Calculation ionic compounds

To date there is no evidence that sodium forms any chloride other than NaCl indeed the electronic theory of valency predicts that Na" and CU, with their noble gas configurations, are likely to be the most stable ionic species. However, since some noble gas atoms can lose electrons to form cations (p. 354) we cannot rely fully on this theory. We therefore need to examine the evidence provided by energetic data. Let us consider the formation of a number of possible ionic compounds and first, the formation of sodium dichloride , NaCl2. The energy diagram for the formation of this hypothetical compound follows the pattern of that for NaCl but an additional endothermic step is added for the second ionisation energy of sodium. The lattice energy is calculated on the assumption that the compound is ionic and that Na is comparable in size with Mg ". The data are summarised below (standard enthalpies in kJ) ... 

The Sodium Chloride and Cesium Chloride Structures.—The agreement found between the observed inter-atomic distances and our calculated ionic radii makes it probable that the crystals considered are built of only slightly deformed ions it should, then, be possible, with the aid of this conception, to explain the stability of one structure, that of sodium chloride, in the case of most compounds, and of the other, that of cesium chloride, in a few cases, namely, the cesium and thallous halides. 

Using Numbers Assume that the cation-anion ratio is 1 1 for KC1. All the solutions are the same concentration, which means that they all contain the same number of moles of ionic compound per liter of solution. Using this information and your results, calculate the cation to anion ratio for each of the known solutions. Record these ratios in Data Table 1. 

The shortest cation-anion distance in an ionic compound corresponds to the sum of the ionic radii. This distance can be determined experimentally. However, there is no straightforward way to obtain values for the radii themselves. Data taken from carefully performed X-ray diffraction experiments allow the calculation of the electron density in the crystal the point having the minimum electron density along the connection line between a cation and an adjacent anion can be taken as the contact point of the ions. As shown in the example of sodium fluoride in Fig. 6.1, the ions in the crystal show certain deviations from spherical shape, i.e. the electron shell is polarized. This indicates the presence of some degree of covalent bonding, which can be interpreted as a partial backflow of electron density from the anion to the cation. The electron density minimum therefore does not necessarily represent the ideal place for the limit between cation and anion. 

Ans. NaCl is an ionic compound it does not form molecules and so does not have a molecular weight. The formula weight of NaCl is 58.5 amu, calculated and used in exactly the same manner as a molecular weight for a molecular compound would be calculated and used. 

The oxidation number, or oxidation state, is the formal charge on an atom calculated on the basis that it is in a wholly ionic compound. Oxidation numbers are assigned according to several rules. 

Results have shown that the properties of solids can usually be modeled effectively if the interactions are expressed in terms of those between just pairs of atoms. The resulting potential expressions are termed pair potentials. The number and form of the pair potentials varies with the system chosen, and metals require a different set of potentials than semiconductors or molecules bound by van der Waals forces. To illustrate this consider the method employed with nominally ionic compounds, typically used to calculate the properties of perfect crystals and defect formation energies in these materials. 

The water that is trapped within the crystal structure of some ionic compounds (the water of hydration) can be removed easily by heating. The amount of the water present in a given sample can be determined by weighing the sample before and after this heating. The weight loss that occurs is the weight of the water in the sample. The percent of water in the hydrate is calculated as it is in a loss-on-drying experiment. 

Le Chatelier s principle is a powerful tool for explaining how a reaction at equilibrium shifts when a stress is placed on the system. In this experiment, you can use Le Chatelier s principle to evaluate the relative solubilities of two precipitates. By observing the formation of two precipitates in the same system, you can infer the relationship between the solubilities of the two ionic compounds and the numerical values of their solubility product constants (K ). You will be able to verify your own experimental results by calculating the molar solubilities of the two compounds using the Ksp for each compound. 

Calculate the molar solubilities of the two ionic compounds from their Ksp values. 

In this section, you learned why solutions of different salts have different pH values. You learned how to analyze the composition of a salt to predict whether the salt forms an acidic, basic, or neutral solution. Finally, you learned how to apply your understanding of the properties of salts to calculate the pH at the equivalence point of a titration. You used the pH to determine a suitable indicator for the titration. In section 9.2, you will further investigate the equilibria of solutions and learn how to predict the solubility of ionic compounds in solution. 

In Investigation 9-A, you will collect solubility data and use these data to determine a Ksp for calcium hydroxide, Ca(OH)2. When you calculate Ksp, you assume that the dissolved ionic compound exists as independent hydrated ions that do not affect one another. This assumption simplifies the investigation, but it is not entirely accurate. Ions do interfere with one another. As a result, the value of Ksp that you calculate will be just an approximation. iCp values that are calculated from data obtained from experiments such as Investgation 9-A are generally higher than the actual values. 

You calculate Qp by substituting the concentration of each ion into the expression. If Qp is larger than K p, the product of the concentrations of the ions is greater than it would be at equilibrium. For the system to attain equilibrium, some of the ions must leave the solution by precipitation. Conversely, if Qp is less than IQp, the product of the concentration of the ions is smaller than it is at equilibrium. Therefore, the solution is not yet saturated and more ions can be added to the solution without any precipitation. The relationship between Qsp and K p for the dissociation of a slightly soluble ionic compound is summarized on the next page. Use the following general equation as a reference. 

When finding the oxidation numbers of elements in ionic compounds, you can work with the ions separately. For example, Na2Cr207 contains two Na" ions, and so sodium has an oxidation number of+1. The oxidation numbers of Cr and 0 can then be calculated as shown in part (c) of the Sample Problem. 

D. Molecular Orbital Calculations. Pharmacological Activity of Meso-ionic Compounds Meso-ionic— Definition and Delineation... 

In the calculations of enthalpies of formation of ionic compounds, the differences from the accepted experimental values (which are very accurate) are probably due to the essential simplicity of the model, rather than having any other significance. If large discrepancies are found between experimental and calculated quantities, this probably means that the model is in error and that the compounds have a considerable covalent character. 

If activity coefficients arc iunorcd la-Mimcd lo bo unity a gro s approximation responsible lor the noii-quantitative connection between chances in Gibbs energy ol solution and actual salt solubilities), it is possible to draw up a table of values of the change in (iibbs energy for the solution of a compound in water that might be expected for various solubilities. Table. Tin contains the calculations of A , (/ for various solubilities of I I ionic compounds. The calculations arc based on the approximate relationship ... 

In principle, we can use the Born-Haber cycle to predict whether a particular ionic compound should be thermodynamically stable, on the basis of calculated values of U, and so proceed to explain all of the chemistry of ionic solids. The relevant quantity is actually the free energy of formation, AGf, and this is calculable if an entropy cycle is set up to complement the Born-Haber enthalpy cycle. However, in practice AHf dominates the energetics of formation of ionic compounds. 

Table VI gives the values of ionization energies which, together with the values of ionic radii and electron affinities, make it possible to calculate the energy changes which occur in the formation of ionic molecules, and to determine which ionic compounds of two elements...

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