Quantitative Analysis of Hydrates: Copper(II) Sulfate and Related Compounds

Hydrate of Copper(II) Sulfate

Fundamental Concepts

The accurate determination of mass is a foundational technique in general chemistry, especially in quantitative analysis. This experiment focuses on the analysis of hydrates, specifically copper(II) sulfate pentahydrate, to determine the percentage of water present in a hydrate sample.

• Mass: A measure of the amount of matter in a sample, typically determined using a balance.

• Quantitative Analysis: The branch of chemistry concerned with measuring the quantities of substances present in a sample.

• Gravimetric Analysis: A quantitative method where the amount of an analyte is determined by measuring mass.

• Hydrate: A compound that contains water molecules weakly bound within its crystal structure. The water can be removed by heating.

Accurate weighing and careful handling of balances are essential for reliable results in gravimetric experiments.

Laboratory Balances

Balances are used to measure mass with high precision. Proper use and care are crucial for accurate measurements.

• Handle balances carefully; do not jar or move them unnecessarily.

• Always use the same balance for a given experiment to ensure consistency.

• Never weigh hot objects; allow them to cool to room temperature before weighing.

• Keep the balance clean and free of chemicals.

• Record all measurements to the correct number of significant figures.

Quantitative Analysis of a Hydrate

Purpose and Overview

This experiment provides practice in gravimetric analysis by determining the percentage of water in a hydrate. The process involves heating a known mass of hydrate to drive off water and measuring the mass loss.

• Hydrate Used: Copper(II) sulfate pentahydrate, CuSO4·5H2O

• Goal: Determine the percentage by mass of water in the hydrate and compare it to the theoretical value.

Experimental Procedure

1. Clean a crucible and lid, then heat over a flame for several minutes. Cool to room temperature.

2. Weigh the empty crucible and lid.

3. Add approximately 2 g of the hydrate to the crucible, weigh, and record the mass.

4. Heat the crucible with the lid slightly ajar to allow water vapor to escape. Heat until the sample appears white or grayish (anhydrous).

5. Cool the crucible to room temperature and weigh again.

6. Repeat heating and weighing until two consecutive masses agree within 0.01 g (constant mass).

7. Dispose of chemicals as instructed.

Balanced Chemical Equation

The dehydration of copper(II) sulfate pentahydrate can be represented as:

$\mathrm{CuSO_4 \cdot 5H_2O (s) \rightarrow CuSO_4 (s) + 5H_2O (g)}$

Data Collection and Calculations

• Record the mass of the empty crucible, crucible plus hydrate, and crucible plus anhydrous salt after heating.

• Calculate the mass of hydrate used and the mass of water lost:

$\text{Mass of hydrate} = \text{(crucible + hydrate)} - \text{(empty crucible)}$

$\text{Mass of water lost} = \text{(crucible + hydrate)} - \text{(crucible + anhydrous salt)}$

• Calculate the percent by mass of water in the hydrate:

$\text{Percent by mass of water} = \frac{\text{mass of water lost}}{\text{mass of hydrate}} \times 100$

• Determine moles of water and moles of anhydrous salt:

$\text{Moles of water} = \frac{\text{mass of water lost}}{18.02\ \text{g/mol}}$

$\text{Moles of CuSO}_4 = \frac{\text{mass of anhydrous salt}}{159.61\ \text{g/mol}}$

• Calculate the experimental formula of the hydrate by finding the ratio of moles of water to moles of CuSO4.

Sample Data Table

Measurement	Value (g)
Mass empty crucible	37.808
Mass crucible + hydrate	40.457
Mass crucible + anhydrous CuSO4 (final)	39.554
Mass of hydrate	2.649
Mass of water lost	0.903

Example Calculations

• Percent by mass of water:

$\frac{0.903\ \text{g}}{2.649\ \text{g}} \times 100 = 34.1\%$

• Moles of water:

$\frac{0.903\ \text{g}}{18.02\ \text{g/mol}} = 0.0501\ \text{mol}$

• Moles of CuSO4:

$\frac{1.746\ \text{g}}{159.61\ \text{g/mol}} = 0.0109\ \text{mol}$

• Experimental formula:

$\frac{0.0501}{0.0109} \approx 4.59 \approx 5$

Thus, the experimental formula is CuSO4·5H2O.

Percent Error Calculation

$\text{Percent error} = \frac{|\text{experimental value} - \text{theoretical value}|}{\text{theoretical value}} \times 100$

For example, if the theoretical percent by mass of water is 36.1% and the experimental value is 39.1%:

$\frac{|39.1 - 36.1|}{36.1} \times 100 = 8.3\%$

Sources of Error and Analysis

• Incomplete dehydration: Not all water is removed if the sample is not heated sufficiently.

• Overheating: May cause decomposition of the salt, leading to mass loss beyond just water.

• Sample loss: Spillage or loss of sample during transfer can affect results.

• Weighing errors: Not allowing the crucible to cool before weighing can cause inaccurate mass readings due to convection currents.

Practice Problems and Applications

• Given the formula NiSO4·6H2O, calculate:

  • Moles of water per mole of hydrate: 6

  • Grams of water in 2.00 g of hydrate: $\frac{6 \times 18.02}{262.85} \times 2.00$

  • Theoretical percent by mass of water: $\frac{6 \times 18.02}{262.85} \times 100$

• Calculate percent error if experimental and theoretical values are known.

Summary of Key Results

• Mass of hydrate used: 2.649 g

• Mass of anhydrous salt: 1.746 g

• Mass of water lost: 0.903 g

• Experimental percent by mass of water: 39.1%

• Theoretical percent by mass of water: 36.1%

• Percent error: 5.5%

• Experimental formula: CuSO4·5H2O

Additional info:

• Hydrates are commonly used to illustrate stoichiometry and gravimetric analysis in introductory chemistry labs.

• Proper technique in heating, cooling, and weighing is essential for minimizing experimental error.