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: $2H_2 + O_2 \rightarrow 2H_2O$

• Decomposition: One substance breaks down into two or more products. Example: $2KClO_3 \rightarrow 2KCl + 3O_2$

• Combustion: A substance reacts with oxygen, releasing energy, usually producing CO2 and H2O. Example: $CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O$

Stoichiometry and Limiting Reactants

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

• Convert between grams and moles using molar mass.

• Use mole ratios from the balanced equation to relate quantities of substances.

• 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: $\text{Percent Yield} = \frac{\text{Actual Yield}}{\text{Theoretical Yield}} \times 100\%$

Chapter 8: Introduction to Solutions and Aqueous Reactions

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): $M = \frac{\text{moles of solute}}{\text{liters of solution}}$

• To find moles: $\text{moles} = M \times \text{volume (L)}$

• Dilution Equation: $M_1V_1 = M_2V_2$

Precipitation Reactions

• Double displacement reactions where an insoluble solid (precipitate) forms.

• Predict products by swapping ions ("swapping dance partners").

• Use solubility guidelines to determine if a precipitate forms.

• Balance both charge and number of atoms in the equation.

• Ionic Equation: Shows all ions present in solution.

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

• Net Ionic Equation: Shows only the ions and molecules directly involved in 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., HC2H3O2, H3PO4).

• Weak Base: Ammonia (NH3).

• Write balanced acid-base neutralization reactions (acid + base → salt + water).

• Write balanced gas-evolving acid-base reactions (e.g., acid + carbonate → CO2 gas).

• Perform titration calculations using $M_1V_1 = M_2V_2$ for monoprotic acids/bases.

Oxidation-Reduction (Redox) Reactions

• Assign oxidation numbers using standard rules (elemental form = 0, group 1 = +1, group 2 = +2, F = -1, etc.).

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

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

• Oxidizing Agent: Causes oxidation (is reduced).

• Reducing Agent: Causes reduction (is oxidized).

Chapter 9: Thermochemistry

Forms of Energy

• Internal Energy (E): The total energy (kinetic + potential) of a system.

• Enthalpy (H): The heat content of a system at constant pressure.

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

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

Exothermic vs. Endothermic Processes

• Exothermic: Releases heat to surroundings ($\Delta H < 0$).

• Endothermic: Absorbs heat from surroundings ($\Delta H > 0$).

Calorimetry

• Measures heat flow in a chemical or physical process.

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

Thermochemical Equations and Hess's Law

• Thermochemical equations show enthalpy changes for reactions.

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

• Add or subtract equations to find $\Delta H$ for a target reaction.

Enthalpy of Formation

• Standard Enthalpy of Formation ($\Delta H_f^\circ$): Enthalpy change for forming 1 mole of a compound from its elements in their standard states.

• Calculating Reaction Enthalpy: $\Delta H_{rxn}^\circ = \sum n\Delta H_f^\circ (\text{products}) - \sum n\Delta H_f^\circ (\text{reactants})$

Chapter 10: Gases (Final Exam Only)

General Properties of Gases

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

Units of Pressure

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

• Conversions: 1 atm = 760 mmHg = 760 torr = 101,325 Pa

Measuring Pressure

• Open-end manometer: measures gas pressure relative to atmospheric pressure.

Gas Laws

• Boyle's Law: $P_1V_1 = P_2V_2$ (at constant T and n)

• Charles' Law: $\frac{V_1}{T_1} = \frac{V_2}{T_2}$ (at constant P and n)

• Gay-Lussac's Law: $\frac{P_1}{T_1} = \frac{P_2}{T_2}$ (at constant V and n)

• Combined Gas Law: $\frac{P_1V_1}{T_1} = \frac{P_2V_2}{T_2}$

Ideal Gas Law and Applications

• Ideal Gas Law: $PV = nRT$

• R = 0.0821 L·atm/(mol·K)

• Use to solve for unknowns (P, V, n, or T).

• Molar Mass: $M = \frac{mRT}{PV}$

• Density: $d = \frac{MP}{RT}$

• Stoichiometry: Use ideal gas law to relate moles of gas to volume under non-standard conditions.

Partial Pressures and Mole Fractions

• Dalton's Law: $P_{total} = P_1 + P_2 + ...$

• Mole Fraction: $X_i = \frac{n_i}{n_{total}}$

• Partial Pressure: $P_i = X_i \times P_{total}$

Effusion and Diffusion

• Effusion: Gas particles escape through a small hole.

• Diffusion: Gas particles spread out to fill a space.

• Graham's Law: $\frac{\text{Rate}_1}{\text{Rate}_2} = \sqrt{\frac{M_2}{M_1}}$

Table: Strong Acids and Bases

Additional info: This guide summarizes the main concepts and skills required for Chapters 7–10, including key equations, definitions, and examples. For practice, refer to the end-of-chapter problems listed in your syllabus.