Ionic solids compounds Salts

Solid Compounds. The tripositive actinide ions resemble tripositive lanthanide ions in their precipitation reactions (13,14,17,20,22). Tetrapositive actinide ions are similar in this respect to Ce . Thus the duorides and oxalates are insoluble in acid solution, and the nitrates, sulfates, perchlorates, and sulfides are all soluble. The tetrapositive actinide ions form insoluble iodates and various substituted arsenates even in rather strongly acid solution. The MO2 actinide ions can be precipitated as the potassium salt from strong carbonate solutions. In solutions containing a high concentration of sodium and acetate ions, the actinide ions form the insoluble crystalline salt NaM02(02CCH2)3. The hydroxides of all four ionic types are insoluble ... 

Boddington and Iqbal [727] have interpreted kinetic data for the slow thermal and photochemical decompositions of Hg, Ag, Na and T1 fulminates with due regard for the physical data available. The reactions are complex some rate studies were complicated by self-heating and the kinetic behaviour of the Na and T1 salts is not described in detail. It was concluded that electron transfer was involved in the decomposition of the ionic solids (i.e. Na+ and Tl+ salts), whereas the rate-controlling process during breakdown of the more covalent compounds (Hg and Ag salts) was probably bond rupture. 

The ionic conductivities of most solid crystalline salts and oxides are extremely low (an exception are the solid electrolytes, which are discussed in Section 8.4). The ions are rigidly held in the crystal lattices of these compounds and cannot move under the effect of applied electric fields. When melting, the ionic crystals break down, forming free ions the conductivities rise drastically and discontinuously, in some cases up to values of over 100 S/m (i.e., values higher than those of the most highly conducting electrolyte solutions). 

A solute may be present as ions or as molecules. We can identify the form of the solute by noting whether the solution conducts an electric current. Because a current is a flow of electric charge, only solutions that contain ions conduct electricity. There is such a tiny concentration of ions in pure water (about 10-7 m) that water alone does not conduct electricity. A substance that dissolves to give a solution that conducts electricity is called an electrolyte. Electrolyte solutions (solutions of electrolytes), which conduct electricity because they contain ions, include aqueous solutions of ionic compounds, such as sodium chloride and potassium nitrate. The ions are not formed when an ionic solid dissolves they exist as separate ions in the solid but become free to move apart in the presence of water (Fig. 1.1). Acids also are electrolytes. Unlike salts, they are molecular compounds in the pure state but form ions when they dissolve. One example is hydrogen chloride, which exists as gaseous HC1 molecules. In solution, however, HCl is called hydrochloric acid and is present as hydrogen ions and chloride ions. 

Although most compounds prepared from neutral parent molecules are neutral DA complexes, in some cases an ionic charge-transfer salt results. In addition, there are compounds prepared from D+p and A-q subunits which have been found to be ionic. Many of the ionic crystals involve A q units that are variants on metal bis-ethylene-1,2-dithiolenes, namely [M(tfd)2] and [M(mnt)2]-2. A variety of planar organic cations have been used to prepare such solids TTF+ M5), POZ+ 56 57), PTZ+ 56 57), and NMP+ 59). A... 

Salts of the bases MOH are crystalline, ionic solids, colorless except where the anion is colored. For the alkali metal ions the energies required to excite electrons to the lowest available empty orbitals could be supplied only by quanta far out in the vacuum ultraviolet (the transition 5p6 —5p56s in Cs+ occurs at 1000 A). However, colored crystals of compounds such as NaCl are sometimes encountered. Color arises from the presence in the lattice of holes and free electrons, called color centers, and such chromophoric disturbances can be produced by irradiation of the crystals with X rays and nuclear radiation. The color results from transitions of the electrons between energy levels in the holes in which they are trapped. These electrons behave in principle similarly to those in solvent cages in the liquid ammonia solutions, but the energy levels are differently spaced and consequently the colors are different and variable. Small excesses of metal atoms produce similar effects, since these atoms form M+ ions and electrons that occupy holes where anions would be in a perfect crystal. 

Stability constants are calculated from the concentrations of the species present in equilibrium mixtures containing the metal ion and the ligand in a wide range of proportions. Activity coefficients are kept constant by appropriate additions of a salt, usually sodium perchlorate, whose ions do not compete with those of the cation and ligand. Concentrations at different ionic strengths are extrapolated to zero ionic strength. It may be necessary to find the number of water molecules displaced at each step the total of these is not necessarily the same as the co-ordination number of the cation in the solid compound. Particularly in a polar solvent such as water, the ligands may not displace all the solvent molecules. 

To conduct an electric current, a substance must satisfy two conditions. First, the substance must contain charged particles. Second, those particles must be free to move. Because ionic compounds are composed of charged particles, you might expect that they could be good conductors. While particles in a solid have some vibrational motion, they remain in fixed locations, as shown by the model in Figure 11a. Therefore, ionic solids, such as salts, generally are not conductors of electric current because the ions cannot move. 

Until now, we have not made a distinction between the dissolving process of a covalent solid, such as sugar, and that of an ionic solid, such as table salt. Surface area and temperature affect both covalent and ionic solids. However, the dissolving of an ionic compound involves a unique factor the separation of ions from the lattice into individual dissolved ions. This process, called dissociation, can be represented as an equation. 

These properties should be treated with caution and are not definitive for ionic solids in general, since some ionic compounds do not show all these properties. For example, ammonium salts such as ammonium nitrate have low melting points ca. 150-300 °C), and some compounds containing doubly charged ions such as magnesium oxide, MgO, have very low solubility. 

Evaporate the water from the above reaction mixture, and the ionic solid barium chloride remains. An ionic compound that results from the reaction of an acid and a base is called a salt. Thus, in a typical aqueous neutralization reaction, the reactants are an acid and a base, and the products are a salt solution and water ... 

To what reaction does the solubility product constant, refer Table 16.1 lists sp values for several ionic solids. For any of these ionic compounds, you should be able to calculate the solubility. What is the solubility of a salt, and what procedures do you follow to calculate the solubility of a salt How would you calculate the value for a salt given the solubility ... 

The common ion effect for ionic solids (salts) is to significantly decrease the solubility of the ionic compound in water. Explain the common ion effect. 

Table salt, an ionic solid, and table sugar, a covalent solid, are similar in appearance. However, these compounds behave differently when heated. Salt does not melt, but sugar melts at a relatively low temperature. Does the type of bonding in a compound affect its properties ... 

In this experiment you will start with two ionic solids, nickel(II) nitrate hexahydrate, Ni(NO3)2 6H2O, and potassium chloride, KCl. You will dissolve these solids in water and use either a temperature change or a concentration change by evaporation to crystallize relatively soluble salts from the aqueous solution. The solubilities at different temperatures of the four possible compounds involved are given in TABLE 19.1. 

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