SELECTED PRINCIPLES: INTRODUCTION - IONIC BONDING Bonding between atoms, which occurs because the resulting molecule or compound has a lower energy than its constituent atoms, is achieved by redistributing the valence (or bonding) electrons. In ionic bonding, this redistribution occurs by the atoms transferring one or more electrons. The term ionic bond describes the electrostatic attraction of two oppositely charged ions in a crystalline lattice.
The Mechanism of Ionic Bonding Formally, the ionic compound sodium fluoride - which contains Na1+ and F1- ions, each with an octet of electrons - results from the transfer of one valence electron from a sodium atom to a fluorine atom; i.e., Na atom + F atom 覧覧覧覧覧ｮ Na1+ion + F1-ion 2,8,1 2,7 2,8 2,8 Despite the attractiveness of such a simple summary, one needs to peer into the mechanism of ionic bonding in order to prevent misconceptions. The formation of solid sodium fluoride, NaF(s), from solid sodium and gaseous difluorine, can be divided into five ergonic processes; i.e., Na(s) 覧覧覧覧覧ｮ Na(g) DH1 Na(g) 覧覧覧覧覧ｮ Na1+(g) + 1e- DH2 F2(g) 覧覧覧覧覧ｮ 2F(g) DH3 F(g) + 1e- 覧覧覧覧覧ｮ F1-(g) DH4 Na1+(g) + F1-(g) 覧覧覧覧覧ｮ Na1+(s)F1-(s) DHL Experiments show that formation of these gaseous ions is endothermic; i.e., SDHI, the value of DH1 + DH2 + DH3 + DH4, is positive. In sharp contrast, formation of the solid from these gaseous ions is exothermic; i.e., the value of DHL, known as the lattice energy, is negative. Solid sodium fluoride forms because the value of DH, DHL - SDHI, is negative; so, the 'driving force behind' the formation of sodium fluoride is the overwhelmimg exothermicity of the attractions of the oppositely charged ions for each other (... and not the formation of ions with an octet). Furthermore, because experiments have shown that, for every combination of metallic and non-metallic element, the corresponding value of SDHI is positive, the formation of an ionic compound is determined largely by the value of DHL (the lattice energy).
This ('long-winded') approach above allows a qualitative understanding of some apparent anomalies; two illustrative examples are given here. Firstly, Na2O, MgO, Al2O3, NaCl, and MgCl2 are all ionic compounds in which the ions have an octet: but, although aluminium and chlorine can readily form ions with an octet, aluminium chloride is a covalent compound. * The explanation is as follows: DHL - SDHI is negative in each of the first five compounds, but DHL - SDHI is positive for the hypothetical ionic compound Al3+(s)3Cl1-(s)
* When gaseous aluminium chloride is cooled, discrete AlCl3 molecules dimerize to form Al2Cl6 (see the electron-structure diagram above).
And secondly, NaCl and MgCl2 are both ionic compounds in which the ions have an octet: but, although transition elements do not form ions with an octet, CuCl2 and FeCl2 are also ionic. The explanation is as before; thus, DHL - SDHI is negative in each of these four compounds.
In the absence of either experimental or calculated data, the student will ask (or even demand) to know how one predicts the type of bonding in any given compound. Although there is no predictive method which is foolproof, two rules of thumb are useful: first, the bonding between non-metallic elements is normally covalent; and second, the bonding between metallic and non-metallic elements is commonly ionic. *
The Structures and Physical Properties of Ionic Compounds The diagram below shows the structure of NaCl; the oppositely charged ions are arranged in a symmetrical pattern within a crystal lattice.
The electrostatic attractions between the oppositely charged ions in the crystal lattices of ionic compounds are strong and omnidirectional: and so, such compounds show several characteristic physical properties; however, only three of these will be considered here. First, they tend to have high melting points; a large amount of thermal energy must be supplied to the crystal lattice before the ions vibrate vigorously enough to overcome the electrostatic forces of attraction. Second, they do not conduct electricity in the solid state, because there are no free-moving ions to carry the current: contrastingly, they do conduct electricity in either the liquid state or when dissolved in water, because the ions are free to move. And third, they tend to be insoluble in organic solvents and soluble in water; the hydration energy released when ions are dissolved in water is often similar to or greater than the lattice energy.
Not surprisingly, these characteristic physical properties are used as evidence for ionic bonding: but there are at least three caveats. First, high melting points are also observed for covalent substances with giant structures [e.g., carbon-graphite and silicon(IV) oxide]. # Second, the measurement of electrical conductivity can be precluded because the compound either thermally decomposes before melting or has a low solubility in water [e.g., many carbonates and nitrates decompose before melting, and most oxides and sulfides are insoluble in water]. And third, a wide variety of covalent compounds are also soluble in water [e.g., many alcohols and carbohydrates with small molar masses].
* In point of fact, the exceptions to this rule of thumb are legion. However, such exceptions usually make no more than 'guest appearances' in an introductory course (... a rare case of 'ignorance is bliss' ?). # Typically, covalent substances with simple molecular structures have low melting points, because the attractive forces between molecules are weak, are insulators of electricity, because there are no free-moving ions, and are soluble in organic solvents such as trichloroethane.
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