Problem 95 - Chemistry Matter and Change

Question

Define the term colligative property.

Step-by-Step Solution

Answer

Colligative properties depend on the quantity, not the type, of solute particles in a solution.

Step 1: Understanding the Basic Concept

Colligative properties are physical properties of solutions that depend on the number of particles in a given amount of solvent, not the type or nature of the particles. This means these properties are affected by how many solute molecules or ions are present rather than the specific chemical identity.

Step 2: Examples of Colligative Properties

There are four main colligative properties: boiling point elevation, freezing point depression, vapor pressure lowering, and osmotic pressure. Each of these properties changes with the concentration of solute particles in the solution.

Step 3: Relating to Real World Applications

An example of a colligative property is boiling point elevation, which explains why adding salt to water makes it boil at a higher temperature. This is because salt increases the number of solute particles in the solution, a direct demonstration of how colligative properties work.

Key Concepts

Boiling Point Elevation

Boiling point elevation is a colligative property that refers to the increase in the boiling point of a solvent when a non-volatile solute is added. It's a direct outcome of the presence of solute particles, which disrupt the normal phase equilibrium in the solvent. This effect can be calculated using the formula:\[ \Delta T_b = i \cdot K_b \cdot m \]where:

• \( \Delta T_b \) is the boiling point elevation,
• \( i \) is the van't Hoff factor indicating the number of particles the solute separates into,
• \( K_b \) is the ebullioscopic constant of the solvent,
• \( m \) is the molality of the solution, a measure of the concentration in moles of solute per kilogram of solvent.

Adding salt to water, for instance, increases the water's boiling point. When cooking pasta, for example, the water will boil at a slightly higher temperature, helping to cook food more efficiently. This is because the salt ions increase the number of particles in the water, illustrating the dependency of this property on particle concentration instead of their identity.

Freezing Point Depression

Freezing point depression is another colligative property where the addition of a solute results in the lowering of the freezing point of the solvent. When solute particles are present, they disrupt the crystal formation needed for a substance to freeze. This phenomenon can be calculated using:\[ \Delta T_f = i \cdot K_f \cdot m \]where:

• \( \Delta T_f \) is the freezing point depression,
• \( i \) is the van't Hoff factor,
• \( K_f \) is the cryoscopic constant, reflecting how much the freezing point lowers for each molal unit of solute added,
• \( m \) is the molality of the solution.

Common in everyday life, this principle explains why salt is used to melt ice on roads during winter. As salt dissolves, it lowers the melting/freezing point of the ice, causing it to melt even at temperatures where pure ice would remain solid. The effect is critical for safe road conditions during icy weather.

Osmotic Pressure

Osmotic pressure is the pressure required to stop the flow of solvent molecules through a semi-permeable membrane from a pure solvent into a solution. This movement arises because molecules from the pure solvent naturally move toward the solution to equalize concentrations. The pressure needed to counterbalance this natural flow is termed osmotic pressure, calculated with:\[ \Pi = i \cdot c \cdot R \cdot T \]where:

• \( \Pi \) is the osmotic pressure,
• \( i \) is the van't Hoff factor,
• \( c \) is the molar concentration of the solution,
• \( R \) is the ideal gas constant,
• \( T \) is the temperature in Kelvin.

This property is vital in biological systems. For example, plants rely on osmotic pressure to absorb water through their roots. It also explains how nutrient and waste transport occur across cell membranes, maintaining essential life processes. These examples highlight the importance of osmotic pressure both in nature and in medical applications like intravenous fluid administration.