As noted at the beginning of the chapter, diamond is a hard, transparent solid graphite is a soft, black solid and the fullerenes have open cage structures. Now let’s apply this procedure to some particular compounds, beginning with one we have already discussed.Īllotropes of an element can have very different physical and chemical properties because of different three-dimensional arrangements of the atoms the number of bonds formed by the component atoms, however, is always the same. This will not change the number of electrons on the terminal atoms. If the central atom has fewer electrons than an octet, use lone pairs from terminal atoms to form multiple (double or triple) bonds to the central atom to achieve an octet. We explain in Section 8.6 'Exceptions to the Octet Rule' that some atoms are able to accommodate more than eight electrons.Ħ. If any electrons are left over, place them on the central atom. These electrons will usually be lone pairs.ĥ. Beginning with the terminal atoms, add enough electrons to each atom to give each atom an octet (two for hydrogen). In H 2O, for example, there is a bonding pair of electrons between oxygen and each hydrogen.Ĥ. Place a bonding pair of electrons between each pair of adjacent atoms to give a single bond. For CO 3 2−, for example, we add two electrons to the total because of the −2 charge.ģ. (Recall from Chapter 6 'The Structure of Atoms' that the number of valence electrons is indicated by the position of the element in the periodic table.) If the species is a polyatomic ion, remember to add or subtract the number of electrons necessary to give the total charge on the ion. Add together the valence electrons from each atom. Determine the total number of valence electrons in the molecule or ion. The central atom is usually the least electronegative element in the molecule or ion hydrogen and the halogens are usually terminal.Ģ. The shapes of the energy versus distance curves in the two figures are similar because they both result from attractive and repulsive forces between charged entities. Notice the similarity between Figure 8.9 'A Plot of Potential Energy versus Internuclear Distance for the Interaction between Two Gaseous Hydrogen Atoms' and Figure 8.1 'A Plot of Potential Energy versus Internuclear Distance for the Interaction between a Gaseous Na', which described a system containing two oppositely charged ions. Thus at intermediate distances, proton–electron attractive interactions dominate, but as the distance becomes very short, electron–electron and proton–proton repulsive interactions cause the energy of the system to increase rapidly. At the observed bond distance, the repulsive and attractive interactions are balanced.Ī plot of the potential energy of the system as a function of the internuclear distance ( Figure 8.9 'A Plot of Potential Energy versus Internuclear Distance for the Interaction between Two Gaseous Hydrogen Atoms') shows that the system becomes more stable (the energy of the system decreases) as two hydrogen atoms move toward each other from r = ∞, until the energy reaches a minimum at r = r 0 (the observed internuclear distance in H 2 is 74 pm). Hence the quantum mechanical probability distributions must be used.įigure 8.8 Attractive and Repulsive Interactions between Electrons and Nuclei in the Hydrogen MoleculeĮlectron–electron and proton–proton interactions are repulsive electron–proton interactions are attractive. Recall from Chapter 6 'The Structure of Atoms' that it is impossible to specify precisely the position of the electron in either hydrogen atom. The electron in one atom is attracted to the oppositely charged proton in the other atom and vice versa ( E Similarly, the protons in adjacent atoms repel each other ( E > 0).The electrons in the two atoms repel each other because they have the same charge ( E > 0).As the two hydrogen atoms are brought together, additional interactions must be considered ( Figure 8.8 'Attractive and Repulsive Interactions between Electrons and Nuclei in the Hydrogen Molecule'): Each hydrogen atom in H 2 contains one electron and one proton, with the electron attracted to the proton by electrostatic forces.

We begin our discussion of the relationship between structure and bonding in covalent compounds by describing the interaction between two identical neutral atoms-for example, the H 2 molecule, which contains a purely covalent bond. To understand the concept of resonance.To use Lewis dot symbols to explain the stoichiometry of a compound.

Perchlorate ion bonding pairs