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Chapter 11: Problem 5

You have read that adding a solute to a solvent can both increase the boiling point and decrease the freezing point. A friend of yours explains it to you like this: "The solute and solvent can be like salt in water. The salt gets in the way of freezing in that it blocks the water molecules from joining together. The salt acts like a strong bond holding the water molecules together so that it is harder to boil." What do you say to your friend?

Short Answer

Expert verified
In short, adding a solute to a solvent causes boiling point elevation and freezing point depression due to colligative properties. The solute particles interfere with the solvent particles' ability to escape into vapor, requiring more energy to boil the solution. Similarly, the solute disrupts the ordered arrangement of solvent particles during freezing, necessitating more energy to be removed for the solution to freeze. Your friend's explanation was partially correct but misunderstood the reason behind the boiling point elevation.

Step by step solution

01

Introduce Colligative Properties

Colligative properties are properties of solutions that depend on the ratio of the solute particles to the solvent particles. These properties include boiling point elevation and freezing point depression. It is important to note that these properties are independent of the nature of the solute, meaning that it doesn't matter if it's salt, sugar, etc.
02

Boiling Point Elevation

When a solute is added to the solvent, it increases the boiling point of the solution. This is because the solute particles interfere with the solvent particles' ability to escape into the vapor phase. Due to the solute particles presence, there are fewer solvent particles that can easily escape. Consequently, more energy (higher temperature) is needed for the solvent particles to overcome the attractive forces between them and become vapor. This increases the boiling point of the solution compared to the pure solvent.
03

Freezing Point Depression

Similarly, adding a solute to the solvent can decrease the freezing point of the solution. When a solute is added to a solvent, it causes disorder between the solvent particles. When the solution freezes, the solute particles get in the way of the solvent particles and disrupt the perfect arrangement of particles needed to form a solid. Therefore, more energy needs to be removed from the system (lower temperature) to overcome the disorder caused by the solute and allow the solvent particles to arrange themselves into a solid (frozen) state.
04

Address Your Friend's Explanation

Your friend's explanation is partially correct. It's true that the solute particles, like salt, get in the way of the water molecules forming a solid (freezing point depression); however, the part about the salt acting like a strong bond holding the water molecules together is not accurate. Instead, the salt (or other solutes) interferes with the water molecules ability to escape the solution, meaning that more energy is needed to boil the solution (boiling point elevation).

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Most popular questions from this chapter

Using the following information, identify the strong electrolyte whose general formula is $$ \mathrm{M}_{x}(\mathrm{~A})_{y} \cdot z \mathrm{H}_{2} \mathrm{O} $$ Ignore the effect of interionic attractions in the solution. a. \(\mathrm{A}^{n-}\) is a common oxyanion. When \(30.0 \mathrm{mg}\) of the anhydrous sodium salt containing this oxyanion \(\left(\mathrm{Na}_{n} \mathrm{~A}\right.\), where \(n=1,2\), or 3 ) is reduced, \(15.26 \mathrm{~mL}\) of \(0.02313 M\) reducing agent is required to react completely with the \(\mathrm{Na}_{n}\) A present. Assume a \(1: 1\) mole ratio in the reaction. b. The cation is derived from a silvery white metal that is relatively expensive. The metal itself crystallizes in a body-centered cubic unit cell and has an atomic radius of \(198.4 \mathrm{pm}\). The solid, pure metal has a density of \(5.243 \mathrm{~g} / \mathrm{cm}^{3}\). The oxidation number of \(\mathrm{M}\) in the strong electrolyte in question is \(+3\). c. When \(33.45 \mathrm{mg}\) of the compound is present (dissolved) in \(10.0 \mathrm{~mL}\) of aqueous solution at \(25^{\circ} \mathrm{C}\), the solution has an osmotic pressure of 558 torr.

Anthraquinone contains only carbon, hydrogen, and oxygen. When \(4.80 \mathrm{mg}\) anthraquinone is burned, \(14.2 \mathrm{mg} \mathrm{CO}_{2}\) and \(1.65 \mathrm{mg} \mathrm{H}_{2} \mathrm{O}\) are produced. The freezing point of camphor is lowered by \(22.3^{\circ} \mathrm{C}\) when \(1.32 \mathrm{~g}\) anthraquinone is dissolved in \(11.4 \mathrm{~g}\) camphor. Determine the empirical and molecular formulas of anthraquinone.

Calculate the molarity and mole fraction of acetone in a \(1.00 \mathrm{~m}\) solution of acetone \(\left(\mathrm{CH}_{3} \mathrm{COCH}_{3}\right)\) in ethanol \(\left(\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{OH}\right) .\) (Density of acetone \(=0.788 \mathrm{~g} / \mathrm{cm}^{3} ;\) density of ethanol \(\left.=0.789 \mathrm{~g} / \mathrm{cm}^{3} .\right)\) Assume that the volumes of acetone and ethanol add.

Erythrocytes are red blood cells containing hemoglobin. In a saline solution they shrivel when the salt concentration is high and swell when the salt concentration is low. In a \(25^{\circ} \mathrm{C}\) aqueous solution of \(\mathrm{NaCl}\), whose freezing point is \(-0.406^{\circ} \mathrm{C}\), erythrocytes neither swell nor shrink. If we want to calculate the osmotic pressure of the solution inside the erythrocytes under these conditions, what do we need to assume? Why? Estimate how good (or poor) of an assumption this is. Make this assumption and calculate the osmotic pressure of the solution inside the erythrocytes.

In flushing and cleaning columns used in liquid chromatography to remove adsorbed contaminants, a series of solvents is used. Hexane \(\left(\mathrm{C}_{6} \mathrm{H}_{14}\right)\), chloroform \(\left(\mathrm{CHCl}_{3}\right)\), methanol \(\left(\mathrm{CH}_{3} \mathrm{OH}\right)\), and water are passed through the column in that order. Rationalize the order in terms of intermolecular forces and the mutual solu- bility (miscibility) of the solvents.
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