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Chapter 13: Problem 36

Explain why pressure substantially affects the solubility of \(\mathrm{O}_{2}\) in water but has little effect on the solubility of \(\mathrm{NaCl}\) in water.

Short Answer

Expert verified
In summary, pressure substantially affects the solubility of \(O_2\) in water because it's a gas, and according to Henry's Law, gas solubility is directly proportional to the applied pressure. On the other hand, the solubility of ionic compounds like \(NaCl\) is mainly influenced by temperature and solvent nature rather than pressure, as described by Le Chatelier's Principle. Therefore, pressure has little effect on the solubility of \(NaCl\) in water.

Step by step solution

01

Understanding solubility

Solubility is the maximum amount of a solute that can dissolve in a solvent at a given temperature and pressure. In this case, the solutes are oxygen and sodium chloride, while the solvent is water.
02

Effect of pressure on gas solubility

According to Henry's Law, the solubility of a gas in a liquid is directly proportional to the partial pressure of the gas above the liquid. When the pressure increases, more gas molecules are forced into the liquid, increasing its solubility. This is why the solubility of oxygen, a gas, in water is significantly affected by pressure. For example, when pressurizing a can of soda, the increased pressure forces more carbon dioxide gas to dissolve in the liquid, which gives the soda its characteristic fizz when opened, and the pressure is released.
03

Effect of pressure on solubility of ionic compounds

When it comes to ionic compounds, such as sodium chloride (NaCl), Le Chatelier's Principle provides insight into the effect of pressure on solubility. Le Chatelier's Principle states that if a stress or change is applied to a system in equilibrium, the system will adjust to counteract the applied stress or change. However, NaCl is an ionic compound, and the dissolution process involves breaking and forming ionic bonds. The solubility of ionic compounds is mainly affected by temperature and the nature of the solvent, while the pressure has minimal influence on solubility.
04

Dissolution of NaCl in water

When NaCl dissolves in water, the ionic bonds between its Na⁺ and Cl⁻ ions are broken, and new hydration shell bonds are formed between the ions and the water molecules. This process is relatively undisturbed by changing pressures because the reaction is not significantly pressure-dependent.
05

Comparing the effect of pressure on O₂ and NaCl solubility

Based on our understanding of solubility, the effect of pressure on solubility is significantly more pronounced in the case of gases like oxygen, as described by Henry's Law, than it is for ionic compounds like sodium chloride, as described by Le Chatelier's Principle. In conclusion, pressure substantially affects the solubility of O₂ in water due to its direct proportionality with gas solubility, as described by Henry's Law. In contrast, pressure has little effect on the solubility of NaCl in water, as ionic solubility is mainly influenced by temperature and solvent nature rather than pressure.

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Most popular questions from this chapter

By referring to Figure 13.18 , determine whether the addition of \(40.0 \mathrm{~g}\) of each of the following ionic solids to \(100 \mathrm{~g}\) of water at \(40^{\circ} \mathrm{C}\) will lead to a saturated solution: (a) \(\mathrm{NaNO}_{3},\) (b) KCl, (c) \(\mathrm{K}_{2} \mathrm{Cr}_{2} \mathrm{O}_{7}\) (d) \(\mathrm{Pb}\left(\mathrm{NO}_{3}\right)_{2}\)

How many grams of ethylene glycol \(\left(\mathrm{C}_{2} \mathrm{H}_{6} \mathrm{O}_{2}\right)\) must be added to \(1.00 \mathrm{~kg}\) of water to produce a solution that freezes at \(-5.00^{\circ} \mathrm{C} ?\)

(a) Calculate the mass percentage of \(\mathrm{Na}_{2} \mathrm{SO}_{4}\) in a solution containing \(10.6 \mathrm{~g} \mathrm{Na}_{2} \mathrm{SO}_{4}\) in \(483 \mathrm{~g}\) water. (b) An ore contains \(2.86 \mathrm{~g}\) of silver per ton of ore. What is the concentration of silver in ppm?

Calculate the molarity of the following aqueous solutions: (a) \(0.540 \mathrm{~g} \mathrm{Mg}\left(\mathrm{NO}_{3}\right)_{2}\) in \(250.0 \mathrm{~mL}\) of solution, (b) \(22.4 \mathrm{~g}\) \(\mathrm{LiClO}_{4} \cdot 3 \mathrm{H}_{2} \mathrm{O}\) in \(125 \mathrm{~mL}\) of solution, (c) \(25.0 \mathrm{~mL}\) of \(3.50 \mathrm{M}\) \(\mathrm{HNO}_{3}\) diluted to \(0.250 \mathrm{~L}\)

(a) Why does a \(0.10 m\) aqueous solution of NaCl have a higher boiling point than a \(0.10 \mathrm{~m}\) aqueous solution of \(\mathrm{C}_{6} \mathrm{H}_{12} \mathrm{O}_{6} ?\) (b) Calculate the boiling point of each solution. (c) The experimental boiling point of the \(\mathrm{NaCl}\) solution is lower than that calculated, assuming that \(\mathrm{NaCl}\) is completely dissociated in solution. Why is this the case?
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