![]() Note that r 0 may differ between the gas-phase dimer and the lattice. In such an arrangement each cation in the lattice is surrounded by more than one anion (typically four, six, or eight) and vice versa, so it is more stable than a system consisting of separate pairs of ions, in which there is only one cation–anion interaction in each pair. ![]() While Equation 4.1.1 has demonstrated that the formation of ion pairs from isolated ions releases large amounts of energy, even more energy is released when these ion pairs condense to form an ordered three-dimensional array. These properties result from the regular arrangement of the ions in the crystalline lattice and from the strong electrostatic attractive forces between ions with opposite charges. They are not easily deformed, and they melt at relatively high temperatures. Ionic compounds are usually rigid, brittle, crystalline substances with flat surfaces that intersect at characteristic angles. Metal ores are commonly combinations of metal atoms with oxygen atoms, and this combination is produced when metals rust, so the process where electrons are transferred to the oxygen atoms from the metal atoms is known as oxidation of the metal and the reverse process, where pure metals are produced is called reduction of the ore to the metal. The reaction of a metal with a nonmetal usually produces an ionic compound that is, electrons are transferred from the metal to the nonmetal. To understand the relationship between the lattice energy and physical properties of an ionic compound.Dangerous spattering of strong acid or base can be avoided if the concentrated acid or base is slowly added to water, so that the heat liberated is largely dissipated by the water.\( \newcommand\) If water is added to a concentrated solution of sulfuric acid (which is 98% H 2SO 4 and 2% H 2O) or sodium hydroxide, the heat released by the large negative Δ H can cause the solution to boil. This phenomenon is particularly relevant for strong acids and bases, which are often sold or stored as concentrated aqueous solutions. If the initial dissolution process is exothermic (Δ H < 0), then the dilution process is also exothermic. The Δ H soln values given previously and in Table 8.2.2 for example, were obtained by measuring the enthalpy changes at various concentrations and extrapolating the data to infinite dilution.īecause Δ H soln depends on the concentration of the solute, diluting a solution can produce a change in enthalpy. The amount of heat released or absorbed when a substance is dissolved is not a constant it depends on the final concentration of the solute. Because of the mass of white sodium acetate that has crystallized, the metal disc is no longer visible. After the hot pack has been agitated, the sodium acetate crystallizes (right) to release heat. ![]() Because the sodium acetate is in solution, you can see the metal disc inside the pack. When the pack is agitated, sodium acetate trihydrate crystallizes, and heat is evolved:įigure 9.5.1 An Instant Hot Pack Based on the Crystallization of Sodium Acetate The hot pack is at room temperature prior to agitation (left). With cooling, an unstable supersaturated solution containing excess solute is formed. ![]() At high temperatures, sodium acetate forms a highly concentrated aqueous solution. If the salt is CaCl 2, heat is released to produce a solution with a temperature of about 90☌ hence the product is an “instant hot compress.” If the salt is NH 4NO 3, heat is absorbed when it dissolves, and the temperature drops to about 0° for an “instant cold pack.”Ī similar product based on the precipitation of sodium acetate, not its dissolution, is marketed as a reusable hand warmer ( Figure 9.5.1). When the pack is twisted or struck sharply, the inner plastic bag of water ruptures, and the salt dissolves in the water. Both types consist of a plastic bag that contains about 100 mL of water plus a dry chemical (40 g of CaCl 2 or 30 g of NH 4NO 3) in a separate plastic pouch. Single-use versions of these products are based on the dissolution of either calcium chloride (CaCl 2, Δ H soln = −81.3 kJ/mol) or ammonium nitrate (NH 4NO 3, Δ H soln = +25.7 kJ/mol). Substances with large positive or negative enthalpies of solution have commercial applications as instant cold or hot packs. Table 9.5.1 Enthalpies of Solution at 25☌ of Selected Ionic Compounds in Water (in kJ/mol)
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