The ionic bond

As you saw on the last page, a sodium atom can lose one electron, and a chlorine atom can gain one, to obtain full outer shells. So, when a sodium atom and a chlorine atom react together, the sodium atom loses its electron to the chlorine atom, and two ions are formed. Here, sodium electrons are shown as and chlorine electrons as x, but remember that all electrons are exactly the same  

The two ions have opposite charges, so they attract each other. The force of attraction between them is strong. It is called an ionic bond, or sometimes an electrovalent bond. 

When sodium reacts with chlorine, billions of sodium and chloride ions form, and are attracted to each other. But the ions do not stay in pairs. 

Instead, they duster together, so that each ion is surrounded by six ions of opposite charge. 

They are held together by strong ionic bonds. (Each ion forms six bonds.) 

The Ionic Bond Although there is no sharp boundary between ionic bonding and covalent bonding, 

Properties of Ionic Several properties distinguish ionic compounds from covalent compounds. These 

Substances may be related rather simply to the crystal structure of ionic compounds, namely, 

Ionic compounds tend to have very low electrical conductivities as solids but conduct electricity quite well when molten. This conductivity is attributed to the presence of ions, atoms charged either positively or negatively, which are free to move under the influence of an electric field. In the solid, the ions are 

1 Some very interesting ionic compounds prove to be exceptions to these rules. They are discussed in Chapter 7. 

The ionic bond is the bond between charged atoms or groups of atoms (complex ions) and is the only one of the four main types of bond that can be satisfactorily described in classical (non-wave-mechanical) terms. Monatomic ions formed by elements of the earlier A subgroups and ions such as N , 0 , etc., F etc., have noble gas configurations, but many transition-metal ions and ions containing two s electrons (such as Tl and Pb ) have less symmetrical structures. We shall not be concerned here with the numerous less stable ionic species formed in the gaseous phase. 

The simplest systems containing ionic bonds are the gaseous molecules of alkali halides and oxides, the structures of which are noted in Chapters 9 and 12 we refer later to the halide molecules in connection with polarization. The importance of the ionic bond lies in the fact that it is responsible for the existence at ordinary temperatures, as stable solids, of numerous metallic oxides and halides (both simple and complex), of some sulphides and nitrides, and also of the very numerous crystalline compounds containing complex ions, particularly oxy-ions, which may be finite (CO3 , NO3, SOl , etc.) or infinite in one, two, or three dimensions. 

The adequacy of a purely electrostatic picture of simple ionic crystals A, X is demonstrated by the agreement between the values of the lattice energy resulting from direct calculation, from the Born-Haber cycle, and in a few cases from direct measurement. 

Recall that ionization ener is the amount of energy required to remove an electron and form a cation. Electron affnity is the amount of energy required to add an electron and form an anion. See Section 6.7, pages 233-237. 

For convenience, imagine that this reaction occurs in separate steps—first the ionization of Li  

Lithium fluoride. Industrially, UF (like most other ionic compounds) is obtairted by purifying minerals containing the compound. 

imagine the two separate ions joining to form a LiF unit  

Many other common reactions lead to the formation of ionic bonds. For instance, calcium bums in oxygen to form calcium oxide  

Assuming that the diatomic O2 molecule first splits into separate oxygen atoms (we will look at the energetics of this step later), we can represent the reaction with Lewis symbols  

Although there is no sharp boundary between ionu bonding and covalent bonding, it is convenient to consider each of these as a separate entity before attempting to discuss molecules and lattices, in which both are important Furthermore, because the purely ionic bond may be described with a simple electrostatic model, it is advantageous to discuss it first The simplicity of the electrostatic.model has caused chemists to think of many solids as systems of ions. We shall see that this view needs some modification, and there are, of course, many solids, ranging from diamond to metals, which require alternative theories of bonding. 

Several properties distinguish ionic compounds from covalent compounds. These may be related rather simply to the crystal structure of ionic compounds, namely, a lattice composed of positive and negative ions in such a way that the attractive forces between oppositely charged ions are maximized and the repulsive forces between ions of the same charge ate minimized. Before discussing some of the possible geometries, a few simple properties of ionic compounds may be mentioned  

Ionic compounds are often soluble in polar solvents with high permittivities (dielectric constants). The energy of interaction of two charged particles is given by 

Another way of looking at this phenomenon is to consider the interaction between the dipole moments of the polar solvent and the ions. Such solvation will provide considerable energy to offset the otherwise unfavorable energetics of breaking up the crystal lattice (see Chapter 8). 

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The semicrystalline, ethylene-based ionomers of commerce are flexible, transparent polymers notable for high strength and elasticity in both soUd and molten states. The ionic bonding is completely reversible (8) and has a strong influence on properties, even at temperatures well above the melting point. 

The ionic bond is the most obvious sort of electrostatic attraction between positive and negative charges. It is typified by cohesion in sodium chloride. Other alkali halides (such as lithium fluoride), oxides (magnesia, alumina) and components of cement (hydrated carbonates and oxides) are wholly or partly held together by ionic bonds. 

So far, we have had to do work to create the ions which will make the ionic bond it does not seem to be a very good start. However, the + and - charges attract each other and if we now bring them together, the force of attraction does work. This force is simply that between two opposite point charges ... 

How much can we bend this bond Well, the electrons of each ion occupy complicated three-dimensional regions (or orbitals ) around the nuclei. But at an approximate level we can assume the ions to be spherical, and there is then considerable freedom in the way we pack the ions round each other. The ionic bond therefore lacks directionality, although in packing ions of opposite sign, it is obviously necessary to make sure that the total charge (+ and -) adds up to zero, and that positive ions (which repel each other) are always separated by negative ions. 

Sq can be calculated from the theoretically derived U(r) curves of the sort described in Chapter 4. This is the realm of the solid-state physicist and quantum chemist, but we shall consider one example the ionic bond, for which U(r) is given in eqn. (4.3). Differentiating once with respect to r gives the force between the atoms, which must, of course, be zero at r = rg (because the material would not otherwise be in equilibrium, but would move). This gives the value of the constant B in equation (4.3) ... 

The stable sodium ion has a positive charge because it is short of one electron and the chlorine atom is negatively charged for the converse reason. Ionic bonds are seldom found in polymers of current interest as plastics materials although the ionic bond is important in ion-exchange resins and in the ionomers (see Chapter 11). 

From Coulombs law, the strength of the ionic bond should depend on two factors ... 

The size of the ions. The ionic bond in NaCl (mp = 801°C) is somewhat stronger than that in KBr (mp = 734°C) because the intemuclear distance is smaller in NaCl ... 

The formulated principals correlating crystal structure features with the X Nb(Ta) ratio do not take into account the impact of the second cation. Nevertheless, substitution of a second cation in compounds of similar types can change the character of the bonds within complex ions. Specifically, the decrease in the ionic radius of the second (outer-sphere) cation leads not only to a decrease in its coordination number but also to a decrease in the ionic bond component of the complex [277]. 

Polyelectrolyte complexes composed of various weight ratios of chitosan and hyaluronic acid were found to swell rapidly, reaching equilibrium within 30 min, and exhibited relatively high swelling ratios of 250-325% at room temperature. The swelling ratio increased when the pH of the buffer was below pH 6, as a result of the dissociation of the ionic bonds, and with increments of temperature. Therefore, the swelling ratios of the films were pH-and temperature-dependent. The amount of free water in the complex films increased with increasing chitosan content up to 64% free water, with an additional bound-water content of over 12% [29]. 

Taber, K. S. (1994). Misunderstanding the ionic bond. Education in Chemistry, 7, 100-102. 

Justi, R., Mendon9a, P. C. C. (2007). Modelling in order to learn an important sub-micro representation the nature of the ionic bond. Paper presented at the VI Conferenee of the European Science Education Research Association, Malmo, Sweden, 21-25 August. 

When an ionic compound dissolves in water, energy is needed to break the ionic bonds of the crystal. As the ions attach to the water molecules and become hydrated, energy is released. The process is endothermic if the energy needed to break the bonds is greater than the energy released when the ions attach to water. 

The method of catalyst immobilisation appeared to affect its performance in catalysis. Catalyst obtained by method II showed a low selectivity in the hydroformylation of 1-octene (l b aldehyde ratio was even lower than 2) at a very high rate and high yields of isomerised alkenes (Table 3.2, entry 2), whereas procedure IV resulted in a catalyst that was highly selective for the linear aldehyde (with a l b ratio of 37) (entry 5). In accordance with examples from literature it is likely that procedure II gave rise to the ionic bonding of ligand-free rhodium cations on the slightly acidic silica surface [29],... 

Ionic compounds, as compared to covalent compounds, tend to have greater densities, higher melting and boiling points, and can be soluble in the very polar solvent, water, if the ionic bond is not too strong. 

The following question remains Why are the sigma-type interactions weaker than the pi-type interactions in this case This is apparently a geometrical effect, enforced by the relative sizes of M and X valence and core orbitals in the ionic bonding environment, and illustrated in the NAO diagrams of Fig. 2.18. We may... 

It is quite difficult to measure an accurate enthalpy of solution A//( olutioni with a calorimeter, but we can measure it indirectly. Consider the example of sodium chloride, NaCl. The ions in solid NaCl are held together in a tight array by strong ionic bonds. While dissolving in water, the ionic bonds holding the constituent ions of Na+ and Cl- in place break, and new bonds form between the ions and molecules of water to yield hydrated species. Most simple ions are surrounded with six water molecules, like the [Na(H20)6]+ ion (VI). Exceptions include the proton with four water molecules (see p. 235) and lanthanide ions with eight. 

Energy is needed to break the ionic bonds in the solid salt and energy is liberated forming hydration complexes like VI. We also break some of the natural hydrogen bonds in the water. The overall change in enthalpy is termed the enthalpy of solution, A// olutioni. Typical values are —207 kJmol-1 for nitric acid 34 kJmol-1 for potassium nitrate and —65.5 kJmol-1 for silver chloride. 

Potassium is a metal, and the polyatomic anion, C104 is a nonmetal therefore, the compound is an ionic solid at room temperature. When the compound is dissolved in water, the ionic bond between the cation, K+, and the polyatomic anion, Cl()4, is broken due to the polarity of the water molecule, resulting in the two aqueous ions, K+ and C104 . 

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