Lab 3: Molar Mass And Freezing Point Depression

Lab 3 Molar Mass And Freezing Point Depression

The objective of this experiment is to determine the molar mass of a known substance using its freezing point depression in a solution. You will be given three solutions in addition to the standard solution (water). The freezing point of water will be determined and then the freezing points of three other solutions will be measured. These solutions are composed of sodium chloride (NaCl), calcium chloride (CaCl2), and glycerol (C3H8O3). The data collected will be used to calculate the molar mass of each solute based on the freezing point depression observed.

Paper For Above instruction

The determination of molar mass via freezing point depression is a classical experiment in physical chemistry, illustrating colligative properties and molecular weight calculations. This experiment involves measuring the change in freezing point of water when various solutes are dissolved, which then allows for the calculation of molar mass based on colligative property equations. This paper discusses the principles, methodology, calculations, and implications of using freezing point depression to find molar masses of solutes such as NaCl, CaCl2, and glycerol.

Introduction

Colligative properties, including freezing point depression, depend solely on the number of solute particles in a solvent and not their identity. When a solute dissolves in a solvent, it disrupts the formation of a solid lattice, lowering the freezing point. Quantitatively, the freezing point depression is proportional to the molality of the solution and the van 't Hoff factor (i), which accounts for dissociation of electrolytes. This experiment aims to utilize this relationship to determine molar masses, highlighting the importance of accurate measurements and control of experimental conditions.

Methodology

The experimental procedure begins with preparing an ice bath with crushed ice and ensuring the temperature probe is properly calibrated. First, the initial freezing point of pure water is determined to establish a baseline. Then, solutions are prepared by adding known masses of NaCl, CaCl2, and glycerol to water, and their freezing points are measured under controlled conditions. The use of a Vernier temperature probe connected to LabQuest 2 interface allows continuous monitoring of temperature changes during the freezing process. By analyzing the flat portion of the temperature-time graph during freezing, the exact freezing point of each solution is obtained.

The solutions are prepared by dissolving a known mass (e.g., 2.0 g) of each solute in water (initially 15.0 mL), with additional solutions made by combining solutes with water and sodium chloride solution to compare effects. The freezing point depression (∆Tf) for each solution is calculated as the difference between the pure water freezing point and the measured freezing point for each solution. The van 't Hoff factor, molality, and molar mass are then calculated through established relationships.

Calculations

The fundamental equation connecting freezing point depression to molar mass is:

∆Tf = iKf m

Where:

  • ∆Tf = freezing point depression (°C)
  • i = van 't Hoff factor
  • Kf = cryoscopic constant (1.86 °C·kg/mol for water)
  • m = molality (mol/kg)

Rearranged to solve for molality:

m = ∆Tf / (iKf)

The molar mass (M) can then be calculated by:

M = (mass of solute in grams) / (number of moles)

Number of moles is determined from molality:

m = (moles of solute) / (kg solvent)

Hence, molar mass M = (mass of solute) / (molality × kg solvent)

For example, for sodium chloride:

moles of NaCl = mass of NaCl / Molar mass of NaCl

Using measured ∆Tf, i (which is 2 for NaCl since it dissociates into 2 particles), and the solvent mass, molar mass can be deduced.

The van 't Hoff factor (i) depends on each solute's dissociation: glycerol (i = 1, non-electrolyte), NaCl (i = 2), CaCl2 (i = 3). The aqueous solutions are prepared with precise concentrations, and the freezing points are measured accurately to minimize error.

Results and Data Analysis

The observed freezing points, ∆Tf, of each solution are tabulated alongside calculated molality, molar mass, and percent error relative to known molar masses (NaCl: 58.44 g/mol, CaCl2: 110.98 g/mol, glycerol: 92.09 g/mol). These calculations validate the theory of colligative properties and demonstrate the effectiveness of freezing point depression in molar mass determination.

Particularly, the mean freezing point during the plateau of the graph provides a precise ∆Tf value, which is then used in calculations. Experimental discrepancies can arise from measurement inaccuracies, incomplete dissociation, or temperature fluctuations. Nonetheless, with proper calibration and technique, results typically align closely with theoretical values.

Discussion

This experiment exemplifies how colligative properties serve as valuable tools in molecular weight determinations. The differences observed between experimental and theoretical molar masses reflect real-world factors such as purity of chemicals, measurement errors, and solution behavior. Notably, the value of i for electrolytes highlights the importance of accounting for dissociation. For instance, calcium chloride’s high dissociation enhances freezing point depression per mol, making it more efficient for practical applications like de-icing.

From an environmental perspective, NaCl and CaCl2 are commonly used for road salting. The data suggests CaCl2 provides a more significant depression per mole, meaning less salt may be necessary for effective melting, which could reduce environmental impact and economic costs. However, its higher corrosion potential must also be considered.

Conclusion

Freezing point depression provides a reliable approach for calculating molar masses of dissolved substances. Precise temperature measurements, proper solution preparation, and understanding dissociation behaviors are critical. This experiment not only reaffirms fundamental thermodynamic principles but also underscores their practical importance in industrial and environmental settings, chiefly in de-icing operations where both efficiency and sustainability are paramount.

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