Chapter 11: Solutions, Solubility, and Units of Concentration

Introduction

This chapter explores how liquids change when other compounds are mixed with them, focusing on the formation of solutions, the principles of solubility, and the quantitative units used to describe solution concentration. Understanding these concepts is fundamental in both chemistry and biology, as many biological processes occur in aqueous solutions.

11.1 Solutions

Definition and Components

• Solution: A homogeneous mixture composed of two or more substances.

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

• Solute: The substance dissolved in the solvent.

Example: In a saltwater solution, water is the solvent and salt (NaCl) is the solute.

11.1 Solubility

Solubility Rules and Dissociation

• Solubility: The ability of a solute to dissolve in a solvent at a specific temperature and pressure.

• Dissociation: The process by which ionic compounds separate into ions in solution.

• Like dissolves like: Polar solvents dissolve polar solutes; non-polar solvents dissolve non-polar solutes.

• Covalent compounds generally do not dissociate in solution.

Example: Table sugar (sucrose) dissolves in water without dissociation.

11.1 Covalent Solubility

Effect of Polarity on Solubility

• Polarity: Molecules with uneven distribution of charge (polar) interact well with other polar molecules.

• Polar compounds dissolve in polar solutions (e.g., water, ethanol, methanol).

• Non-polar compounds dissolve in non-polar solutions (e.g., methane, pentane, hydrocarbons).

Example Table:

11.3 Solutions: Saturation and Solubility

Types of Solutions Based on Concentration

• Saturated solution: Contains the maximum amount of solute that can dissolve at a given temperature.

• Unsaturated solution: Contains less solute than the maximum amount that can dissolve.

• Supersaturated solution: Contains more solute than is normally possible at a given temperature; usually unstable.

Example: If a solution at 340 K has a concentration of 0.52 M and is cooled to 290 K without precipitation, it becomes supersaturated.

11.3 Solubility as an Equilibrium Process

Dynamic Equilibrium in Solutions

• Solubility is an equilibrium process where the rate of dissolving equals the rate of precipitation.

• Forward process: Undissolved solute dissolves.

• Reverse process: Dissolved solute precipitates out.

• At equilibrium, the amount of dissolved and undissolved solute remains constant.

11.3 Solid Solubility

Effect of Temperature and Pressure

• Increasing temperature usually increases the solubility of solids in liquids, but not always.

• Pressure has no effect on the solubility of solids in liquids.

11.3 Gas Solubility

Effect of Temperature and Pressure

• Solubility of gases in water decreases as temperature increases.

• Solubility of gases in water increases as pressure increases (Henry's Law).

Example: Carbonated beverages are bottled under high pressure to increase the solubility of CO2 gas.

Summary of Solubility

11.4 Units of Concentration

Quantitative Measures of Solution Concentration

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

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

• Mole fraction (X): Ratio of moles of a component to the total moles in the solution.

Key Equations:

• Molarity:

• Molality:

• Mole fraction (for solvent):

Additional info: Molecular mass (MM) is always expressed in g/mol. To convert between grams and moles, use:

• Grams to moles:

• Moles to grams:

Examples

• Given a solute dissolved in 755 mL of water (density = 1.00 g/mL), calculate M, m, and Xsolvent.

• Calculate the molality (m) for a 0.125 M solution of KBr, given the density of the solution is 1.13 g/mL and the molar mass of KBr is 119.00 g/mol.

Additional info: To solve such problems, convert all volumes to liters or kilograms as appropriate, and use the provided densities to find mass if needed.