Chapter 7: Chemical Reactions and Chemical Quantities

Balancing Chemical Equations

Balancing chemical equations ensures the law of conservation of mass is obeyed. Each side of the equation must have the same number of atoms for each element.

• Coefficients are used to balance equations, not subscripts.

• Balance elements that appear in only one reactant and one product first.

• Balance polyatomic ions as units if they appear unchanged on both sides.

Types of Chemical Reactions

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

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

• Combustion: A substance reacts with oxygen, releasing energy, usually producing CO2 and H2O. Example:

Stoichiometry and Mole Relationships

Stoichiometry involves using balanced equations to relate amounts of reactants and products.

• Use mole ratios from the balanced equation to convert between substances.

• Conversions may involve grams, moles, or particles.

Example: How many moles of CO2 are produced from 2 moles of C2H6 in the reaction ? Answer:

Limiting Reactant, Theoretical Yield, and Percent Yield

• Limiting Reactant: The reactant that is completely consumed first, limiting the amount of product formed.

• Theoretical Yield: The maximum amount of product possible from given reactants.

• Percent Yield: The ratio of actual yield to theoretical yield, expressed as a percentage.

Formulas:

• Percent Yield:

Chapter 8: Introduction to Solutions and Aqueous Reactions

General Properties of Solutions

• Electrolyte: A substance that dissolves in water to produce ions, conducting electricity (e.g., NaCl).

• Nonelectrolyte: A substance that dissolves in water but does not produce ions (e.g., sugar).

• Weak Electrolyte: Partially ionizes in solution (e.g., acetic acid).

• Molarity (M):

• To find moles:

• To find volume:

• Dilution Equation:

Precipitation Reactions

• Involve the formation of an insoluble product (precipitate) when two solutions are mixed.

• Usually double displacement reactions:

• Use solubility guidelines to predict if a precipitate forms.

• Balance both atoms and charges in the equation.

• Ionic Equation: Shows all ions present.

• Spectator Ions: Ions that do not participate in the reaction.

• Net Ionic Equation: Shows only the ions and compounds that change during the reaction.

Acid-Base Reactions

• Strong Acids: Completely ionize in water (e.g., HCl, HNO3, H2SO4).

• Strong Bases: Completely ionize in water (e.g., NaOH, KOH).

• Weak Acids: Partially ionize (e.g., acetic acid, phosphoric acid).

• Weak Base: Ammonia (NH3).

• Acid formulas usually start with H (e.g., H3PO4).

• Neutralization Reaction: Acid + Base → Salt + Water.

• Some acid-base reactions evolve gas (e.g., acid + carbonate).

• Titration: Analytical method to determine concentration using a reaction with known stoichiometry.

Oxidation-Reduction (Redox) Reactions

• Oxidation Number: Assigned to atoms to track electron transfer.

• Rules for assigning oxidation numbers (in order):

  • Elemental form: 0

  • Monatomic ion: charge of ion

  • Fluorine: -1; Oxygen: usually -2; Hydrogen: +1 with nonmetals, -1 with metals

  • Sum in a compound: 0; in a polyatomic ion: equals ion charge

• Oxidation: Loss of electrons (increase in oxidation number).

• Reduction: Gain of electrons (decrease in oxidation number).

• Oxidizing Agent: Causes oxidation (is reduced).

• Reducing Agent: Causes reduction (is oxidized).

Chapter 9: Thermochemistry

Forms of Energy

• Internal Energy (E): Total energy (kinetic + potential) within a system.

• Enthalpy (H): Heat content at constant pressure.

• Heat (q): Energy transferred due to temperature difference.

• Work (w): Energy transferred when an object is moved by a force.

First Law of Thermodynamics:

Exothermic and Endothermic Processes

• Exothermic: Releases heat to surroundings ().

• Endothermic: Absorbs heat from surroundings ().

Calorimetry

• Measures heat flow in a chemical or physical process.

• q = mC\Delta T, where m = mass, C = specific heat, = temperature change.

Thermochemical Equations and Hess's Law

• Thermochemical equations show enthalpy changes with reactions.

• Hess's Law: The enthalpy change for a reaction is the same, no matter how many steps it is carried out in.

• Enthalpy changes are additive for sequential reactions.

Enthalpy of Formation

• Standard Enthalpy of Formation (): Enthalpy change for forming 1 mole of a compound from its elements in their standard states.

• For an element in its standard state, .

• Calculating Reaction Enthalpy:

Chapter 10: Gases (Final Exam Only)

General Properties of Gases

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

• Gas particles move rapidly and are far apart.

Units of Pressure and Measurement

• Common units: atmosphere (atm), torr, mmHg, pascal (Pa).

• Conversions:

• Pressure measured with a manometer in the lab.

Gas Laws

• Boyle's Law: (at constant T, n)

• Charles' Law: (at constant P, n)

• Gay-Lussac's Law: (at constant V, n)

• Combined Gas Law:

Ideal Gas Law and Applications

• Ideal Gas Law:

• R = 0.0821 L·atm/(mol·K)

• Can be rearranged to solve for any variable.

• Molar Mass:

• Density:

• Stoichiometry: Use ideal gas law to relate moles and volume in reactions involving gases.

Partial Pressures and Mole Fractions

• Dalton's Law:

• Mole Fraction:

• Partial Pressure:

Effusion and Diffusion

• Effusion: Passage of gas through a tiny opening.

• Diffusion: Mixing of gases due to random motion.

• Graham's Law: Rate of effusion is inversely proportional to the square root of molar mass.

Appendix: Solubility Guidelines Table (for Precipitation Reactions)

Additional info: Table entries inferred from standard solubility rules for clarity.