Chapter 7: Chemical Reactions & Quantities

The Mole

The mole is a fundamental unit in chemistry used to count particles such as atoms, molecules, or ions. One mole contains Avogadro's number of particles, which is .

• Definition: A mole is the amount of substance containing as many entities as there are atoms in 12 grams of carbon-12.

• Avogadro's Number: particles/mole

• Example: 1 mole of water contains molecules of water.

Molar Mass Concept & Calculations

Molar mass is the mass of one mole of a substance, expressed in grams per mole (g/mol).

• Calculation: Add the atomic masses of all atoms in a formula.

• Formula:

• Example: Molar mass of H2O: g/mol

Writing Chemical Equations

Chemical equations represent chemical reactions using symbols and formulas.

• Reactants: Substances consumed in the reaction.

• Products: Substances formed in the reaction.

• Example:

Identifying Types of Chemical Reactions

Chemical reactions are classified based on the changes occurring.

• Combination (Synthesis): Two or more substances form one product.

• Decomposition: One substance breaks into two or more products.

• Single Replacement: One element replaces another in a compound.

• Double Replacement: Two compounds exchange ions.

• Combustion: A substance reacts with oxygen, producing energy.

• Example: (combustion)

Mole Relationships in Chemical Reactions

Stoichiometry uses balanced equations to relate moles of reactants and products.

• Coefficients: Indicate mole ratios.

• Example: In , 2 moles of H2 react with 1 mole of O2 to produce 2 moles of H2O.

Mass Relationships in Chemical Reactions

Mass relationships are determined using molar masses and stoichiometry.

• Steps: Convert mass to moles, use mole ratios, convert moles to mass.

• Formula:

• Example: How many grams of H2O are produced from 4 g of H2?

Energy in Chemical Reactions (Endothermic vs. Exothermic)

Energy changes accompany chemical reactions.

• Exothermic: Releases energy (heat).

• Endothermic: Absorbs energy (heat).

• Example: Combustion reactions are exothermic.

Chapter 8: Gases, Liquids, and Solids

Properties of Gases (Kinetic Molecular Theory)

The kinetic molecular theory explains the behavior of gases based on particle motion.

• Assumptions: Gas particles are in constant, random motion; collisions are elastic; volume of particles is negligible.

• Properties: Gases have low density, are compressible, and fill their containers.

• Example: Air in a balloon expands to fill the shape.

Boyle’s Law

Boyle’s Law describes the relationship between pressure and volume at constant temperature.

• Formula:

• Inverse Relationship: As pressure increases, volume decreases.

• Example: Compressing a syringe increases pressure.

Charles’s Law

Charles’s Law relates volume and temperature at constant pressure.

• Formula:

• Direct Relationship: As temperature increases, volume increases.

• Example: A hot air balloon expands as it is heated.

Gay-Lussac’s Law

Gay-Lussac’s Law relates pressure and temperature at constant volume.

• Formula:

• Direct Relationship: As temperature increases, pressure increases.

• Example: Pressure in a tire increases on a hot day.

Combined Gas Law

The combined gas law combines Boyle’s, Charles’s, and Gay-Lussac’s laws.

• Formula:

• Application: Used when more than one variable changes.

Avogadro’s Law

Avogadro’s Law relates volume and number of moles at constant temperature and pressure.

• Formula:

• Direct Relationship: More moles, greater volume.

• Example: Adding gas to a balloon increases its size.

Dalton’s Law (Partial Pressures)

Dalton’s Law states that the total pressure of a gas mixture is the sum of the partial pressures of each component.

• Formula:

• Example: Air is a mixture of N2, O2, CO2, etc.

Chapter 9: Solutions

Solutions: Solute & Solvent

A solution is a homogeneous mixture of two or more substances.

• Solute: The substance dissolved.

• Solvent: The substance doing the dissolving (usually present in greater amount).

• Example: Salt (solute) dissolved in water (solvent).

Electrolytes vs. Nonelectrolytes

Electrolytes conduct electricity in solution; nonelectrolytes do not.

• Electrolytes: Ionic compounds that dissociate in water.

• Nonelectrolytes: Molecular compounds that do not dissociate.

• Example: NaCl is an electrolyte; sugar is a nonelectrolyte.

Saturated vs. Unsaturated Solutions

A saturated solution contains the maximum amount of solute; an unsaturated solution can dissolve more solute.

• Saturated: No more solute dissolves.

• Unsaturated: More solute can dissolve.

• Example: Adding sugar to tea until no more dissolves.

Solubility vs. Temperature

Solubility is the amount of solute that can dissolve in a solvent at a given temperature.

• Solids: Solubility increases with temperature.

• Gases: Solubility decreases with temperature.

• Example: More sugar dissolves in hot water than cold.

Mass Percent

Mass percent expresses concentration as grams of solute per 100 grams of solution.

• Formula:

• Example: 5 g NaCl in 100 g solution is 5% mass percent.

Volume Percent

Volume percent is used for liquid solutions.

• Formula:

• Example: 10 mL ethanol in 100 mL solution is 10% volume percent.

Mass/Volume Percent

Mass/volume percent expresses concentration as grams of solute per 100 mL of solution.

• Formula:

• Example: 1 g NaCl in 100 mL solution is 1% mass/volume percent.

Molarity

Molarity (M) is the number of moles of solute per liter of solution.

• Formula:

• Example: 0.5 moles NaCl in 1 L solution is 0.5 M.

Solution vs. Colloid vs. Suspension

Mixtures are classified by particle size and behavior.

Hypertonic vs. Hypotonic vs. Isotonic

These terms describe the relative concentration of solutes in solutions separated by a membrane.

• Hypertonic: Higher solute concentration outside the cell; water leaves the cell.

• Hypotonic: Lower solute concentration outside the cell; water enters the cell.

• Isotonic: Equal solute concentration; no net water movement.

• Example: IV solutions are isotonic to prevent cell damage.