Explain the bonding nature of ionic compounds. Relating microscopic bonding properties to macroscopic solid properties.

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The substances described in the preceding discussion are composed of molecules that are electrically neutral; that is, the number of positively-charged protons in the nucleus is equal to the number of negatively-charged electrons. In contrast, ions are atoms or assemblies of atoms that have a net electrical charge. Ions that contain fewer electrons than protons have a net positive charge and are called cations. Conversely, ions that contain more electrons than protons have a net negative charge and are called anions. Ionic compounds contain both cations and anions in a ratio that results in no net electrical charge.

In covalent compounds, electrons are shared between bonded atoms and are simultaneously attracted to more than one nucleus. In contrast, ionic compounds contain cations and anions rather than discrete neutral molecules. Ionic compounds are held together by the attractive electrostatic interactions between cations and anions. In an ionic compound, the cations and anions are arranged in space to form an extended three-dimensional array that maximizes the number of attractive electrostatic interactions and minimizes the number of repulsive electrostatic interactions (Figure $$\PageIndex{1}$$). As shown in Equation $$\ref{Eq1}$$, the electrostatic energy of the interaction between two charged particles is proportional to the product of the charges on the particles and inversely proportional to the distance between them:

\< \text {electrostatic energy} \propto {Q_1Q_2 \over r} \label{Eq1}\>

where $$Q_1$$ and $$Q_2$$ are the electrical charges on particles 1 and 2, and $$r$$ is the distance between them. When $$Q_1$$ and $$Q_2$$ are both positive, corresponding to the charges on cations, the cations repel each other and the electrostatic energy is positive. When $$Q_1$$ and $$Q_2$$ are both negative, corresponding to the charges on anions, the anions repel each other and the electrostatic energy is again positive. The electrostatic energy is negative only when the charges have opposite signs; that is, positively charged species are attracted to negatively charged species and vice versa.

api/deki/files/128311/clipboard_eb3eac2b922a33e35b9db86e87afa383b.png?revision=1" />Figure $$\PageIndex{2}$$: The Effect of Charge and Distance on the Strength of Electrostatic Interactions. As the charge on ions increases or the distance between ions decreases, so does the strength of the attractive (−…+) or repulsive (−…− or +…+) interactions. The strength of these interactions is represented by the thickness of the arrows.

If the electrostatic energy is positive, the particles repel each other; if the electrostatic energy is negative, the particles are attracted to each other.

One example of an ionic compound is sodium chloride (NaCl; Figure $$\PageIndex{3}$$), formed from sodium and chlorine. In forming altoalsimce.orgical compounds, many elements have a tendency to gain or lose enough electrons to attain the same number of electrons as the noble gas closest to them in the periodic table. When sodium and chlorine come into contact, each sodium atom gives up an electron to become a Na+ ion, with 11 protons in its nucleus but only 10 electrons (like neon), and each chlorine atom gains an electron to become a Cl− ion, with 17 protons in its nucleus and 18 electrons (like argon), as shown in part (b) in Figure $$\PageIndex{1}$$. Solid sodium chloride contains equal numbers of cations (Na+) and anions (Cl−), thus maintaining electrical neutrality. Each Na+ ion is surrounded by 6 Cl− ions, and each Cl− ion is surrounded by 6 Na+ ions. Because of the large number of attractive Na+Cl− interactions, the total attractive electrostatic energy in NaCl is great.

Figure $$\PageIndex{3}$$: Sodium Chloride: an Ionic Solid. The planes of an NaCl crystal reflect the regular three-dimensional arrangement of its Na+ (purple) and Cl− (green) ions.

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Consistent with a tendency to have the same number of electrons as the nearest noble gas, when forming ions, elements in groups 1, 2, and 3 tend to lose one, two, and three electrons, respectively, to form cations, such as Na+ and Mg2+. They then have the same number of electrons as the nearest noble gas: neon. Similarly, K+, Ca2+, and Sc3+ have 18 electrons each, like the nearest noble gas: argon. In addition, the elements in group 13 lose three electrons to form cations, such as Al3+, again attaining the same number of electrons as the noble gas closest to them in the periodic table. Because the lanthanides and actinides formally belong to group 3, the most common ion formed by these elements is M3+, where M represents the metal. Conversely, elements in groups 17, 16, and 15 often react to gain one, two, and three electrons, respectively, to form ions such as Cl−, S2−, and P3−. Ions such as these, which contain only a single atom, are called monatomic ions. The charges of most monatomic ions derived from the main group elements can be predicted by simply looking at the periodic table and counting how many columns an element lies from the extreme left or right. For example, barium (in Group 2) forms Ba2+ to have the same number of electrons as its nearest noble gas, xenon; oxygen (in Group 16) forms O2− to have the same number of electrons as neon; and cesium (in Group 1) forms Cs+, which has the same number of electrons as xenon. Note that this method is ineffective for most of the transition metals. Some common monatomic ions are listed in Table $$\PageIndex{1}$$.

Table $$\PageIndex{1}$$: Some Common Monatomic Ions and Their Names Group 1Group 2Group 3Group 13Group 15Group 16Group 17
Li+ lithium Be2+ beryllium N3− nitride (azide) O2− oxide F− fluoride
Na+ sodium Mg2+ magnesium Al3+ aluminum P3− phosphide S2− sulfide Cl− chloride
K+ potassium Ca2+ calcium Sc3+ scandium Ga3+ gallium As3− arsenide Se2− selenide Br− bromide
Rb+ rubidium Sr2+ strontium Y3+ yttrium In3+ indium Te2− telluride I− iodide
Cs+ cesium Ba2+ barium La3+ lanthanum