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Chapter 11: Problem 29

Adding a solute to a solvent extends the liquid phase over a larger temperature range. Explain this statement.

Short Answer

Expert verified
When a solute is added to a solvent, it disrupts the natural arrangement of the solvent molecules and forms intermolecular interactions. This leads to a depression of the freezing point and an elevation of the boiling point of the solution. Consequently, the temperature range over which the solution exists as a liquid is extended. This effect of the solute on the phase diagram is what the statement "adding a solute to a solvent extends the liquid phase over a larger temperature range" refers to.

Step by step solution

01

Understanding Phase Changes

In order to explain this statement, it is important to understand the concept of phase changes. A substance, such as a solvent, exists in different phases (solid, liquid, and gas) depending on the temperature. The phase change occurs when a substance changes from one phase to another, such as from solid to liquid or liquid to gas. The temperature at which a phase change occurs is called the phase change temperature.
02

Solute-Solvent Interactions

When a solute is added to a solvent, it disrupts the natural arrangement of the solvent molecules. This disruption occurs because the solute forms intermolecular interactions, such as hydrogen bonds, dipole-dipole interactions, and van der Waals forces, with the solvent molecules. These interactions will affect the energy required for the solvent molecules to escape or overcome the forces holding them together.
03

Effect on Melting Point and Freezing Point

The presence of a solute typically leads to a depression of the freezing point (melting point) and an elevation of the boiling point. This occurs because the solute molecules disrupt the normal arrangement of the solvent molecules. Consequently, it requires more energy to overcome these interactions and move the molecules further apart in the phase change from solid to liquid (freezing point depression) and less energy for the liquid to solid phase change (melting point elevation).
04

Effect on Boiling Point

Similarly, the presence of a solute in a solvent can raise the boiling point of the solution. This is because the solute-solvent interactions force the solvent molecules to use more energy to overcome these additional intermolecular forces and move further apart in the phase change from liquid to gas. As a result, the boiling point of the solution will be higher than that of the pure solvent.
05

Extending the Liquid Phase Temperature Range

Now, recalling the concept of liquid phase extension: when a solute is added to a solvent, the freezing point is depressed, and the boiling point is elevated. Consequently, the temperature range over which the solution exists as a liquid is extended. So, the statement "adding a solute to a solvent extends the liquid phase over a larger temperature range" refers to this effect of the solute on the phase diagram and the subsequent broadening of the liquid phase temperature range.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Freezing Point Depression
When a solute is added to a solvent, it causes a phenomenon known as freezing point depression. This means that the temperature at which the liquid turns into a solid is lowered. How does this happen?
The solute particles disrupt the orderly arrangement of solvent molecules, preventing them from forming a solid structure easily. Consequently, a lower temperature is needed to freeze the solution.
  • The solute interferes with the formation of the crystal lattice of the solid phase.
  • This disruption requires the system to release more thermal energy (heat) to solidify.
Freezing point depression is a key concept in colligative properties, which depend on the quantity of solute particles rather than their identity. So, whether you add sugar or salt, the key factor is the number of particles disrupting the solvent's molecular structure.
Boiling Point Elevation
Just as adding a solute to a solvent causes freezing point depression, it also leads to boiling point elevation. Essentially, this means the temperature at which the liquid turns into a gas is raised.
But why does this occur? It boils down to the intermolecular interactions.
  • The solute molecules create additional forces that the solvent molecules must overcome to vaporize.
  • As a result, more energy (higher temperature) is needed for the solvent molecules to escape into the vapor phase.
Thus, the boiling point of the solution increases when solute is added. This is also a colligative property, dependent on the number of solute particles. It's a crucial principle for many practical applications, like antifreeze in car radiators, which allow the liquid to remain in its current phase over a broader temperature range.
Intermolecular Forces
Intermolecular forces are forces of attraction or repulsion between molecules. They are a key factor in understanding phase changes and the impact of solutes on solvents.
These forces include:
  • Hydrogen bonds - strong attractions between molecules with hydrogen atoms bonded to electronegative elements like oxygen or nitrogen.
  • Dipole-dipole interactions - forces between polar molecules with positive and negative ends.
  • Van der Waals forces - weaker interactions primarily due to temporary dipoles in otherwise non-polar molecules.
When a solute is added, these intermolecular forces need to be overcome for phase changes to occur. Solute molecules create new intermolecular forces with solvent molecules, affecting the thermal energy required for changes between phases. This affects both boiling and freezing points, showing how crucial intermolecular forces are in determining the physical properties of liquids.
Temperature Range Extension
When a solute modifies the solvent's boiling and freezing points, it extends the temperature range where the liquid phase exists.
This is crucial in many real-world applications where maintaining a liquid state over varied temperatures is essential.
  • This principle is harnessed in antifreeze solutions, preventing water freezing in car radiators and ensuring smoother operations in cold weather.
  • In cooking, adding salt to boiling water increases the boiling point, allowing for faster cooking by reaching higher temperatures.
Thus, the statement of extending the liquid phase over a larger temperature range highlights the role of solutes in broadening the temperature spectrum where a solvent retains its liquid state. This concept underlies much of the work in chemistry and engineering, impacting everything from culinary to industrial applications.

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Most popular questions from this chapter

The freezing point of \(t\) -butanol is \(25.50^{\circ} \mathrm{C}\) and \(K_{\mathrm{f}}\) is \(9.1^{\circ} \mathrm{C} \cdot \mathrm{kg} / \mathrm{mol}\) Usually \(t\) -butanol absorbs water on exposure to air. If the freezing point of a 10.0 -g sample of \(t\) -butanol is \(24.59^{\circ} \mathrm{C},\) how many grams of water are present in the sample?

You have a solution of two volatile liquids, A and B (assume ideal behavior). Pure liquid A has a vapor pressure of 350.0 torr and pure liquid B has a vapor pressure of 100.0 torr at the temperature of the solution. The vapor at equilibrium above the solution has double the mole fraction of substance A that the solution does. What is the mole fraction of liquid A in the solution?

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: 鈥淭he 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.鈥� What do you say to your friend?

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

Calculate the solubility of \(\mathrm{O}_{2}\) in water at 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 \((\mathrm{mol} / \mathrm{L})\)
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