Properties of Solutions – General Chemistry Study Notes

Chapter 13: Properties of Solutions

Introduction to Solutions

Solutions are homogeneous mixtures composed of two or more different chemical substances. The study of solutions is fundamental in chemistry, as it explains how substances mix, interact, and dissolve. Understanding the properties of solutions is essential for predicting solubility, reaction outcomes, and physical properties.

• Solution: A homogeneous mixture of two or more substances (ions, molecules, gases).

• Solvent: The substance present in the greatest amount; it dissolves the solute.

• Solute: The substance present in a smaller quantity; it is dissolved by the solvent.

• Types of Solutions:

  • Gas Solution: e.g., air (mixture of gases)

  • Liquid Solution: e.g., salt water (NaCl in H2O)

  • Solid Solution: e.g., alloys (22K gold: Au 91.67%, Ag 5%, Cu 2%)

• Homogeneous Mixture: Uniform composition throughout; does not settle over time.

Intermolecular Forces in Solutions

The formation and properties of solutions depend on the types of intermolecular forces present between solute and solvent particles. These forces determine solubility, boiling point, melting point, and viscosity.

• Dispersion Forces: Weak forces present in all molecules, especially nonpolar substances.

• Dipole-Dipole Interactions: Occur between polar molecules.

• Hydrogen Bonding: Special dipole-dipole interaction when H is bonded to O, N, or F.

• Ion-Dipole Interactions: Occur between ions and polar molecules (e.g., Na+ in water).

• Covalent Bonding: Strong bonds within molecules, not typically involved in solution formation.

Example: The principal interaction between Na+ and water is ion-dipole.

Properties Affected by Intermolecular Forces

• Boiling Point

• Melting Point

• Viscosity

• Color is not affected by intermolecular forces.

Hydrogen Bonding

Hydrogen bonding occurs when hydrogen is bonded to a highly electronegative atom (O, N, F) that has non-bonding electrons. This leads to strong intermolecular attractions.

• Electronegativity: O (3.5), N (3.0), F (4.0), Cl (3.0)

• Example: H2O exhibits hydrogen bonding; HCl does not, as Cl is less effective at hydrogen bonding than O, N, or F.

Spontaneity of Solution Formation

The mixing of substances to form solutions is often spontaneous, driven by enthalpy and entropy changes.

• Enthalpy Change ($\Delta H$): Energy change due to breaking and forming interactions.

• Entropy Change ($\Delta S$): Increase in disorder when substances mix.

• Gibbs Free Energy ($\Delta G$): Determines spontaneity: $\Delta G = \Delta H - T\Delta S$

Example: Mixing gases is spontaneous because $\Delta H \approx 0$ and $\Delta S$ increases.

Solvation and Hydration

Solvation is the process where solute particles are surrounded by solvent molecules. When water is the solvent, this process is called hydration.

• Solvent-Solvent Interactions: Must be overcome for solvation to occur.

• Solute-Solute Interactions: Must be broken up.

• Solvent-Solute Interactions: Must be favorable for solution formation.

Energy of Solution Formation

The enthalpy change of solution formation ($\Delta H_{soln}$) is the sum of three components:

• $\Delta H_{solute}$: Energy to separate solute particles (endothermic)

• $\Delta H_{solvent}$: Energy to separate solvent particles (endothermic)

• $\Delta H_{mix}$: Energy released when solute and solvent interact (exothermic)

Equation:

$\Delta H_{soln} = \Delta H_{solute} + \Delta H_{solvent} + \Delta H_{mix}$

Lattice Energy (LE): For ionic compounds, the energy required to separate ions in the solid lattice.

Example: For NaCl, LE = 788 kJ/mol.

Solubility and Saturation

Solubility is the maximum amount of solute that can dissolve in a solvent at a given temperature and pressure. The process is dynamic, involving both dissolution and crystallization.

• Unsaturated Solution: Contains less solute than the maximum possible at equilibrium.

• Saturated Solution: Contains the maximum amount of solute at equilibrium; excess solute remains undissolved.

• Supersaturated Solution: Contains more solute than the equilibrium maximum; unstable and can precipitate excess solute.

Example: NaCl in water: If more than the maximum dissolves, the solution is supersaturated.

Factors Affecting Solubility

Several factors influence the solubility of substances:

• Solute-Solvent Interactions: Similar intermolecular forces increase solubility ("like dissolves like").

• Pressure Effects: Significant for gases; increased pressure increases gas solubility in liquids.

• Temperature Effects: Generally, solubility of solids increases with temperature; gas solubility decreases.

Liquid-Liquid Interactions

• Miscible: Liquids that mix in all proportions (e.g., ethanol and water).

• Immiscible: Liquids that do not mix, forming separate layers (e.g., oil and water).

Pressure Effects and Henry's Law

Pressure has little effect on the solubility of solids and liquids, but greatly affects gases. Henry's Law quantifies this relationship.

• Henry's Law: The solubility of a gas in a liquid is proportional to the partial pressure of the gas above the solution.

Equation:

$S_g = k P_g$

• $S_g$ = solubility of the gas (mol/L)

• $k$ = Henry's Law constant (depends on gas, solvent, and temperature)

• $P_g$ = partial pressure of the gas (atm)

Example: The solubility of N2 in water at 25°C and 0.78 atm is $5.3 \times 10^{-4}$ M. To find $k$:

$k = \frac{S_g}{P_g} = \frac{5.3 \times 10^{-4}}{0.78}$

Solubility of Gases in Water

Generally, the solubility of gases in water increases with increasing molar mass.

Gas	Molar Mass (g/mol)	Solubility (M) at 20°C, 1 atm
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

Additional info: Table values inferred from context and standard chemistry data.

Physical vs. Chemical Change in Solution Formation

Dissolving is a physical change; the original solute can be recovered by evaporating the solvent. If a chemical reaction occurs, the solute cannot be recovered by simple evaporation.

• Physical Change: Dissolution (e.g., salt in water)

• Chemical Change: Reaction (e.g., acid-base reaction)

Summary Table: Types of Intermolecular Forces

Type of Force	Example	Relative Strength
Dispersion	He-He	Weak
Dipole-Dipole	HCl-HCl	Moderate
Hydrogen Bonding	H2O-H2O	Strong
Ion-Dipole	Na+-H2O	Very Strong

Additional info: Table constructed for clarity and completeness.