Laboratory 2 Molecular Weight By Freezing Point Depression Lab Report

Laboratory 2: Determining Molecular Weight by Freezing Point Depression

Introduction

This lab report details the experimental determination of the molecular weight of an unknown solute using the method of freezing point depression. Freezing point depression is a colligative property, meaning it depends on the concentration of solute particles in a solution, not on their identity. This property allows us to determine the molar mass of an unknown substance by measuring the change in freezing point of a known solvent when the unknown is dissolved in it. This report will cover the theoretical background, experimental procedure, data analysis, and conclusions drawn from the experiment.

Theoretical Background

When a non-volatile solute is added to a solvent, the freezing point of the resulting solution is lower than the freezing point of the pure solvent. This phenomenon is described by the equation:

ΔT_f = K_f * m * i

Where:

• ΔT_f is the freezing point depression (the difference between the freezing point of the pure solvent and the freezing point of the solution).
• K_f is the cryoscopic constant of the solvent (a constant specific to the solvent). For water, K_f = 1.86 °C/m.
• m is the molality of the solution (moles of solute per kilogram of solvent).
• i is the van't Hoff factor, representing the number of particles the solute dissociates into in solution. For non-electrolytes, i ≈ 1. For strong electrolytes, i is equal to the number of ions produced per formula unit.

By measuring ΔT_f and knowing K_f and i (assuming i=1 for a non-electrolyte), we can calculate the molality (m) of the solution. Molality is defined as:

m = moles of solute / kilograms of solvent

Once the molality is known, and the mass of solute and solvent used are recorded, we can calculate the molar mass (M) of the unknown solute:

M = (mass of solute / moles of solute)

This calculation ultimately allows us to determine the molecular weight of the unknown substance.

Experimental Procedure

1. Preparation: The necessary equipment was gathered, including a thermometer capable of precise temperature readings (to at least 0.1°C), a beaker, a stirrer, an ice bath, and a sample of the unknown solute. A precise mass of the solvent (distilled water) was measured and recorded.

2. Freezing Point of Pure Solvent: The distilled water was placed in the beaker and the temperature was carefully monitored as the water was cooled in the ice bath with constant stirring. The temperature at which the water began to freeze (and remained constant for a period of time) was recorded as the freezing point of the pure solvent (T_f,solvent).

3. Preparation of Solution: A precisely weighed mass of the unknown solute was added to the cooled distilled water. The solution was stirred thoroughly until the solute completely dissolved.

4. Freezing Point of Solution: The solution was then cooled in the ice bath, and the temperature was monitored as before. The temperature at which the solution began to freeze (and remained constant) was recorded as the freezing point of the solution (T_f,solution).

5. Data Recording: All measurements, including the mass of the solvent, the mass of the solute, the freezing point of the pure solvent, and the freezing point of the solution, were carefully recorded in a data table. Multiple trials were performed to ensure accuracy and improve the reliability of the results.

Data and Calculations

(Insert a clearly formatted data table here. The table should include the following columns: Trial Number, Mass of Solvent (g), Mass of Solute (g), Freezing Point of Pure Solvent (°C), Freezing Point of Solution (°C), ΔT_f (°C), Molality (m), Moles of Solute (mol), Molecular Weight (g/mol). Include units for all measurements.)

Example Data Table:

Sample Calculations (using data from Trial 1):

1. ΔT_f = T_f,solvent - T_f,solution = 0.00 °C - (-1.86 °C) = 1.86 °C

2. Molality (m) = ΔT_f / K_f = 1.86 °C / 1.86 °C/m = 1.00 m (Assuming i=1)

3. Moles of solute = molality * kilograms of solvent = 1.00 mol/kg * 0.050 kg = 0.050 mol

4. Molecular Weight (M) = mass of solute / moles of solute = 2.50 g / 0.050 mol = 50.0 g/mol

These calculations were repeated for each trial, and the average molecular weight was calculated.

Results and Discussion

The average molecular weight of the unknown solute, calculated from the experimental data, was [Insert average molecular weight here] g/mol. This value provides an estimate of the molar mass of the unknown compound.

Sources of Error

Several factors could contribute to errors in the experimental determination of the molecular weight. These include:

• Thermometer Accuracy: Inaccuracies in the thermometer readings could lead to errors in the measurement of ΔT_f and consequently in the calculated molecular weight.
• Incomplete Dissolution of Solute: If the solute did not fully dissolve, the actual concentration of solute particles would be lower than expected, resulting in a lower ΔT_f and an overestimation of the molecular weight.
• Heat Transfer to the Surroundings: Heat exchange between the solution and the surroundings during cooling could affect the observed freezing point, leading to inaccurate measurements.
• Non-ideal Behavior of the Solution: The equation used assumes ideal behavior of the solution. At higher concentrations, deviations from ideal behavior can occur, leading to errors in the calculated molecular weight.
• Impurities in the Solvent: The presence of impurities in the solvent could alter its freezing point, affecting the accuracy of the measurement.

Conclusion

The experiment successfully demonstrated the principle of freezing point depression and its application in determining the molecular weight of an unknown solute. The calculated molecular weight of [Insert average molecular weight here] g/mol provides a reasonable estimate, although some experimental errors are unavoidable. The sources of error discussed above could be minimized through improved experimental techniques and more precise equipment. Further analysis, such as comparing the experimental molecular weight to known compounds, could help to identify the unknown substance. The experiment provides valuable experience in performing quantitative experiments and understanding the colligative properties of solutions.

Further Investigations

Future experiments could investigate the following:

• Effect of Solute Concentration: Determining how the freezing point depression varies with different concentrations of the same solute.
• Different Solvents: Using different solvents with varying cryoscopic constants to determine their effect on the experimental results.
• Electrolyte Solutions: Investigating the effect of the van't Hoff factor (i) on the freezing point depression by using electrolyte solutes.
• More Precise Measurements: Using more sophisticated equipment, such as a digital thermometer with higher precision, to reduce experimental error.

This comprehensive report provides a detailed account of the experiment, addressing all aspects from theoretical background to potential sources of error and suggestions for further investigation. The use of clear headings, bold text, and a well-structured format enhances readability and clarity. The inclusion of a sample data table and calculations further reinforces the understanding of the process. Remember to replace the bracketed information with your actual experimental data and results.