Ch.8 - Basic Concepts of Chemical Bonding

Brown - Chemistry: The Central Science 15th Edition

Brown15th EditionChemistry: The Central ScienceISBN: 9780137542970Not the one you use?Change textbook

Brown 15th Edition

Ch.8 - Basic Concepts of Chemical Bonding

Problem 26a

Chapter 8, Problem 26a

(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?

Verified step by step guidance

Understand that lattice energy is the energy required to separate one mole of an ionic solid into its gaseous ions.

Recall that lattice energy is directly proportional to the product of the charges of the ions. Therefore, as the charges of the ions increase, the lattice energy increases.

Remember that lattice energy is inversely proportional to the distance between the ions, which is related to the size of the ions. As the sizes of the ions increase, the distance between them increases, leading to a decrease in lattice energy.

Consider the formula for lattice energy: \( U = \frac{k \cdot Q_1 \cdot Q_2}{r} \), where \( U \) is the lattice energy, \( k \) is a constant, \( Q_1 \) and \( Q_2 \) are the charges of the ions, and \( r \) is the distance between the centers of the ions.

Summarize: (i) Lattice energy increases with increasing ion charges, and (ii) lattice energy decreases with increasing ion sizes.

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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 a more stable ionic solid, as it reflects stronger ionic bonds resulting from the electrostatic attraction between oppositely charged ions.

Charge of Ions

The charge of ions significantly affects lattice energy. As the charges of the ions increase, the electrostatic attraction between them becomes stronger, leading to a higher lattice energy. For example, a compound formed from divalent ions (e.g., Mg²⁺ and O²⁻) will have a greater lattice energy than one formed from monovalent ions (e.g., Na⁺ and Cl⁻) due to the increased charge.

Size of Ions

The size of the ions also influences lattice energy. As the size of the ions increases, the distance between the centers of the ions increases, which weakens the electrostatic attraction and results in lower lattice energy. Therefore, larger ions lead to a decrease in lattice energy, as seen when comparing NaCl to KCl, where K⁺ is larger than Na⁺.

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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

NaCl and KF have the same crystal structure. The only difference between the two is the distance that separates cations and anions. (a) The lattice energies of NaCl and KF are given in Table 8.1. Based on the lattice energies, would you expect the Na─Cl or the K─F distance to be longer?

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Textbook Question

Which of the following trends in lattice energy is due to differences in ionic radii: a. NaCl > RbBr > CsBr, b. BaO > KF, c. SrO > SrCl2?

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Textbook Question

NaCl and KF have the same crystal structure. The only difference between the two is the distance that separates cations and anions. (b) Use the ionic radii given in Figure 7.8 to estimate the Na─Cl and K─F distances.

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Textbook Question

The substances NaF and CaO are isoelectronic (have the same number of valence electrons). (b) What are the charges of each of the anions in each compound?

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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.

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