Question

Explain each of the following statements: (a) The boiling point of seawater is higher than that of pure water. (b) Carbon dioxide escapes from the solution when the cap is removed from a carbonated soft drink bottle. (c) Molal and molar concentrations of dilute aqueous solutions are approximately equal. (d) In discussing the colligative properties of a solution (other than osmotic pressure), it is preferable to express the concentration in units of molality rather than in molarity. (e) Methanol (b.p. \(65^{\circ} \mathrm{C}\) ) is useful as an antifreeze, but it should be removed from the car radiator during the summer season.

Short answer

(a) Seawater's higher boiling point is due to dissolved salts. (b) CO2 escapes as pressure is reduced. (c) Molal and molar are nearly equal in dilute solutions. (d) Molality is temperature-independent for colligative properties. (e) Methanol should be removed in summer due to low boiling point.

Step-by-step solution

Understanding Boiling Points

(a) Seawater contains dissolved salts and other impurities. These impurities increase the boiling point of water by causing a boiling point elevation, a colligative property. Therefore, seawater boils at a higher temperature than pure water due to its solute particles.

Gas Solubility and Pressure

(b) Carbon dioxide is dissolved under pressure in a carbonated drink. When the bottle cap is removed, the pressure inside decreases, reducing the solubility of carbon dioxide in the liquid. As a result, carbon dioxide escapes the solution, forming bubbles.

Concentration Units in Dilution

(c) In dilute aqueous solutions, the density of the solution is close to that of pure water and the mass of the solvent is nearly equal to the volume of the solution. Hence, molarity (mol/L) and molality (mol/kg of solvent) become approximately equal in value under these conditions.

Colligative Properties and Units

(d) Colligative properties depend on the number of solute particles, not their identity. Molality is preferred over molarity for these calculations because molality is temperature-independent, whereas molarity can change with temperature due to the expansion or contraction of the solution.

Seasonal Use of Antifreeze

(e) Methanol lowers the freezing point of water, making it useful as an antifreeze. However, during summer, temperatures rise, and methanol's low boiling point (65°C) could cause it to evaporate, possibly leading to loss of antifreeze properties and potential damage to the cooling system if it boils off.

Key concepts

Boiling Point Elevation

When you add a solute like salt to a solvent such as water, an interesting phenomenon called boiling point elevation occurs. This means that the boiling point of the solution is higher than that of the pure solvent. Why does this happen? It's all about vapor pressure. In a solution, solute particles disrupt the solvent's ability to evaporate. This requires more energy (higher temperature) to reach the boiling point.

In simple terms, the added solute particles take up space at the liquid's surface. This decreases the rate at which molecules escape into the vapor phase. So, the liquid must be heated to a higher temperature for the vapor pressure to equal atmospheric pressure.

This principle explains why seawater boils at a higher temperature than pure water. Its dissolved salts increase the solvent’s boiling point, demonstrating a key colligative property.

Gas Solubility

When you open a carbonated drink, the fizz you see is a fascinating display of gas solubility in action. Carbonated drinks contain carbon dioxide (CO₂) under high pressure, making it more soluble in the liquid. Once the cap is removed, the pressure drops, reducing CO₂ solubility. As a result, the gas escapes, forming bubbles.

Solubility of a gas depends significantly on pressure, described by Henry's Law, which states that the solubility of a gas is directly proportional to the pressure of the gas above the liquid. This is why your soda stays fizzy when closed but loses its fizz more quickly once opened.

• High pressure = more gas dissolves.
• Lower pressure = gas escapes.

Understanding this concept not only helps in chemistry but also explains many everyday phenomena.

Molality vs Molarity

To accurately describe concentrations in solutions, one should consider the units of molality and molarity. Both are measures of concentration, but they are quite different.

Molarity (M) is defined as the number of moles of solute per liter of solution, while molality (m) is the number of moles of solute per kilogram of solvent. It's important to note the differences in usage.

• Molarity depends on the total volume of the solution, which can change with temperature.
• Molality, however, depends on the mass of the solvent, remaining unaffected by temperature changes.

This means that molality is preferred for certain calculations, like colligative properties, where temperature independence is crucial. In dilute solutions, however, molarity and molality values can be close, which helps when using either measurement.

Temperature Dependence of Concentration Units

Temperature can have a striking impact on concentration measurements, particularly when using molarity. As temperature changes, so does the volume of the solution due to expansion or contraction. Since molarity is based on volume, it is directly affected by these temperature shifts.

Conversely, molality remains unaffected because it's based on the mass of the solvent, which stays constant with temperature changes. This makes molality a more reliable measure for experiments when temperature varies.

Choosing the right concentration unit is essential for precise scientific measurements. It's especially important in solutions where colligative properties are studied, as they depend solely on the number of solute particles, not the volume of the solution.

So remember: use molality in scenarios where temperature change might affect your results, as it's the more stable and consistent choice.