Chapter 13: Solutions – Study Notes for Introductory Chemistry

Solutions in Chemistry

Introduction to Solutions

Solutions are a fundamental concept in chemistry, describing homogeneous mixtures of two or more substances. Understanding solutions is essential for explaining many natural phenomena and industrial processes.

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

• Solvent: The majority component in a solution.

• Solute: The minority component in a solution.

• Aqueous solution: A solution in which water is the solvent.

• Example: The mixture of carbon dioxide and water at the bottom of Lake Nyos is a solution.

Tragedy in Cameroon: Chemistry in Context

Lake Nyos Disaster

The Lake Nyos tragedy illustrates the importance of understanding solutions and gas solubility. A sudden release of dissolved carbon dioxide from the lake led to a deadly event.

• Lake Nyos is a volcanic crater lake in Cameroon, West Africa.

• Molten volcanic rock beneath the lake produces carbon dioxide gas, which dissolves in the lake water.

• The concentration of dissolved gases increases with pressure; at the lake bottom, high pressure allows large amounts of CO2 to dissolve.

• A sudden drop in pressure (due to rising bubbles) caused CO2 to escape, displacing air and causing fatalities.

• Engineers now vent CO2 from the lake to prevent future disasters.

Types of Solutions

Common Types of Solutions

Solutions can exist in various phases, depending on the physical states of the solute and solvent.

Solution Phase	Solute Phase	Solvent Phase	Example
Gaseous solutions	Gas	Gas	Air (mainly oxygen and nitrogen)
Liquid solutions	Gas	Liquid	Soda water (CO2 and water)
Liquid solutions	Liquid	Liquid	Vodka (ethanol and water)
Solid solutions	Solid	Solid	Brass (copper and zinc)
Liquid solutions	Solid	Liquid	Seawater (salt and water)

Common Laboratory Solvents

Common Polar Solvents	Common Nonpolar Solvents
Water (H2O)	Hexane (C6H14)
Acetone (CH3COCH3)	Diethyl ether (CH3CH2OCH2CH3)
Methyl alcohol (CH3OH)	Toluene (C7H8)

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

How Solids Dissolve in Water

Solubility Process

When a solid is placed in water, the process of dissolution depends on the competition between solute-solute, solvent-solvent, and solvent-solute attractions.

• If solvent-solute attractions are stronger than solute-solute and solvent-solvent attractions, the solid dissolves.

• Example: NaCl dissolves in water because water molecules (dipoles) surround and separate Na+ and Cl- ions.

Solubility and Saturation

Definitions and Types

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

• Solubility: Amount of solute (usually in grams) that dissolves in a certain amount of liquid.

• Saturated solution: Contains the maximum amount of solute under given conditions.

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

• Supersaturated solution: Contains more than the normal maximum amount of dissolved solute; excess solute may precipitate.

• Example: The CO2-water mixture at Lake Nyos became supersaturated when pressure dropped, causing gas to escape.

Solubility Rules and Examples

Ionic and Molecular Solids

• Ionic solids: Solubility depends on the relative strength of ion-ion and solvent-ion attractions.

• Example: Calcium carbonate (CaCO3) is insoluble because Ca2+ and CO32- attraction is stronger than water's attraction.

• Molecular solids: Polar molecular solids (e.g., table sugar) are soluble in water; nonpolar solids (e.g., lard) are not.

Electrolyte and Nonelectrolyte Solutions

Electrical Conductivity

• Electrolyte solutions: Contain dissolved ions and conduct electricity (e.g., salt solution).

• Nonelectrolyte solutions: Contain dissolved molecules and do not conduct electricity (e.g., sugar solution).

Temperature Dependence of Solubility

Solids and Gases

• Solids: Solubility in water generally increases with increasing temperature.

• Example: Potassium nitrate solubility increases from 30 g/100 g water at 20°C to 80 g/100 g water at 50°C.

• Gases: Solubility in liquids decreases with increasing temperature but increases with increasing pressure.

Recrystallization and Purification

Recrystallization Technique

• Used to purify solids by dissolving them in hot solvent and allowing them to crystallize as the solution cools.

• Impurities are excluded from the crystalline structure.

• Example: Making rock candy by cooling a saturated sugar solution and allowing crystals to form on a string.

Solutions of Gas in Water

Gas-Liquid Solutions

• Many liquids contain dissolved gases (e.g., oxygen in lake water, CO2 in soda).

• Solubility of gases decreases with increasing temperature (e.g., warm soda fizzes more than cold soda).

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

Henry's Law: The higher the pressure of a gas above a liquid, the more soluble the gas is in the liquid.

Solution Concentration

Mass Percent

• Mass percent: Grams of solute per 100 g of solution.

$\text{mass percent} = \frac{\text{mass solute}}{\text{mass solute} + \text{mass solvent}} \times 100$

• Also expressed as parts per million (ppm) and parts per billion (ppb).

Molarity

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

$\text{molarity (M)} = \frac{\text{moles solute}}{\text{liters solution}}$

• Used to prepare solutions of specific concentration.

Molality

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

$\text{molality (m)} = \frac{\text{moles solute}}{\text{kilograms solvent}}$

Solution Dilution

Dilution Equation

• Used to prepare less concentrated solutions from stock solutions.

$M_1 V_1 = M_2 V_2$

• $M_1$ and $V_1$ are the molarity and volume of the concentrated solution; $M_2$ and $V_2$ are for the diluted solution.

Solution Stoichiometry

Using Concentration and Volume in Reactions

• Volume and concentration can be used to calculate moles of reactants or products.

• Stoichiometric coefficients from balanced equations are used to relate quantities.

Colligative Properties

Freezing Point Depression and Boiling Point Elevation

• Adding a nonvolatile solute lowers the freezing point and raises the boiling point of a solution.

• These effects depend only on the number of solute particles, not their identity.

Freezing Point Depression:

$\Delta T_f = m \times K_f$

Boiling Point Elevation:

$\Delta T_b = m \times K_b$

• $K_f$ and $K_b$ are constants specific to the solvent.

Osmosis and Osmotic Pressure

Osmosis

• Osmosis is the flow of solvent from a less concentrated solution to a more concentrated solution through a semipermeable membrane.

• Osmotic pressure is the pressure required to stop osmotic flow.

• Living cell membranes act as semipermeable membranes.

Medical Applications

• Hypoosmotic solutions: Lower osmotic pressure than bodily fluids; pump water into cells.

• Hyperosmotic solutions: Higher osmotic pressure than bodily fluids; draw water out of cells.

• Isoosmotic solutions: Equal osmotic pressure to bodily fluids; used in intravenous (IV) solutions.

Summary Table: Key Solution Concepts

Concept	Definition/Formula
Mass Percent	$\frac{\text{mass solute}}{\text{mass solute} + \text{mass solvent}} \times 100$
Molarity (M)	$\frac{\text{moles solute}}{\text{liters solution}}$
Molality (m)	$\frac{\text{moles solute}}{\text{kilograms solvent}}$
Dilution Equation	$M_1 V_1 = M_2 V_2$
Freezing Point Depression	$\Delta T_f = m \times K_f$
Boiling Point Elevation	$\Delta T_b = m \times K_b$

Learning Objectives

• Define solution, solute, and solvent.

• Relate solubility of solids in water to temperature.

• Relate solubility of gases in liquids to temperature and pressure.

• Calculate mass percent, molarity, and molality.

• Use dilution equation and stoichiometry in solution calculations.

• Explain colligative properties and osmosis.