Consider the ionic compounds KF, NaCl, NaBr, and LiCl. (b) Based on your answer to part (a), arrange these four compounds in order of decreasing lattice energy. (c) Check your predictions in part (b) with the experimental values of lattice energy from Table 8.1. Are the predictions from ionic radii correct?

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To arrange the ionic compounds in order of decreasing lattice energy, we need to consider the factors that affect lattice energy: the charges of the ions and the distance between them (ionic radii). Lattice energy is directly proportional to the product of the charges and inversely proportional to the distance between the ions.

Since all the compounds have ions with charges of +1 and -1, the charge factor is the same for all. Therefore, we focus on the ionic radii. Smaller ions will have higher lattice energies because they can get closer together, increasing the electrostatic attraction.

Compare the ionic radii of the cations: Li⁺, Na⁺, and K⁺. Li⁺ has the smallest radius, followed by Na⁺, and then K⁺. For the anions, F⁻ is smaller than Cl⁻ and Br⁻.

Based on ionic radii, predict the order of decreasing lattice energy: LiCl (smallest cation and anion combination), NaCl, NaBr, and KF (largest cation and anion combination).

To verify the predictions, compare them with experimental lattice energy values from Table 8.1. Check if the order based on ionic radii matches the experimental data, and note any discrepancies.

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Key Concepts

Here are the essential concepts you must grasp in order to answer the question correctly.

Lattice Energy

Lattice energy is the energy released when gaseous ions combine to form an ionic solid. It is a measure of the strength of the forces between the ions in an ionic compound. Higher lattice energy indicates stronger ionic bonds, which typically results from smaller ionic radii and higher charges on the ions. Understanding lattice energy is crucial for predicting the stability and solubility of ionic compounds.

Ionic Radii

Ionic radii refer to the effective size of an ion in a crystal lattice. The size of the ions affects the distance between them, which in turn influences the lattice energy. Generally, smaller ions lead to stronger attractions and higher lattice energies. Recognizing the trends in ionic radii, such as cation size decreasing across a period and anion size increasing down a group, is essential for comparing the lattice energies of different ionic compounds.

Trends in Lattice Energy

The trends in lattice energy can be predicted based on the charges and sizes of the ions involved. Compounds with ions of higher charges and smaller radii typically exhibit greater lattice energies. For example, comparing KF, NaCl, NaBr, and LiCl, one can analyze the ionic charges and sizes to predict their lattice energies. Understanding these trends allows for a systematic approach to arranging compounds by their lattice energies.

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Lattice Energy

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Textbook Question

(a) Does the lattice energy of an ionic solid increase or decrease (i) as the charges of the ions increase, (ii) as the sizes of the ions increase?

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(b) Arrange the following substances not listed in Table 8.1 according to their expected lattice energies, listing them from lowest lattice energy to the highest: MgS, KI, GaN, LiBr.

Textbook Question

Consider the ionic compounds KF, NaCl, NaBr, and LiCl. (a) Use ionic radii (Figure 7.8) to estimate the cation–anion distance for each compound.

Textbook Question

Which of the following trends in lattice energy is due to differences in ionic radii? (a) LiF > NaF > CsF, (b) CaO > KCl, (c) PbS > Li2O.

Textbook Question

Energy is required to remove two electrons from Ca to form Ca2+, and energy is required to add two electrons to O to form O2 - . Yet CaO is stable relative to the free elements. Which statement is the best explanation? (a) The lattice energy of CaO is large enough to overcome these processes. (b) CaO is a covalent compound, and these processes are irrelevant. (c) CaO has a higher molar mass than either Ca or O. (d) The enthalpy of formation of CaO is small. (e) CaO is stable to atmospheric conditions.

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Textbook Question

List the individual steps used in constructing a Born–Haber cycle for the formation of BaI₂ from the elements. Which of the steps would you expect to be exothermic?

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