What is the specific definition of solubility for a solid in a liquid?

Solubility is the maximum mass of solute that can dissolve in $100\text{ g}$ of a solvent at a specific temperature. It is a measurement of the saturation point of a specific solute-solvent pair.

What is the mathematical formula used to calculate solubility from experimental data?

The formula is $\text{Solubility} = \frac{\text{mass of solute (g)}}{\text{mass of water (g)}} \times 100$. This expresses the result in the standard units of grams per $100\text{ g}$ of water.

How does solubility differ from the concentration of a solution?

Concentration is a general measure of how much solute is in a solvent, whereas solubility is the specific concentration at the point of saturation. A solution can have many concentrations, but only one solubility value at a given temperature.

What is the difference between a saturated and an unsaturated solution?

A saturated solution contains the maximum amount of dissolved solute at a given temperature, often evidenced by excess solid. An unsaturated solution contains less than the maximum amount and can still dissolve more solute.

How does the solubility of most solids change as the temperature of the solvent increases?

For most solids, solubility increases as temperature increases because the added thermal energy helps overcome the forces holding the solute particles together. This allows more solute to be integrated into the solvent phase.

Why must the solution be warmed above the target temperature when preparing the saturated solution?

Warming ensures that more than enough solute can be dissolved initially. As the solution cools to the target temperature, the excess solute crystallizes out, ensuring the remaining liquid is perfectly saturated at that exact temperature.

What does the term 'heating to constant mass' mean in a practical context?

It is the process of heating, cooling, and weighing a sample repeatedly until two consecutive mass readings are identical. This procedure proves that all water has been evaporated and only the dry solid remains.

What experimental error occurs if 'spitting' takes place during the evaporation of the solvent?

Spitting causes a loss of both solvent and dissolved solute from the evaporating basin. This results in a measured mass of dry solute that is too low, leading to an underestimate of the actual solubility.

Why is it a mistake to allow undissolved solid to enter the evaporating basin when sampling the saturated solution?

The undissolved solid was not part of the liquid equilibrium. If transferred, it will be weighed as part of the 'dissolved' solute mass, leading to an artificially high and incorrect solubility value.

What is the consequence of not drying the evaporating basin completely before the final weighing?

Residual water will add to the mass of the solute, making it appear as though more solid was dissolved than actually was. This results in an overestimate of the solubility of the substance.

Practical: Investigate the Solubility of a Solid in Water at a Specific Temperature | Pearson Edexcel IGCSE Chemistry

1. Principles of Chemistry

Practical: Investigate the Solubility of a Solid in Water at a Specific Temperature

Summary

This practical investigation determines the solubility of a solid by creating a saturated solution at a precisely controlled temperature, then isolating and weighing the dry solute after complete evaporation of the solvent. The core knowledge focuses on the relationship between mass, temperature, and saturation limits, requiring rigorous techniques like 'heating to constant mass' to ensure quantitative accuracy.

1. Definition & Core Concepts

Solubility is defined as the maximum mass of a solute that can dissolve in a specific volume of solvent (usually $100\text{ g}$ of water) at a given temperature. It represents a physical property of a substance that varies significantly depending on the thermal energy of the system.

A saturated solution is a chemical equilibrium state where the solvent has dissolved the maximum possible amount of solute at that specific temperature. In a practical setting, this is confirmed by the presence of undissolved solid at the bottom of the container, indicating no more can be integrated into the liquid phase.

The solute refers to the solid substance being dissolved (e.g., copper(II) sulfate), while the solvent is the liquid medium (usually water) that facilitates the dissolution process.

2. Underlying Principles

The investigation relies on the principle of mass conservation during state changes; while water is removed via evaporation, the mass of the non-volatile solute remains constant. By measuring the mass of the solution before and after evaporation, the mass of both the solvent and solute can be determined independently.

Solubility is highly temperature-dependent; for most solids, an increase in temperature increases kinetic energy, allowing more solute particles to be broken from the solid lattice and surrounded by solvent molecules. Therefore, the temperature must be held constant and measured precisely at the moment the saturated solution is sampled.

Heating to constant mass is a critical procedural principle used to verify complete dehydration. By repeatedly heating, cooling, and weighing the sample until the mass no longer changes, the experimenter ensures that all volatile solvent has been removed, leaving only the pure dry solute.

Isometric diagram of a lab setup showing an evaporating basin on a tripod being heated by a Bunsen burner to obtain dry solute crystals.

3. Methods & Techniques

4. Key Distinctions

5. Exam Strategy & Tips

6. Common Pitfalls & Misconceptions

Practical: Investigate the Solubility of a Solid in Water at a Specific Temperature

Summary

1. Definition & Core Concepts

The solute refers to the solid substance being dissolved (e.g., copper(II) sulfate), while the solvent is the liquid medium (usually water) that facilitates the dissolution process.

2. Underlying Principles

Isometric diagram of a lab setup showing an evaporating basin on a tripod being heated by a Bunsen burner to obtain dry solute crystals.

3. Methods & Techniques

Step 1: Preparation of Saturated Solution: Heat water slightly above the target temperature and stir in the solute until excess solid remains. Cooling the solution back to the exact target temperature ensures it is perfectly saturated at that point.

Step 2: Measurement of Masses: Record the mass of the empty basin ( $m_1$ ), the basin with the saturated solution ( $m_2$ ), and finally the basin with the dry solid ( $m_3$ ). Accurate weighing is the foundation of the final calculation.

Step 3: Controlled Evaporation: Heat the solution gently to remove the solvent. Rapid boiling should be avoided to prevent 'spitting,' where small droplets of solution are lost, leading to an underestimate of the solute mass.

Step 4: Final Verification: Use the heating-to-constant-mass technique to confirm all moisture is gone. This involves multiple cycles of heating, cooling in a desiccator, and reweighing until the mass value stabilizes.

4. Key Distinctions

Solubility vs. Concentration: While concentration refers to any ratio of solute to solvent, solubility refers specifically to the concentration at the saturation limit. A solution can be concentrated but not saturated, or saturated but dilute (for poorly soluble substances).
Evaporation vs. Spitting: Evaporation is the desired phase change of the solvent, while spitting is a mechanical loss of solution. Spitting introduces a significant error because it removes both solvent and dissolved solute, whereas evaporation only removes the solvent.
Saturated vs. Unsaturated:

Feature	Saturated Solution	Unsaturated Solution
Definition	Contains max solute possible	Can dissolve more solute
Visual Cue	Excess solid at bottom	All solid dissolved
Use in Practical	Required for solubility calc	Not suitable for calc

5. Exam Strategy & Tips

The Calculation Flow: Always follow a three-step subtraction process. First, subtract basin mass from the initial solution mass to find the total solution mass. Second, subtract basin mass from the dry solid mass to find solute mass. Third, subtract solute mass from total solution mass to find solvent (water) mass.
Solubility Units: Ensure your final answer is expressed as $g$ per $100\text{ g}$ of water. If you found $5\text{ g}$ of solute in $25\text{ g}$ of water, you must scale it up: $\frac{5}{25} \times 100 = 20\text{ g}/100\text{g}$ .
Reasonableness Check: Check the values against common solubility curves. Most salts have solubilities between $10\text{ g}$ and $100\text{ g}$ per $100\text{ g}$ of water; if your answer is outside this range, re-verify your subtractions.
Identifying Errors: If an exam question mentions a 'higher than expected' solubility result, check if undissolved solid was accidentally transferred into the basin. If the result is 'lower than expected,' check if spitting occurred during heating.

6. Common Pitfalls & Misconceptions

The 'Spitting' Error: Students often use high heat to speed up the experiment, which causes the solution to spit. This results in an underestimate of the solute mass ( $m_3$ ) and an incorrect solubility value.

Transferring Solid: When pouring the saturated solution into the basin, any undissolved solid that falls in will be counted as 'dissolved' once the water is gone. This leads to an artificially high solubility result because that solid was never part of the dissolved equilibrium.

Incomplete Drying: Failing to heat to constant mass leaves residual water trapped in the crystals. This increases the measured mass of the 'dry' solid, resulting in an overestimate of solubility.