Chapter 9: Solutions – Fundamentals of General, Organic, and Biological Chemistry

Chapter 9: Solutions

9.1 Mixtures and Solutions

Mixtures are combinations of two or more substances that are not chemically bonded. They can be classified based on the uniformity of their composition and the size of their particles.

• Heterogeneous mixtures: Mixtures in which the composition is not uniform throughout. Different regions have different compositions. Example: Rocky Road ice cream, chocolate chip cookies, pot pie.

• Homogeneous mixtures: Mixtures with uniform composition throughout. Example: Seawater, salt water, coffee, air.

• Homogeneous mixtures can be further classified as:

  • Solutions: Homogeneous mixtures with particle sizes less than 2.0 nm (nanometers). Example: Air, seawater, gasoline, wine.

  • Colloids: Homogeneous mixtures with particle sizes between 2.0 and 500 nm. Example: Butter, milk, fog, pearl.

Type of Mixture	Particle Size	Examples	Characteristics
Solution	<2.0 nm	Air, seawater, gasoline, wine	Transparent to light, does not separate on standing, nonfilterable
Colloid	2.0–500 nm	Butter, milk, fog, pearl	Often murky or opaque to light; does not separate on standing; nonfilterable
Heterogeneous	>500 nm	Blood, paint, aerosol sprays	Murky or opaque to light; separates on standing; filterable

• Parts of a solution:

  • Solute: The substance that is dissolved.

  • Solvent: The substance in which the solute is dissolved. In liquid solutions, the major component is usually the solvent.

Type of Solution	Example
Gas in gas	Air (O2, N2, Ar, etc.)
Gas in liquid	Seltzer water (CO2 in water)
Gas in solid	H2 in palladium metal
Liquid in liquid	Gasoline (mixture of hydrocarbons)
Liquid in solid	Dental amalgam (mercury in silver)
Solid in liquid	Seawater (NaCl and other salts in water)
Solid in solid	Metal alloys such as 14-karat gold (Au, Ag, Cu)

9.2 The Solution Process

The ability of a solute to dissolve in a solvent (solubility) depends on the relative strength of the attractions between solute and solvent particles compared to the attractions within the pure substances.

• "Like dissolves like": Polar solvents dissolve polar and ionic solutes; nonpolar solvents dissolve nonpolar solutes.

• Example: Water (polar) dissolves NaCl (ionic), but oil (nonpolar) does not mix with water.

• Solvation (hydration): When ionic compounds dissolve in water, water molecules surround the ions, stabilizing them by electrical attraction.

• Dissolution is a physical change—the chemical identities of the solute and solvent are retained.

• Enthalpy change (ΔH): Dissolution can be exothermic (releases heat) or endothermic (absorbs heat).

Worked Example 9.1

• (a) CCl4 and hexane: Both are nonpolar; will form a solution.

• (b) Octane and methyl alcohol: Octane is nonpolar, methyl alcohol is polar; will not form a solution.

9.3 Solubility

Solubility describes how much solute can dissolve in a given amount of solvent at a specific temperature.

• Miscible: Two substances that are mutually soluble in all proportions (e.g., ethyl alcohol and water).

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

• Solubility: The maximum amount of a substance that will dissolve in a given amount of solvent at a specified temperature.

9.4 The Effect of Temperature on Solubility

Temperature significantly affects solubility, but the effect depends on the state of the solute.

• Solids: Most become more soluble as temperature increases.

• Gases: Become less soluble as temperature increases.

• Supersaturated solution: Contains more solute than a saturated solution at a given temperature; unstable and will precipitate if disturbed.

Worked Example 9.2

• Solubility of O2 at 25°C: 8.3 mg/L; at 35°C: 7.0 mg/L.

• Percent change: $\frac{8.3 - 7.0}{8.3} \times 100 = 16$

9.5 The Effect of Pressure on Solubility: Henry's Law

Henry's law states that the solubility of a gas in a liquid is directly proportional to the partial pressure of the gas above the liquid, provided temperature is constant.

• Equation: $C = kP_{gas}$, where C is the concentration (solubility) of the gas, k is a constant, and $P_{gas}$ is the partial pressure of the gas.

• Increasing pressure increases gas solubility; decreasing pressure decreases it.

9.6 Units of Concentration

Concentration expresses the amount of solute in a given amount of solution. Several units are commonly used:

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

• Volume/volume percent (v/v)%: $\frac{\text{volume of solute (mL)}}{\text{volume of solution (mL)}} \times 100$

• Mass/volume percent (m/v)%: $\frac{\text{mass of solute (g)}}{\text{volume of solution (mL)}} \times 100$

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

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

• Molarity (M): $\frac{\text{moles of solute}}{\text{liters of solution}}$

9.7 Dilution

Solutions are often prepared by diluting a more concentrated solution. The amount of solute remains constant; only the volume changes.

• Dilution equation: $M_1V_1 = M_2V_2$

• Where $M_1$ and $V_1$ are the initial molarity and volume, and $M_2$ and $V_2$ are the final molarity and volume.

9.8 Ions in Solution: Electrolytes

When ionic compounds dissolve in water, they dissociate into ions, which can conduct electricity.

• Strong electrolytes: Ionize completely in water (e.g., NaCl).

• Weak electrolytes: Partially ionize in water (e.g., acetic acid).

• Nonelectrolytes: Do not produce ions in water (e.g., sugar).

• Equivalent (Eq): The amount of an ion equal to 1 mole of charge. $1\,\text{Eq} = 1000\,\text{mEq}$

9.9 Properties of Solutions

Colligative properties depend on the number of dissolved solute particles, not their identity.

• Vapor pressure lowering: Solutions have lower vapor pressure than pure solvents.

• Boiling point elevation: Solutions boil at higher temperatures than pure solvents. $\Delta T_b = 0.51^\circ\text{C} \times \text{mol solute particles per kg water}$

• Freezing point depression: Solutions freeze at lower temperatures than pure solvents. $\Delta T_f = 1.86^\circ\text{C} \times \text{mol solute particles per kg water}$

• Osmosis: The movement of solvent through a semipermeable membrane from a region of lower solute concentration to higher solute concentration.

9.10 Osmosis and Osmotic Pressure

Osmosis is vital in biological systems, as cell membranes are semipermeable.

• Osmotic pressure (π): The pressure required to prevent osmosis. $\pi = nRT$ (for ideal solutions)

• Osmolarity (osmol/L): The total molarity of all particles in solution.

• Isotonic: Same osmolarity as body fluids.

• Hypotonic: Lower osmolarity than body fluids.

• Hypertonic: Higher osmolarity than body fluids.

9.11 Dialysis

Dialysis is the process by which small solute particles and solvent molecules pass through a semipermeable membrane, but larger particles (like proteins) do not.

• Hemodialysis: A medical procedure that uses dialysis to remove waste products from the blood when the kidneys are not functioning properly.

• Proteins do not cross the membrane, helping to maintain osmotic balance in blood plasma.