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Chapter 12: Problem 3

Briefly describe the solution process at the molecular level. Use the dissolution of a solid in a liquid as an example.

Short Answer

Expert verified
Dissolution of a solid in a liquid involves breaking of the solute intermolecular forces by the solvent molecules, leading to the dispersion of solute in the solvent. The extent of dissolution depends on the solute-solvent interactions and the balance between them, which is defined as solubility. Once the solution is saturated, it achieves equilibrium where the rate of dissolution equals the rate of crystallization.

Step by step solution

01

Understanding Dissolution

The dissolution of a solid in a liquid starts when the solvent molecules (of the liquid) come into contact with the solute molecules (of the solid). The solvent molecules apply a force that breaks the intermolecular forces of the solute, allowing the solute molecules to disperse within the solvent.
02

Role of Solute-Solvent Interactions

Solute-solvent interactions play a vital role in the process of dissolution. Depending on the characteristics of the molecules involved, solute-solvent interactions can be of different types such as ion-dipole forces, hydrogen bonding, or van der Waals forces. If the solute-solvent forces are stronger than the solute-solute and solvent-solvent interactions, the solute will dissolve in the solvent.
03

Concept of Solubility

Solubility refers to the maximum amount of solute that can dissolve in a particular solvent at a given temperature. It's determined by the balance of intermolecular forces between the solute and the solvent molecules. If the resulting solution achieves equilibrium, where the rate of dissolution equals the rate of crystallization, the solution is said to be saturated.

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Most popular questions from this chapter

Write the equation relating osmotic pressure to the concentration of a solution. Define all the terms and specify their units.

Arrange the following solutions in order of decreasing freezing point: \(0.10 \mathrm{~m} \mathrm{Na}_{3} \mathrm{PO}_{4}, 0.35 \mathrm{~m} \mathrm{NaCl}\) \(0.20 \mathrm{~m} \mathrm{MgCl}_{2}, 0.15 \mathrm{~m} \mathrm{C}_{6} \mathrm{H}_{12} \mathrm{O}_{6}, 0.15 \mathrm{~m} \mathrm{CH}_{3} \mathrm{COOH}\)

An example of the positive deviation shown in Figure 12.8 (a) is a solution made of acetone \(\left(\mathrm{CH}_{3} \mathrm{COCH}_{3}\right)\) and carbon disulfide \(\left(\mathrm{CS}_{2}\right) .\) (a) Draw Lewis structures of these molecules. Explain the deviation from ideal behavior in terms of intermolecular forces. (b) A solution composed of 0.60 mole of acetone and 0.40 mole of carbon disulfide has a vapor pressure of \(615 \mathrm{mmHg}\) at \(35.2^{\circ} \mathrm{C} .\) What would be the vapor pressure if the solution behaved ideally? The vapor pressure of the pure solvents at the same temperature are acetone: \(349 \mathrm{mmHg} ;\) carbon disulfide: \(501 \mathrm{mmHg}\). (c) Predict the sign of \(\Delta H_{\text {soln. }}\)

Ammonia \(\left(\mathrm{NH}_{3}\right)\) is very soluble in water, but nitrogen trichloride \(\left(\mathrm{NCl}_{3}\right)\) is not. Explain.

Write the equations relating boiling-point elevation and freezing-point depression to the concentration of the solution. Define all the terms, and give their units.
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