Chapter 9: Solutions

9.1 Mixtures and Solutions

This section introduces the classification of mixtures and the distinction between solutions, colloids, and heterogeneous mixtures.

• Heterogeneous Mixtures: Non-uniform composition; different regions have different properties. Example: Rocky Road ice cream.

• Homogeneous Mixtures: Uniform composition throughout. Example: Seawater.

• Solutions: Homogeneous mixtures with particles the size of ions or small molecules.

• Colloids: Homogeneous mixtures with particle diameters from 2 to 500 nm.

• Solute: The substance dissolved in a solvent.

• Solvent: The substance in which the solute is dissolved; usually the major component.

Additional info: Table 9.1 and 9.2 summarize characteristics and types of solutions.

9.2 The Solution Process

The formation of solutions depends on the interactions between solute and solvent particles, primarily their polarity and intermolecular forces.

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

• Solvation/Hydration: Water molecules surround ions, stabilizing them by electrical attraction.

• Physical Change: Dissolution retains chemical identities of components.

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

Example: Hexane and carbon tetrachloride (both nonpolar) form a solution; octane (nonpolar) and methyl alcohol (polar) do not.

9.3 Solubility

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

• Miscible: Substances mutually soluble in all proportions (e.g., ethyl alcohol and water).

• Saturated Solution: Contains maximum dissolved solute at equilibrium.

• Solubility: Maximum amount of solute that dissolves in a given solvent at a specified temperature.

Example: At 20°C, 35.8 g NaCl dissolves in 100 mL water; 204 g sucrose dissolves under same conditions.

9.4 Effect of Temperature on Solubility

Temperature affects solubility differently for solids and gases.

• Solids: Generally more soluble as temperature increases.

• Gases: Less soluble as temperature increases.

• Supersaturated Solution: Contains more solute than a saturated solution; unstable and precipitates when disturbed.

Example: Solubility of oxygen in water decreases from 8.3 mg/L at 25°C to 7.0 mg/L at 35°C.

9.5 Effect of Pressure on Solubility: Henry’s Law

Pressure influences the solubility of gases in liquids, described by Henry’s Law.

• Henry’s Law: Solubility of a gas is directly proportional to its partial pressure at constant temperature.

• Le Châtelier’s Principle: Increasing pressure increases gas solubility; decreasing pressure decreases it.

Equation:

Where C is concentration, k is a constant, and P is partial pressure.

Example: Solubility of oxygen in blood decreases with altitude as partial pressure drops.

9.6 Units of Concentration

Concentration expresses the amount of solute in a given amount of solution, using various units.

• Mass/Mass Percent (m/m)%:

• Volume/Volume Percent (v/v)%:

• Mass/Volume Percent (m/v)%:

• Parts per Million (ppm):

• Parts per Billion (ppb):

• Molarity (M):

Example: Calculating mass of gold in an 18-carat ring, volume of methyl alcohol for a solution, or molarity of sulfuric acid in water.

9.7 Dilution

Dilution involves adding solvent to a concentrated solution to achieve a desired concentration.

• Key Principle: Amount of solute remains constant; only volume changes.

• Dilution Equation:

• Dilution Factor: Ratio of initial to final volumes.

Example: Diluting 75 mL of 3.5 M glucose to 450 mL; preparing 750 mL of 0.32 M NaOH from 1.0 M stock.

9.8 Ions in Solution: Electrolytes

Electrolytes are substances that produce ions in solution and conduct electricity.

• Strong Electrolytes: Ionize completely in water.

• Weak Electrolytes: Partially ionize in water.

• Nonelectrolytes: Do not produce ions in water.

• Equivalent (Eq): Amount of ion equal to 1 mol of charge; 1 mEq = 0.001 Eq.

Example: Calculating milligrams of potassium ion in blood using mEq/L.

9.9 Properties of Solutions

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

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

• Boiling Point Elevation: Solutions boil at higher temperatures. (where is boiling point elevation constant, is molality)

• Freezing Point Depression: Solutions freeze at lower temperatures. (where is freezing point depression constant)

• Osmosis: Solvent passes through a semipermeable membrane toward higher solute concentration.

Example: Boiling point of 0.75 mol KBr in 1 kg water increases by 0.77°C; freezing point depression for tomato cells calculated from solute concentration.

9.10 Osmosis and Osmotic Pressure

Osmosis is the movement of solvent across a semipermeable membrane, and osmotic pressure is the pressure required to prevent this movement.

• Osmotic Pressure (π): (where M is molarity, R is gas constant, T is temperature in Kelvin)

• Osmolarity: Sum of molarities of all dissolved particles per liter.

• Isotonic: Same osmolarity as blood plasma.

• Hypotonic: Lower osmolarity than plasma.

• Hypertonic: Higher osmolarity than plasma.

Example: Calculating osmolarity of glucose solution; mass of NaCl needed for a specific osmolarity.

9.11 Dialysis

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

• Hemodialysis: Used to cleanse blood in patients with kidney malfunction.

• Colloidal Proteins: Do not cross membranes; regulate osmolarity and water movement in blood.

• Filtration vs. Reabsorption: Blood pressure pushes water out; osmotic pressure of proteins draws water in.

Example: Hemodialysis removes waste like urea from blood; proteins maintain osmotic balance.

Additional info: Academic context was added to clarify equations, units, and physiological relevance. Tables were reconstructed for clarity.