Sodium chloride enthalpy

Let us consider the formation of sodium chloride from its elements. An energy (enthalpy) diagram (called a Born-Haber cycle) for the reaction of sodium and chlorine is given in Figure 3.7. (As in the energy diagram for the formation of hydrogen chloride, an upward arrow represents an endothermic process and a downward arrow an exothermic process.)... 

A/ij the lattice energy of sodium chloride this is the heat liberated when one mole of crystalline sodium chloride is formed from one mole of gaseous sodium ions and one mole of chloride ions, the enthalpy of formation of sodium chloride. 

Sodium Chlorite. The standard enthalpy, Gibbs free energy of formation, and standard entropy for aqueous chlorite ions ate AH° = —66.5 kJ/mol ( — 15.9 kcal/mol), AG = 17.2 kJ/mol (4.1 kcal/mol), and S° = 0.1883 kJ/(molK) (0.045 kcal/(molK)), respectively (107). The thermal decomposition products of NaClO, in the 175—200°C temperature range ate sodium chlorate and sodium chloride (102,109) ... 

In the second hypothetical step, we imagine the gaseous ions plunging into water and forming the final solution. The molar enthalpy of this step is called the enthalpy of hydration, AHhvd, of the compound (Table 8.7). Enthalpies of hydration are negative and comparable in value to the lattice enthalpies of the compounds. For sodium chloride, for instance, the enthalpy of hydration, the molar enthalpy change for the process... 

When we include the data, the limiting enthalpy of solution of sodium chloride, the enthalpy change for the process... 

Because the enthalpy of solution is positive, there is a net inflow of energy as heat when the solid dissolves (recall Fig. 8.23b). Sodium chloride therefore dissolves endothermically, but only to the extent of 3 kj-mol-1. As this example shows, the overall change in enthalpy depends on a very delicate balance between the lattice enthalpy and the enthalpy of hydration. 

The enthalpy of formation of a compound is a so-called thermodynamic state function, which means that the value depends only on the initial and final states of the system. When the formation of crystalline NaCl from the elements is considered, it is possible to consider the process as if it occurred in a series of steps that can be summarized in a thermochemical cycle known as a Born-Haber cycle. In this cycle, the overall heat change is the same regardless of the pathway that is followed between the initial and final states. Although the rate of a reaction depends on the pathway, the enthalpy change is a function of initial and final states only, not the pathway between them. The Born-Haber cycle for the formation of sodium chloride is shown as follows ... 

Some values for the enthalpy of formation of Schottky defects in alkali halides of formula MX that adopt the sodium chloride structure are given in Table 2.1. The experimental determination of these values (obtained mostly from diffusion or ionic conductivity data (Chapters 5 and 6) is not easy, and there is a large scatter of values in the literature. The most reliable data are for the easily purified alkali halides. Currently, values for defect formation energies are more often obtained from calculations (Section 2.10). 

Pure potassium bromide, KBr, which adopts the sodium chloride structure, has the fraction of empty cation sites due to Schottky defects, ncv/Nc, equal to 9.159xl0-21 at 20°C. (a) Estimate the enthalpy of formation of a Schottky defect, Ahs. (b) Calculate the number of anion vacancies per cubic meter of KBr at 730°C (just below the melting point of KBr). The unit cell of KBr is cubic with edge length a = 0.6600 nm and contains four formula units of KBr. 

The following table gives the values of the fraction of Schottky defects, S/N, in a crystal of NaBr, with the sodium chloride structure, as a function of temperature. Estimate the formation enthalpy of the defects. 

Worked Example 3.15 What is the lattice enthalpy A //(lattice) of sodium chloride at 25 °C ... 

Figure 3.8 Born-Haber cycle constructed to obtain the lattice enthalpy A//(E, lce) of sodium chloride. All arrows pointing up represent endothermic processes and arrows pointing down represent exothermic processes (the figure is not drawn to scale)...

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. 

Silvester, L. F. Pitzer, K. S. "Thermodynamics of Electrolytes. 8. High-Temperature Properties, Including Enthalpy and Heat Capacity with Application to Sodium Chloride" J. Phys. Chem., 1977, 81, 1822. 

As the enthalpy of the dissolved sodium chloride in its standard state according to Henry s law is that of the infinitely dilute solution, A// , for the reaction in Equation (20.74) is... 

Q Calculate the standard enthalpies of formation for the compounds (i) sodium chloride and (ii) potassium iodide. The interionic distances in the compounds are 282 and 353 pm, respectively. Compare your answers with the accepted experimental values for these quantities, which are -411 and -327.6 kJ mol1, respectively. 

FIGURE 8.23 The enthalpy of solution, AHSC, is the sum of the enthalpy change required to separate the molecules or ions of the solute (the lattice enthalpy, AH,) and the enthalpy change accompanying their hydration, AHhvd. The outcome is finely balanced (a) in some cases, it is exothermic (b) in others, it is endothermic. The figures in (b) refer to sodium chloride, in kilojoules per mole (not to scale). For gaseous solutes, the lattice enthalpy is 0 because the molecules are already widely separated. 

To understand the values in Table 8.6, we can think of dissolving as a two-step process (Fig. 8.23). In the first hypothetical step, we imagine the ions separating from the solid to form a gas of ions. The change in enthalpy accompanying this highly endothermic step is the lattice enthalpy, AHL, of the solid, which was introduced in Section 6.20 (see Table 6.3 for values). The lattice enthalpy of sodium chloride (787 kj-mol-1), for instance, is the molar enthalpy change for the process... 

Now we bring the two steps of the dissolving process together and calculate the energy for the overall change. As we see from Fig. 8.23, the limiting enthalpy of solution of sodium chloride, the enthalpy change for the process... 

L. F. Silvester and K. S. Pitzer, Thermodynamics of electrolytes. 8. High-temperature properties, including enthalpy and heat capacity, with application to sodium chloride , J. Phys. Chem., 81, 1822-1828 (1977). 

The anhydrous salt explodes on impact [1], and decomposes violently at 200°C [2]. The trihydrate is also percussion sensitive [3] though other sources suggest that in clean, grease- and oil-free equipment the anhydrous salt is shockproof [4]. A bottle of the purified anhydrous salt exploded then burnt on opening, (transition to sodium chloride and oxygen has an enthalpy of 1.2 kJ/g), presumably initiated by friction in... 

Given the enthalpy of formation of an ionic solid, an experimental lattice energy can be obtained by thermochemical analysis. For example, the formation of crystalline sodium chloride is broken down as follows ... 

The atomisation enthalpy of elemental sodium Afl%tom, the first ionisation energy of atomic sodium Iu the dissociation enthalpy D of gaseous chlorine, the electron attachment energy Ex of atomic chlorine and the enthalpy of formation A//)1 of crystalline sodium chloride can all be taken from standard tabulations of experimental data. An experimental lattice energy UL is thus given by ... 

This, of course, is always negative, and plays the same role in aqueous thermochemistry as the lattice energy does in the energetics of ionic solids. The hydration enthalpy cannot be measured directly, and many thermodynamicists frown upon this or any other single-ion quantity. For example, the enthalpy of solution of sodium chloride can be measured and subjected to the following analysis ... 

The standard change of enthalpy for the formation of sodium chloride in an infinitely dilute aqueous solution is thus given by the sum of the standard changes of enthalpy for the last two changes of state. Therefore,... 

The second step shown here combines the terms AHZ (for expanding the solvent) and AH3 (for solvent-solute interactions) and is called the enthalpy (heat) of hydration (AH d). This term represents the enthalpy change associated with the dispersal of a gaseous solute in water. Thus the standard enthalpy of solution for dissolving sodium chloride is the sum of AH and AH yd-... 

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