How does temperature affect the solubility of a compound?

Solubility typically increases with increasing temperature. For example, sugar dissolves more readily in hot coffee than in iced coffee. Sodium carbonate demonstrates this trend, with solubility rising from 7 g per 100 mL at 0°C to 44 g per 100 mL at 100°C, though some compounds show exceptions to this pattern.

What is a saturated solution and why does additional solute remain undissolved?

A saturated solution contains the maximum amount of solute dissolved in a given solvent volume. At this equilibrium point, dissolved ions balance with solid solute, so any additional solute cannot dissolve and remains as a precipitate. For instance, 36 grams of sodium chloride in 100 mL of water at room temperature creates a saturated solution.

How do you calculate the equilibrium constant for a dissolution process?

For dissolution equilibria, the equilibrium constant Kc uses concentrations of dissolved products in the numerator and reactants in the denominator, each raised to their stoichiometric coefficient. For ionic salts, the solubility product constant Ksp simplifies this by setting solid concentrations to 1, yielding Ksp = [A]^x[B]^y for compounds AxBy.

What does a negative ΔG value tell you about a dissolution reaction?

A negative ΔG indicates the reaction is spontaneous, meaning the system has sufficient energy to proceed without external input. Gibbs free energy, calculated from ΔG = ΔH – TΔS, reveals whether a dissolution process releases or requires energy. Negative values show the system naturally favors the dissolved state over the undissolved crystalline form.

How do enthalpy and entropy influence whether a solute dissolves?

Enthalpy measures heat content change; positive ΔH indicates endothermic dissolution (absorbs heat), while negative ΔH indicates exothermic dissolution (releases heat). Entropy measures disorder; dissolution increases solute disorder but orders solvent molecules around it. Together, these properties determine whether the system prefers the ordered crystalline state or the disordered dissolved state.

Why does sodium chloride dissolve differently in various solvents?

Solubility varies by solvent because different solvents have different polarities and solvation capabilities. Sodium chloride has a solubility of 36 g per 100 mL in water, but only 1.1 g per 100 mL in methanol and 0.034 g per 100 mL in dimethylformamide. Polar ionic compounds dissolve best in polar solvents that can effectively solvate the dissociated ions.

Solubility

Concepts Instructor Prep Student Protocol

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Solubility is a measure of a solute's ability to dissolve in a solvent. Different solutes have different solubilities. For example, sodium chloride has a solubility of 39 grams in 100 milliliters of water, while silver chloride is only 0.002 grams.

So why is this? Solubility depends on the physical and chemical properties of both the solute and solvent. You've seen this yourself. For example, you know that butter doesn't dissolve in water, but it does dissolve in olive oil.

The term 'like dissolves like' reminds us that a solvent dissolves a solute with a similar polarity. So polar solvents dissolve polar solutes, while non-polar solvents dissolve non-polar solutes. In addition, we can influence solubility through other factors like pH, temperature, and pressure.

For example, solubility typically increases with increasing temperature. Think about dissolving sugar in iced coffee versus hot coffee. Now, if a solute does dissolve in a solvent, only so much can be dissolved. When this limit is reached, the dissociated ions are in equilibrium with the solid and the solution is saturated. This means that if you add more of the solute, it won't dissolve.

We can quantify this equilibrium using an equilibrium constant K_c. This equation uses the concentrations of the dissolved products in the numerator and reactants in the denominator, each raised to the power of their stoichiometric coefficient. We can use this equilibrium constant to understand the thermodynamics of the dissolution process as it goes from its initial undissolved state to its final dissolved state.

The enthalpy, H, of a solution is a quantitative measure of the total heat content of the system, and ΔH describes the change in this heat content. If ΔH is positive, it indicates that the reaction absorbed heat, or is endothermic. When ΔH is negative, it indicates that the reaction releases heat, or is exothermic. Next, entropy, S, describes the degree of disorder in a system. ΔS is 0 for a reversible reaction, but is otherwise positive, as the disorder of a system prefers to increase.

Finally, the Gibbs free energy, G, is a measure of the energy that can be used to do work. It is calculated from ΔH and ΔS and is dependent on temperature. When ΔG is positive, the reaction is not spontaneous and energy must be put in for the reaction to proceed. When ΔG is less than 0, it indicates that the reaction is spontaneous.

These properties tell us a lot about how a solute dissolves in a solvent. For example, we can use enthalpy and entropy to learn whether or not the solute prefers to remain undissolved in its ordered crystalline form, or disordered in solution. And we can use Gibbs energy to learn whether or not we need to put energy in, via heat, to dissolve a substance.

In this lab, you'll explore the solubility of a compound at varying temperatures, and use titration to determine the exact concentration of the saturated solution. Then, you'll use your data to calculate the thermodynamic properties of the dissolution.