0.4 units, we consider the bond to be wholly nonpolar. If the difference is between 0.4 and 1.7 electronegativity units, the bond is considered to be polar covalent. Remember that electronegativities have been calculated from fairly imprecise data for particular bonding situations. Electronegativity is useful in predicting the nature of a bond and for comparing bond types, but the prediction is only an approximation. Remember too that no sharp distinction exists between ionic, polar covalent, and nonpolar bonds; rather, they form a continuum. Even the most ionic bond (between cesium and fluorine) has some covalent character, and only bonds between atoms of the same element have no ionic character. In these bonds, the atom with the higher electronegativity will be the negative end of the bond and, in extreme situations, will become the negative ion. To show these partial charges on a polar covalent bond, we mark the positive end of the bond with a and the negative end of the bond with a . Table 7.1 summarizes these data. TABLE 7.1 Guidelines for predicting bond type from electronegativity data
Difference in electronegativity ( EN)
Type of bond predominant Example EN More positive atom > 1.7 ionic NaCl 2.1 sodium 0.4 – 1.7 polar covalent C-Cl 1.5 carbon covalent H-H 0.0 neither C – H 0.4 neither
Predict the nature of the bond between the following pairs of atoms as predominantly nonpolar covalent, polar covalent, or ionic. For each polar covalent bond, use a small Greek letter
to show which atoms bears a partial positive charge (
) and which a partial negative charge (
a. S-O b. C-O c. Al-F
Solution a. The electronegativity of oxygen is 3.5 and that of sulfur is 2.5 The difference is 1.0 unit; we predict the S-O bond to be polar covalent. The oxygen is partially negative, and the sulfur is partially positive, so we write: .
b. The electronegativity difference between oxygen and carbon is 1.0 unit ( 3.5 – 2.5). Therefore, we predict the C – O bond to be polar covalent. Because oxygen is the more electronegative of the two, it carries the negative charge.
c. The electronegativity difference between fluorine and aluminum is 2.5 units (4.0 – 1.5). Therefore, we predict the Al – F bond to be largely ionic. The aluminum forms a cation, the fluoride an anion.
C. Single, Double, and Triple Bonds A covalent bond represents the sharing of electrons between two atoms. Single bonds result from the sharing of a single pair of electrons. The covalent bonds shown in Figure 7.2 are single bonds. Usually, as in the hydrogen molecule, each atom forming the bond contributes one electron to the bond. Sometimes, as in the reaction of ammonia, NH3, with a hydrogen ion, H+, to form the ammonium ion, NH4+, both electrons come from the same atom: It is common practice to use a dash to represent a pair of electrons. In this text we will use dashes for shared electrons and dots for unshared (lone-pair) electrons. With this notation, the above equation is written: In the ammonia molecule, the nitrogen shares a pair of electrons with each of the three hydrogens. In each bond, one electron comes from nitrogen and one from hydrogen. The nitrogen still has an unshared pair of electrons. A hydrogen ion has no electrons; the single hydrogen electron was lost when the atom became an ion and gained a positive charge. When the hydrogen ion bonds to the ammonia molecule, both electrons of the bond come from the nitrogen. A bond in which one atom has donated both electrons is often referred to as a coordinate covalent bond. It is most important to realize that the different name refers only to the method of formation. Once the ammonium ion is formed, all hydrogen-nitrogen bonds in the ion are equivalent. Notice, too, that the entire ammonium ion now carries a positive charge, denoted by placing brackets around the ion and writing a superscript +. In addition to single bonds, there are double bonds and triple bonds. A double bond represents the sharing of four electrons by two atoms. The bond between carbon and oxygen is often a double bond, as in formaldehyde, CH2O. Here carbon is singly bonded to each of the hydrogens and doubly bonded to oxygen. Of this double bond, two electrons have come from carbon and two from oxygen. The single carbon-hydrogen bonds are nonpolar ( EN = 0.4); the double carbon-oxygen bond is polar covalent ( EN = 1.0). Note that each atom in the diagram of formaldehyde now follows the octet rule. Each hydrogen has two electrons; the carbon and the oxygen have eight electrons each. Notice too that the oxygen has two pairs of unshared electrons. Such an unshared pair is sometimes known as a lone pair. We will see that the negative end of a polar bond often holds unshared electron pairs. A triple bond is formed when two atoms share six electrons (three pairs). The nitrogen molecule contains a triple bond. Its structure is Each nitrogen donates three electrons to the bond and retains a lone pair. Back   Home   Next