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Calculate The Boiling Point of Each of The Following Solutions

Reviewed by Calculator Editorial Team

Determining the boiling point of solutions is essential in chemistry and chemical engineering. The boiling point elevation formula allows you to calculate how much a solvent's boiling point increases when a solute is added. This guide explains the formula, provides a calculator, and includes practical examples.

Introduction

The boiling point of a solution is higher than that of the pure solvent due to the presence of solute particles. This phenomenon is known as boiling point elevation. The increase in boiling point depends on the concentration of the solute and the solvent's properties.

Key factors affecting boiling point elevation include:

  • The number of solute particles dissolved in the solvent
  • The nature of the solute-solvent interactions
  • The temperature of the solution

Understanding boiling point elevation is crucial for processes like distillation, where precise temperature control is necessary.

Boiling Point Elevation Formula

The boiling point elevation (ΔTb) can be calculated using the following formula:

ΔTb = i × Kb × m

Where:

  • ΔTb = Boiling point elevation (in °C)
  • i = Van't Hoff factor (dimensionless)
  • Kb = Ebullioscopic constant (in °C·kg/mol)
  • m = Molality of the solution (in mol/kg)

The Van't Hoff factor accounts for the number of particles a solute dissociates into. For example, NaCl dissociates into 2 ions, so i = 2.

The ebullioscopic constant (Kb) is specific to each solvent. For water, Kb is approximately 0.512 °C·kg/mol.

How to Calculate Boiling Points

Step-by-Step Calculation Process

  1. Determine the molality (m) of the solution using the formula: m = moles of solute / kg of solvent
  2. Identify the Van't Hoff factor (i) based on the solute's dissociation behavior
  3. Find the ebullioscopic constant (Kb) for the solvent
  4. Plug these values into the boiling point elevation formula
  5. Calculate the boiling point elevation (ΔTb)
  6. Add ΔTb to the pure solvent's boiling point to get the solution's boiling point

Note: The boiling point of the pure solvent is typically 100°C for water at standard conditions.

Worked Examples

Example 1: Sodium Chloride in Water

Calculate the boiling point of a solution containing 25 g of NaCl in 250 g of water.

  1. Calculate moles of NaCl: 25 g / 58.44 g/mol ≈ 0.428 mol
  2. Calculate molality: 0.428 mol / 0.250 kg ≈ 1.712 mol/kg
  3. Van't Hoff factor (i) for NaCl = 2
  4. Ebullioscopic constant (Kb) for water = 0.512 °C·kg/mol
  5. ΔTb = 2 × 0.512 × 1.712 ≈ 1.74 °C
  6. Boiling point = 100°C + 1.74°C = 101.74°C

Example 2: Sucrose in Water

Calculate the boiling point of a solution containing 50 g of sucrose in 500 g of water.

  1. Calculate moles of sucrose: 50 g / 342.3 g/mol ≈ 0.146 mol
  2. Calculate molality: 0.146 mol / 0.500 kg ≈ 0.292 mol/kg
  3. Van't Hoff factor (i) for sucrose = 1 (does not dissociate)
  4. Ebullioscopic constant (Kb) for water = 0.512 °C·kg/mol
  5. ΔTb = 1 × 0.512 × 0.292 ≈ 0.15 °C
  6. Boiling point = 100°C + 0.15°C = 100.15°C

Frequently Asked Questions

What is the difference between boiling point elevation and freezing point depression?
Both phenomena are colligative properties that depend on the number of solute particles. Boiling point elevation increases the boiling point, while freezing point depression decreases the freezing point. The formulas are similar but use different constants (Kb for boiling point, Kf for freezing point).
How does temperature affect boiling point elevation?
The boiling point elevation formula assumes ideal behavior. In reality, at very high temperatures, the solvent's properties may change, and the formula becomes less accurate. For precise measurements, temperature corrections may be needed.
Can boiling point elevation be used for non-aqueous solvents?
Yes, the principle applies to any solvent. The ebullioscopic constant (Kb) and the pure solvent's boiling point will differ for each solvent. Common solvents like ethanol and benzene have their own Kb values.