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Chapter 13: Problem 99

How does each of the following affect the solubility of an ionic compound: (a) lattice energy, (b) solvent (polar versus nonpolar), (c) enthalpies of hydration of cation and anion?

Short Answer

Expert verified
The solubility of an ionic compound is directly influenced by Lattice energy, Solvent type, and Enthalpies of hydration. Higher lattice energy tends to make compounds less soluble. Polar solvents enhance solubility of ionic compounds while non-polar solvents hinder it. Greater enthalpies of cation and anion hydration boost solubility as the process releases energy which can overcome the lattice energy.

Step by step solution

01

Understanding the effect of Lattice Energy

Lattice energy refers to the energy released when constituent ions of an ionic compound come together to form the lattice structure. Compounds with high lattice energy are more difficult to break apart, which makes them generally less soluble in a solvent. Thus, the greater the lattice energy, the less soluble the compound.
02

Understanding the effect of Solvent (Polar versus nonpolar)

The type of solvent significantly affects the solubility of ionic compounds. Generally, ionic compounds are more soluble in polar solvents (like water) because polar solvents can dissolve the ions in the ionic compound due to the charge-charge interaction. In the reverse, ionic compounds are usually less soluble in non-polar solvents (like oil) because non-polar solvents are not capable of dissolving ions effectively.
03

Understanding the effect of Enthalpies of Hydration of Cation and Anion

Enthalpy of hydration refers to the energy change when one mole of gaseous ions is completely hydrated. Greater enthalpies of hydration (both cations and anions) tend to increase the solubility of an ionic compound. When ions are surrounded by water molecules (a process called hydration), energy is released. If this energy release is significant, it can help to overcome the lattice energy, allowing the ionic compound to dissolve.

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Most popular questions from this chapter

The molar mass of benzoic acid \(\left(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{COOH}\right)\) determined by measuring the freezing-point depression in benzene is twice that expected for the molecular formula, \(\mathrm{C}_{7} \mathrm{H}_{6} \mathrm{O}_{2}\). Explain this apparent anomaly.

A mixture of liquids \(A\) and \(B\) exhibits ideal behavior. At \(84^{\circ} \mathrm{C},\) the total vapor pressure of a solution containing 1.2 moles of \(A\) and 2,3 moles of \(B\) is 331 \(\mathrm{mmHg}\). Upon the addition of another mole of \(\mathrm{B}\) to the solution, the vapor pressure increases to 347 \(\mathrm{mmHg}\). Calculate the vapor pressures of pure \(\mathrm{A}\) and \(\mathrm{B}\) at \(84^{\circ} \mathrm{C}\)

Calculate the molality of each of the following aqueous solutions: (a) \(2.50 \mathrm{M} \mathrm{NaCl}\) solution (density of solution \(=1.08 \mathrm{~g} / \mathrm{mL}\) ), (b) 48.2 percent by mass KBr solution.

Using Henry's law and the ideal gas equation to prove the statement that the volume of a gas that dissolves in a given amount of solvent is independent of the pressure of the gas. (Hint: Henry's law can be modified as \(n=k P\), where \(n\) is the number of moles of the gas dissolved in the solvent.)

Explain why molality is used for boiling-point elevation and freezing-point depression calculations and molarity is used in osmotic pressure calculations.
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