1. The heat of solution of salta may either be endothermic or exothermic. What factors determine this?
2. As a solution freezes, why does the freezing temperature continue to decrease?
3. When the concentration of very dilute aqueous solutions are calculated, why are the values of molarity and molality the same?
1. Basically how much energy is need to break about the crystal lattice vs the energy produced by hydration of the resulting ions.
2. As the "pure" water freezes, the solution that is left is more concentrated in the solute; hence, a more concentrated solution lowers the freezing point.
3. Molarity =# mols/L solution; molality = # mols/kg solvent. As the # mols becomes smaller and smaller, at "infinite" dilution density of the solutins are the same because the L of solution is closer and closer to a kg of solvent.
1. The heat of solution of a salt can be either endothermic or exothermic, depending on several factors. The key factors that determine this are the energy required to break the crystal lattice of the salt and the energy released when the salt molecules dissolve and are hydrated by water molecules.
To understand this conceptually, imagine salt crystals composed of positively charged cations (e.g., Na+) and negatively charged anions (e.g., Cl-). When these salt crystals are placed in water, the water molecules surround the ions, breaking the ionic bonds and allowing the ions to become solvated.
If the energy required to break the crystal lattice is less than the energy released during hydration, the process is exothermic. This means that more energy is released as heat during dissolution than is absorbed from the surroundings. In this case, the temperature of the solution may increase.
Conversely, if the energy required to break the crystal lattice is greater than the energy released during hydration, the process is endothermic. This means that more energy is absorbed from the surroundings than is released as heat during dissolution. In this case, the temperature of the solution may decrease.
Factors such as the nature of the salt (e.g., its crystal structure, bond strength), the size and charge of the ions, and the solvent (e.g., water) all play a role in determining whether the heat of solution is endothermic or exothermic.
2. When a solution freezes, the freezing temperature continues to decrease because freezing is a process of the formation of a solid from a liquid. In a solution, the presence of solute molecules disrupts the regular arrangement of solvent molecules, making it more difficult for the solvent molecules to come together and form a solid crystal lattice.
The freezing point of a solution is lower than that of the pure solvent because adding a solute lowers the vapor pressure of the solvent. This means that the solvent molecules have a reduced tendency to escape from the liquid phase and form a solid. Therefore, the solution requires a lower temperature to reach the point where the liquid turns into a solid.
Additionally, the solute molecules can become incorporated into the solid crystal lattice as the solution freezes, further lowering the freezing point. The solute molecules take up space in the lattice and alter the organization of the solvent molecules, resulting in a decrease in freezing temperature compared to the pure solvent.
3. When calculating the concentration of very dilute aqueous solutions, both molarity (M) and molality (m) may yield the same values. Molarity refers to the number of moles of solute per liter of solution, while molality refers to the number of moles of solute per kilogram of solvent.
At very low concentrations, the density of the solution approaches that of the solvent, so the volume of the solution is effectively equal to the mass of the solvent. In these extremely dilute solutions, the difference in scale between a liter of solution and a kilogram of solvent becomes negligible.
Since the differences in volume and mass are small, the values of molarity and molality can be considered approximately equal. As the concentration of the solution decreases, the difference between molarity and molality becomes less important and both can be used interchangeably to measure the concentration of the solute.