Unit 8: Solubility & Saturated Solutions
Exploring the fundamental concepts of how substances dissolve and the limits of solubility.
8.1 Definitions (Solute, Solvent, Solution, Solubility)
To understand solubility, we must first define the components of a mixture.
- Solvent: The substance that does the dissolving. It is typically the component present in the largest amount. The most common solvent in chemistry is water ($H_2O$).
- Solute: The substance that is being dissolved in the solvent.
- Solution: A homogeneous mixture formed when a solute dissolves in a solvent. Solutions where water is the solvent are called aqueous solutions.
Solubility is a quantitative measure of the maximum amount of a solute that can dissolve in a given amount of solvent at a specific temperature to form a stable solution. It is typically expressed as concentration, most commonly in moles per cubic decimeter ($mol/dm^3$), also known as molarity (M).
Solubility is temperature-dependent. For most solids, solubility increases as temperature increases. For gases, solubility decreases as temperature increases.
Solved Examples:
-
In a sugar-water solution, identify the solute and the solvent.
Solution: The sugar is the solute, and the water is the solvent. -
The solubility of potassium nitrate ($KNO_3$) at 20 °C is 3.16 mol/dm³. What
does this statement mean?
Solution: It means that at 20 °C, a maximum of 3.16 moles of potassium nitrate can be dissolved in 1 cubic decimeter (1 liter) of water to form a stable solution.
8.2 Saturated Solutions & Dynamic Equilibrium
A saturated solution is a solution that contains the maximum possible amount of dissolved solute at a given temperature. If any more solute is added to a saturated solution, it will not dissolve and will remain as a solid at the bottom of the container. A solution that contains less than the maximum amount of solute is called an unsaturated solution.
In a saturated solution with excess undissolved solid, a state of dynamic equilibrium exists. This means that the process of dissolving and the process of crystallization (the solute coming out of solution) are occurring at the exact same rate.
$$ \text{Solute (solid)} \rightleftharpoons \text{Solute (dissolved)} $$
Although the overall concentration of the solution remains constant, at the microscopic level, particles are continuously moving from the solid state to the dissolved state and back again.
Solved Examples:
-
At 25 °C, the solubility of NaCl is 6.2 mol/dm³. A student prepares a
solution with a concentration of 5.0 mol/dm³. Is this solution saturated or
unsaturated?
Solution: The solution is unsaturated because its concentration (5.0 mol/dm³) is less than the maximum possible solubility (6.2 mol/dm³). More NaCl could still be dissolved. -
What would you observe if you added a small crystal of a solute to its
supersaturated solution?
Solution: A supersaturated solution is an unstable state containing more dissolved solute than a saturated solution. Adding a "seed" crystal provides a surface for crystallization to begin, and you would observe a rapid crystallization of the excess solute from the solution. -
Explain the concept of dynamic equilibrium in a saturated salt
solution.
Solution: In a saturated solution with undissolved salt, ions from the solid crystal are constantly dissolving into the water, while an equal number of dissolved ions are precipitating back onto the crystal. The rates of these two opposing processes are equal, so there is no net change in concentration.
8.3 Electrolytes (Strong, Weak, Non-electrolytes)
An electrolyte is a substance that produces an electrically conducting solution when dissolved in a polar solvent, such as water. The conductivity is due to the presence of mobile ions.
- Strong Electrolytes: These are solutes that dissociate completely into
their constituent ions when dissolved in water. All soluble ionic compounds (like NaCl)
and strong acids (like HCl) are strong electrolytes.
$NaCl(s) \rightarrow Na^+(aq) + Cl^-(aq)$ (100% dissociation) - Weak Electrolytes: These solutes only partially dissociate into ions in
solution. An equilibrium is established between the undissociated molecules and the
ions. Weak acids (like acetic acid) and weak bases (like ammonia) are weak electrolytes.
$NH_3(g) + H_2O(l) \rightleftharpoons NH_4^+(aq) + OH^-(aq)$ (Partial dissociation) - Non-electrolytes: These substances dissolve in water but do not produce
any ions. The solution contains only neutral molecules and does not conduct electricity.
Most molecular compounds, like sugar (sucrose) and ethanol, are non-electrolytes.
$C_{12}H_{22}O_{11}(s) \rightarrow C_{12}H_{22}O_{11}(aq)$ (No dissociation)
Solved Examples:
-
Is potassium chloride (KCl) a strong electrolyte, weak electrolyte, or
non-electrolyte? Explain.
Solution: KCl is a soluble ionic compound. Therefore, it is a strong electrolyte because it dissociates completely into K⁺(aq) and Cl⁻(aq) ions in water. -
A solution of substance X barely conducts electricity. How would you
classify substance X?
Solution: Since it conducts electricity only slightly, it must be producing a low concentration of ions. Therefore, substance X is a weak electrolyte.
8.4 Measuring Solubility (Acids & Bases by Titration)
The solubility of a substance can be determined experimentally. For soluble or sparingly soluble acids and bases, a common method is to prepare a saturated solution and then determine its concentration using an acid-base titration.
The general procedure is:
- Prepare a saturated solution of the substance by adding an excess of the solid to water and stirring until equilibrium is reached.
- Filter the solution to remove any undissolved solid. The clear liquid (filtrate) is the saturated solution.
- Accurately measure a known volume of the saturated solution (the analyte).
- Titrate this solution with a standard solution of a strong acid or base (the titrant) using a suitable indicator.
- Use the titration results ($M_aV_a = n_a M_bV_b$) to calculate the concentration of the saturated solution. This concentration is the solubility of the substance at that temperature.
Solved Examples:
-
Why must the saturated solution be filtered before titration?
Solution: Filtering removes the excess, undissolved solid. If the solid were present during the titration, it would continue to dissolve as the dissolved substance is neutralized by the titrant, leading to an inaccurate (too high) measurement of the solubility. -
A student titrates 25.0 cm³ of a saturated calcium hydroxide solution with
0.050 mol/dm³ HCl. The average titre is 20.0 cm³. Calculate the solubility of
calcium hydroxide in mol/dm³.
Solution:
Equation: $Ca(OH)_2 + 2HCl \rightarrow CaCl_2 + 2H_2O$. The mole ratio is 1:2.
Moles of HCl = $0.050 \times (20.0/1000) = 0.001$ mol.
Moles of Ca(OH)₂ = $0.001 / 2 = 0.0005$ mol.
Concentration (Solubility) = Moles / Volume = $0.0005 / (25.0/1000) = 0.020$ mol/dm³. -
What is a "standard solution" in the context of a titration?
Solution: A standard solution is a solution whose concentration is known accurately. It is used as the titrant to determine the unknown concentration of the analyte.