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Chapter 11: Problem 46

Which ion in each of the following pairs would you expect to be more strongly hydrated? Why? a. \(\mathrm{Na}^{+}\) or \(\mathrm{Mg}^{2+}\) d. \(\mathrm{F}^{-}\) or \(\mathrm{Br}^{-}\) b. \(\mathrm{Mg}^{2+}\) or \(\mathrm{Be}^{2+}\) e. \(\mathrm{Cl}^{-}\) or \(\mathrm{ClO}_{4}^{-}\) c. \(\mathrm{Fe}^{2+}\) or \(\mathrm{Fe}^{3+}\) f. \(\mathrm{ClO}_{4}^{-}\) or \(\mathrm{SO}_{4}^{2-}\)

Short Answer

Expert verified
In summary, comparing the charge densities of each ion pair: a. Mg2+ is more strongly hydrated than Na+. b. Be2+ is more strongly hydrated than Mg2+. c. Fe3+ is more strongly hydrated than Fe2+. d. F- is more strongly hydrated than Br-. e. Cl- is more strongly hydrated than ClO4-. f. SO42- is more strongly hydrated than ClO4-.

Step by step solution

01

Pair a: Na+ or Mg2+

Na+ has a charge of +1 and a size of approximately 1.02 Å, while Mg2+ has a charge of +2 and a size of approximately 0.72 Å. Since Mg2+ has a higher charge (+2) and a smaller size (0.72 Å), its charge density is higher. Therefore, Mg2+ is expected to be more strongly hydrated than Na+.
02

Pair b: Mg2+ or Be2+

Both Mg2+ and Be2+ have a charge of +2. But their sizes are different – Mg2+ has a size of approximately 0.72 Å, while Be2+ has a size of approximately 0.45 Å. Since Be2+ is smaller than Mg2+ and has the same charge, Be2+ has a higher charge density. Therefore, Be2+ is expected to be more strongly hydrated than Mg2+.
03

Pair c: Fe2+ or Fe3+

For this pair, we have the same element but different charges. Fe2+ has a charge of +2 and a size of approximately 0.76 Å, while Fe3+ has a charge of +3 and a size of approximately 0.64 Å. Since Fe3+ has a higher charge and a smaller size, its charge density is higher. Therefore, Fe3+ is expected to be more strongly hydrated than Fe2+.
04

Pair d: F- or Br-

F- has a charge of -1 and a size of approximately 1.36 Å, while Br- has a charge of -1 and a size of approximately 1.96 Å. Since F- has a higher charge density (due to its smaller size) and the same charge, F- is expected to be more strongly hydrated than Br-.
05

Pair e: Cl- or ClO4-

Cl- has a charge of -1 and a size of approximately 1.81 Å, while ClO4- has a charge of -1 and a size of approximately 3.10 Å (considering the entire ion). Although both ions have the same charge, Cl- is smaller in size, leading to a higher charge density. Therefore, Cl- is expected to be more strongly hydrated than ClO4-.
06

Pair f: ClO4- or SO42-

For this pair, both ions have a charge of -2. ClO4- has a size of approximately 3.10 Å (considering the entire ion) while SO42- has a size of approximately 2.4 Å. Since SO42- has a smaller size and the same charge, its charge density is higher. Therefore, SO42- is expected to be more strongly hydrated than ClO4-.

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Most popular questions from this chapter

An aqueous solution is \(1.00 \% \mathrm{NaCl}\) by mass and has a density of \(1.071 \mathrm{~g} / \mathrm{cm}^{3}\) at \(25^{\circ} \mathrm{C}\). The observed osmotic pressure of this solution is \(7.83\) atm at \(25^{\circ} \mathrm{C}\).

A \(1.37-M\) solution of citric acid \(\left(\mathrm{H}_{3} \mathrm{C}_{6} \mathrm{H}_{5} \mathrm{O}_{7}\right)\) in water has a density of \(1.10 \mathrm{~g} / \mathrm{cm}^{3}\). Calculate the mass percent, molality, mole fraction, and normality of the citric acid. Citric acid has three acidic protons.

Calculate the solubility of \(\mathrm{O}_{2}\) in water a partial pressure of \(\mathrm{O}_{2}\) of 120 torr at \(25^{\circ} \mathrm{C}\). The Henry's law constant for \(\mathrm{O}_{2}\) is \(1.3 \times 10^{-3} \mathrm{~mol} / \mathrm{L} \cdot\) atm for Henry's law in the form \(C=k P\), where \(C\) is the gas concentration \((\mathrm{mol} / \mathrm{L})\).

The lattice energy" of NaI is \(-686 \mathrm{~kJ} / \mathrm{mol}\), and the enthalpy of hydration is \(-694 \mathrm{~kJ} / \mathrm{mol}\). Calculate the enthalpy of solution per mole of solid NaI. Describe the process to which this enthalpy change applies.

The high melting points of ionic solids indicate that a lot of energy must be supplied to separate the ions from one another. How is it possible that the ions can separate from one another when soluble ionic compounds are dissolved in water, often with essentially no temperature change?
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