Properties of Solutions – GOB Chemistry Study Notes

Properties of Solutions

Introduction to Solutions

Solutions are homogeneous mixtures composed of two or more pure substances. In a solution, the solute is the substance present in the smaller amount, while the solvent is present in the larger amount. The solute is uniformly dispersed throughout the solvent, resulting in a single phase.

• Homogeneous mixture: Uniform composition throughout.

• Solute: Substance dissolved (smaller amount).

• Solvent: Substance doing the dissolving (larger amount).

• Example: In a soft drink, sugar and CO2 are solutes, water is the solvent.

Solutions can exist in different states depending on the physical states of the solute and solvent.

State of Solution	State of Solvent	State of Solute	Example
Gas	Gas	Gas	Air
Liquid	Liquid	Gas	Oxygen in water
Liquid	Liquid	Liquid	Alcohol in water
Liquid	Liquid	Solid	Salt in water
Solid	Solid	Gas	Hydrogen in palladium
Solid	Solid	Liquid	Mercury in silver
Solid	Solid	Solid	Silver in gold

The Solution Process

The formation of solutions is driven by two main factors: the natural tendency toward mixing (increasing entropy) and the effect of intermolecular forces. Mixing increases randomness (entropy), which is thermodynamically favored.

• Entropy: A measure of disorder or randomness. Mixing increases entropy.

• Intermolecular forces: Attractions between solute and solvent molecules must be strong enough to allow mixing.

Types of intermolecular forces include dispersion, dipole-dipole, hydrogen bonding, and ion-dipole interactions.

Saturated, Unsaturated, and Supersaturated Solutions

Solubility is the maximum amount of solute that can dissolve in a given amount of solvent at a specific temperature. Solutions can be classified as:

• Saturated: Contains the maximum amount of dissolved solute at equilibrium. Additional solute will not dissolve.

• Unsaturated: Contains less solute than the maximum amount; more solute can dissolve.

• Supersaturated: Contains more solute than is normally possible at that temperature; unstable and can crystallize easily.

Dynamic equilibrium exists in saturated solutions between dissolved and undissolved solute.

Factors Affecting Solubility

Solubility depends on the nature of the solute and solvent, temperature, and pressure (for gases). The principle "like dissolves like" applies: polar solutes dissolve in polar solvents, nonpolar in nonpolar.

• Solute-solvent interactions: The more similar the intermolecular attractions, the greater the solubility.

• Examples: Glucose (hydrogen bonding) is very soluble in water; cyclohexane (dispersion forces only) is not.

Vitamins illustrate solubility differences: Vitamin A (nonpolar) dissolves in fats, Vitamin C (polar) dissolves in water.

Solubility of Gases

• Solubility of gases in water increases with molecular mass due to stronger dispersion forces.

• Solubility of gases decreases with increasing temperature.

• Solubility of gases increases with increasing pressure (Henry's Law).

Gas	Molar Mass (g/mol)	Solubility (M)
N2	28.0	0.69 × 10-3
O2	32.0	1.38 × 10-3
Ar	39.9	1.50 × 10-3
Kr	83.8	2.79 × 10-3

Henry's Law

The solubility of a gas in a liquid is directly proportional to the partial pressure of the gas above the solution:

$S_g = k P_g$

• Sg: Solubility of the gas (usually in molarity)

• k: Henry's Law constant (depends on gas, solvent, temperature)

• Pg: Partial pressure of the gas

Applications include carbonated beverages and scuba diving safety.

Temperature Effects

• Solubility of most solid solutes in liquid solvents increases with temperature.

• Solubility of gases in liquids decreases with increasing temperature.

Concentration Units

Concentration expresses the amount of solute in a given quantity of solvent or solution. Common units include:

• Molarity (M): Moles of solute per liter of solution.

$M = \frac{\text{moles of solute}}{\text{liters of solution}}$

• Molality (m): Moles of solute per kilogram of solvent.

$m = \frac{\text{moles of solute}}{\text{kilograms of solvent}}$

• Mass percent (%): $\frac{\text{mass of solute}}{\text{total mass of solution}} \times 100$

• Parts per million (ppm): $\frac{\text{mass of solute}}{\text{total mass of solution}} \times 10^6$

• Parts per billion (ppb): $\frac{\text{mass of solute}}{\text{total mass of solution}} \times 10^9$

• Mole fraction (χ): $\chi_A = \frac{\text{moles of A}}{\text{total moles of all components}}$

Preparation of Standard Solutions

• Calculate required mass of solute using molarity equation.

• Weigh solute, dissolve in water, transfer to volumetric flask, dilute to mark, mix thoroughly.

Dilution

To prepare a dilute solution from a concentrated stock:

$M_1 V_1 = M_2 V_2$

• M1, V1: Initial concentration and volume

• M2, V2: Final concentration and volume

Colligative Properties

Colligative properties depend only on the number of solute particles, not their identity. These include:

• Vapor-pressure lowering

• Boiling-point elevation

• Freezing-point depression

• Osmosis

Vapor Pressure Lowering (Raoult's Law)

The vapor pressure of a solution is lower than that of the pure solvent. Raoult's Law quantifies this:

$P_A = X_A P_A^\circ$

• PA: Vapor pressure of solvent A in solution

• XA: Mole fraction of solvent A

• PA∘: Vapor pressure of pure solvent A

Boiling-Point Elevation

The boiling point of a solution is higher than that of the pure solvent:

$\Delta T_b = i K_b m$

• ΔTb: Boiling point elevation

• Kb: Molal boiling-point elevation constant

• m: Molality of solution

• i: van 't Hoff factor (number of particles per formula unit)

Freezing-Point Depression

The freezing point of a solution is lower than that of the pure solvent:

$\Delta T_f = i K_f m$

• ΔTf: Freezing point depression

• Kf: Molal freezing-point depression constant

• m: Molality of solution

• i: van 't Hoff factor

Osmosis and Osmotic Pressure

Osmosis is the movement of solvent through a semipermeable membrane from a region of lower solute concentration to higher solute concentration. The pressure required to stop osmosis is the osmotic pressure (π):

$\pi = MRT$

• π: Osmotic pressure

• M: Molarity of solution

• R: Gas constant

• T: Temperature in Kelvin

In biological systems, osmotic pressure differences can cause cells to shrink (crenation) or burst (hemolysis).

Determination of Molar Mass Using Colligative Properties

Colligative properties such as osmotic pressure, freezing-point depression, and boiling-point elevation can be used to determine the molar mass of an unknown solute. By measuring the effect and knowing the amount of solute and solvent, the molar mass can be calculated using the relevant equations above.