What Is 0.0098 Boiling Point

What is the Boiling Point of 0.0098 Molar Solution? Understanding Colligative Properties

Determining the boiling point of a 0.0098 molar (M) solution requires understanding colligative properties, specifically boiling point elevation. This article will delve into the science behind boiling point elevation, explain how to calculate the boiling point of a 0.0098 M solution, discuss the factors influencing this calculation, and address frequently asked questions. We'll focus on providing a clear and comprehensive explanation, suitable for students and anyone interested in learning more about solution chemistry.

Introduction to Colligative Properties

Colligative properties are properties of solutions that depend on the concentration of solute particles, not their identity. This means that the property is affected by the number of particles dissolved, regardless of what those particles are. Four main colligative properties are:

• Vapor pressure lowering: The presence of a non-volatile solute reduces the vapor pressure of the solvent.
• Boiling point elevation: The boiling point of a solution is higher than that of the pure solvent.
• Freezing point depression: The freezing point of a solution is lower than that of the pure solvent.
• Osmotic pressure: The pressure required to prevent osmosis (the movement of solvent across a semipermeable membrane).

In this context, we are concerned with boiling point elevation. Adding a solute to a solvent makes it harder for the solvent molecules to escape into the gas phase, thus requiring a higher temperature to reach the boiling point.

Boiling Point Elevation: The Science Behind It

The extent of boiling point elevation is directly proportional to the molality (m) of the solution, a measure of concentration defined as moles of solute per kilogram of solvent. This relationship is expressed by the equation:

ΔTb = Kb * m * i

Where:

• ΔTb is the boiling point elevation (the difference between the boiling point of the solution and the boiling point of the pure solvent).
• Kb is the ebullioscopic constant (or molal boiling point elevation constant) of the solvent. This is a solvent-specific constant that represents the increase in boiling point for a 1 molal solution.
• m is the molality of the solution (moles of solute per kilogram of solvent).
• i is the van't Hoff factor. This factor accounts for the number of particles the solute dissociates into in solution. For non-electrolytes (substances that do not dissociate into ions), i = 1. For strong electrolytes (substances that completely dissociate into ions), i is equal to the number of ions produced per formula unit. For weak electrolytes, i is between 1 and the number of ions.

The importance of the van't Hoff factor (i) cannot be overstated. For example, a 0.0098 M solution of glucose (a non-electrolyte) will have a different boiling point elevation than a 0.0098 M solution of NaCl (a strong electrolyte that dissociates into two ions, Na+ and Cl−). For NaCl, i = 2.

Calculating the Boiling Point of a 0.0098 M Solution

To calculate the boiling point of a 0.0098 M solution, we need several pieces of information:

1. The identity of the solute and solvent: This is crucial for determining the van't Hoff factor (i) and the ebullioscopic constant (Kb) of the solvent.
2. The molality (m) of the solution: While we are given the molarity (0.0098 M), we need to convert it to molality. This conversion requires knowing the density of the solution, which is often not readily available and may vary with temperature. For dilute solutions, molarity and molality are often approximately equal, but this approximation is less accurate for more concentrated solutions. For simplicity and because the solution is dilute, we'll assume molarity and molality are approximately equal in this example. Thus m ≈ 0.0098 mol/kg.
3. The ebullioscopic constant (Kb) of the solvent: This value is specific to the solvent used. For water, Kb = 0.512 °C/m.

Example Calculation (Assuming the solute is a non-electrolyte and the solvent is water):

Let's assume our 0.0098 M solution is a non-electrolyte dissolved in water. Therefore, i = 1.

ΔTb = Kb * m * i = (0.512 °C/m) * (0.0098 m) * (1) = 0.0050176 °C

This means the boiling point is elevated by approximately 0.005 °C. Since the boiling point of pure water is 100 °C, the boiling point of this 0.0098 M solution would be approximately 100.005 °C.

Example Calculation (Assuming the solute is NaCl and the solvent is water):

Now let's consider a 0.0098 M solution of NaCl in water. NaCl is a strong electrolyte with i = 2.

ΔTb = Kb * m * i = (0.512 °C/m) * (0.0098 m) * (2) = 0.0100352 °C

In this case, the boiling point elevation is approximately 0.01 °C, resulting in a boiling point of approximately 100.01 °C.

Factors Influencing Boiling Point Elevation

Several factors can influence the accuracy of boiling point elevation calculations:

• Ion pairing: In solutions of strong electrolytes, ion pairing can occur, reducing the effective number of particles and lowering the boiling point elevation. This is more significant at higher concentrations.
• Solvent interactions: Strong interactions between solute and solvent molecules can affect the boiling point elevation.
• Non-ideality: The equation for boiling point elevation is based on ideal solution behavior. Deviations from ideal behavior can occur, particularly at higher concentrations.
• Experimental error: Errors in measuring the concentration of the solution or the boiling point can affect the accuracy of the calculation.

Frequently Asked Questions (FAQ)

• Q: Why is the boiling point elevation so small in this case?

  • A: The boiling point elevation is directly proportional to the molality of the solution. Since the solution is very dilute (0.0098 M), the elevation is small.

• Q: Can I use molarity instead of molality in the calculation?

  • A: For dilute solutions, molarity and molality are approximately equal. However, for more concentrated solutions, this approximation is inaccurate, and molality should be used.

• Q: What if the solute is a weak electrolyte?

  • A: For a weak electrolyte, the van't Hoff factor (i) will be between 1 and the number of ions produced upon dissociation. The exact value of i needs to be determined experimentally or estimated based on the degree of dissociation.

• Q: What happens if I don't know the identity of the solute?

  • A: Without knowing the identity of the solute, you cannot accurately determine the van't Hoff factor (i), making it impossible to accurately calculate the boiling point elevation.

• Q: What are the practical applications of understanding boiling point elevation?

  • A: Boiling point elevation is used in various applications, including determining the molar mass of unknown substances, purifying substances through fractional distillation, and understanding the behavior of solutions in various industrial processes.

Conclusion

Determining the exact boiling point of a 0.0098 M solution requires knowing the identity of the solute and solvent to accurately determine the van't Hoff factor (i) and the ebullioscopic constant (Kb). While we can approximate the boiling point using the given molarity and assuming the solution is dilute, more precise calculations require converting molarity to molality and accounting for non-ideal behaviors. The examples provided demonstrate how seemingly small changes in solute concentration and properties can cause noticeable changes in the boiling point, highlighting the importance of colligative properties in chemistry. This understanding is fundamental to various scientific fields and industrial applications. Remember that this explanation provides a foundational understanding and more complex scenarios may require more advanced calculations and considerations.