GOB Chemistry Exam 2 Study Guide: Chemical Reactions, Acids & Bases, and Hydrocarbons

Chapter 8: Chemical Reactions and Stoichiometry

Parts and Balancing of Chemical Equations

Chemical equations represent the reactants and products in a chemical reaction. Balancing ensures the law of conservation of mass is obeyed.

• Reactants are substances consumed; products are substances formed.

• Balancing involves adjusting coefficients so the number of atoms of each element is equal on both sides.

• Example: $2H_2 + O_2 ightarrow 2H_2O$

Types of Chemical Reactions

Chemical reactions are classified by how reactants change into products.

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

• Decomposition: A single compound breaks down into two or more simpler substances.

• Single Displacement: An element replaces another in a compound.

• Double Displacement: Exchange of ions between two compounds.

Mole Relationships and Calculations

The mole is a fundamental unit for counting particles in chemistry.

• Avogadro's Number: $1 ext{ mol} = 6.02 imes 10^{23}$ (atoms, molecules, or ions)

• Use molar ratios from balanced equations to convert between moles of reactants and products.

• Example: If 2 moles of $H_2$ react, 2 moles of $H_2O$ are produced.

Molar Mass and Mass Calculations

Molar mass is the mass of one mole of a substance, used to convert between mass and moles.

• Molar Mass: Sum of atomic masses in a chemical formula (g/mol).

• Example: Molar mass of $H_2O$ is $2 imes 1.01 + 16.00 = 18.02$ g/mol.

• To find mass: $ext{mass} = ext{moles} imes ext{molar mass}$

Stoichiometry and Limiting Reactant

Stoichiometry uses balanced equations to predict amounts of products and reactants.

• Determine the limiting reactant by comparing mole ratios.

• Calculate the amount of product formed from a given amount of reactant.

• Dilution Equation: $M_1V_1 = M_2V_2$

Chapter 9: Energy, Equilibrium, and Acids & Bases

Potential and Kinetic Energy

Energy in chemical systems can be stored (potential) or in motion (kinetic).

• Potential Energy: Stored energy due to position or composition.

• Kinetic Energy: Energy of motion.

Endothermic vs. Exothermic Reactions

Reactions are classified by energy flow.

• Exothermic: Release energy; products have less energy than reactants.

• Endothermic: Absorb energy; products have more energy than reactants.

Reading Energy Diagrams

Energy diagrams show the energy changes during a reaction.

• Activation energy is the energy required to start a reaction.

• Exothermic reactions have products lower than reactants; endothermic, higher.

Chemical Equilibrium

At equilibrium, the rate of the forward reaction equals the rate of the reverse reaction.

• Equilibrium Constant: $K = rac{[ ext{products}]}{[ ext{reactants}]}$

• Le Châtelier’s Principle: System shifts to counteract changes in concentration, pressure, or temperature.

Acids and Bases: Definitions and Properties

Acids and bases are defined by their ability to donate or accept protons.

• Bronsted-Lowry Acid: Proton (H+) donor.

• Bronsted-Lowry Base: Proton (H+) acceptor.

• Strong acids/bases dissociate completely; weak acids/bases only partially.

pH, Buffers, and Neutralization

pH measures the acidity or basicity of a solution.

• pH Equation: $pH = - ext{log}[ ext{H}_3 ext{O}^+]$

• Water Ionization Constant: $K_w = [ ext{H}_3 ext{O}^+][ ext{OH}^-] = 1.0 imes 10^{-14}$

• Buffer: Solution that resists changes in pH; made from weak acid and its conjugate base.

• Neutralization: Acid reacts with base to form water and a salt.

Acid Dissociation Constants

Acid strength is measured by the acid dissociation constant ($K_a$).

• Equation: $K_a = rac{[ ext{A}^-][ ext{H}_3 ext{O}^+]}{[ ext{HA}]}$

• pKa: $pK_a = - ext{log} K_a$

• Henderson-Hasselbalch Equation: $pH = pK_a + ext{log} rac{[ ext{A}^-]}{[ ext{HA}]}$

Table: Common Acids and Their $K_a$ Values

Acid	Formula	$K_a$
Phosphoric acid	H3PO4	7.5 × 10-3
Nitrous acid	HNO2	4.5 × 10-4
Hydrofluoric acid	HF	3.5 × 10-4
Formic acid	HCOOH	1.8 × 10-4
Acetic acid	CH3COOH	4.3 × 10-5
Carbonic acid	H2CO3	4.3 × 10-7
Hydrosulfuric acid	H2S	9.1 × 10-8
Dihydrogen phosphate ion	H2PO4-	6.2 × 10-8
Hydrocyanic acid	HCN	4.9 × 10-10
Bicarbonate ion	HCO3-	5.6 × 10-11
Hydrogen phosphate	HPO42-	2.2 × 10-13

Chapter 13: Hydrocarbons and Organic Chemistry

Hydrocarbon Structure and Nomenclature

Hydrocarbons are organic compounds composed of carbon and hydrogen. Their structure and naming follow IUPAC rules.

• Condensed Structural Formula: Shows the arrangement of atoms without displaying all bonds.

• Line Structure: Simplified representation where lines represent bonds between carbon atoms.

• Alkanes: Single bonds only; Alkenes: At least one double bond; Alkynes: At least one triple bond.

• Example: Ethane: CH3CH3; Ethene: CH2CH2

Physical Properties of Hydrocarbons

Hydrocarbons vary in solubility and physical properties based on structure.

• Most hydrocarbons are nonpolar and insoluble in water.

• Solubility decreases as the hydrocarbon chain length increases.

Isomerism and Stereochemistry

Isomers have the same molecular formula but different structures.

• Structural Isomers: Differ in connectivity of atoms.

• Stereoisomers: Same connectivity, different spatial arrangement.

• Cis/Trans Isomers: Occur in alkenes due to restricted rotation around double bonds.

Reactions of Hydrocarbons

Hydrocarbons undergo characteristic reactions such as substitution, addition, and elimination.

• Hydrogenation: Addition of hydrogen to double or triple bonds.

• Halogenation: Addition of halogen atoms (Cl, Br, etc.) to hydrocarbons.

• Example: Halogenation of benzene can occur at three positions: ortho, meta, para.

Appendix: Periodic Table

The periodic table organizes elements by increasing atomic number and groups elements with similar chemical properties.

• Groups: Vertical columns; elements share similar properties.

• Periods: Horizontal rows; properties change progressively.

• Useful for determining atomic masses, valence electrons, and element classification.

Key Equations and Constants

• Avogadro's Number: $1 ext{ mol} = 6.02 imes 10^{23}$ (atoms, molecules, or ions)

• Molarity: $M = rac{ ext{mol}}{L}$

• Dilution Equation: $M_1V_1 = M_2V_2$

• Equilibrium Constant: $K = rac{[ ext{products}]}{[ ext{reactants}]}$

• Water Ionization Constant: $K_w = [ ext{H}_3 ext{O}^+][ ext{OH}^-] = 1.0 imes 10^{-14}$

• pH Equation: $pH = - ext{log}[ ext{H}_3 ext{O}^+]$

• Hydronium Ion Concentration: $[ ext{H}_3 ext{O}^+] = 10^{-pH}$

• Acid Dissociation Constant: $K_a = rac{[ ext{A}^-][ ext{H}_3 ext{O}^+]}{[ ext{HA}]}$

• pKa: $K_a = 10^{-pK_a}$

• Henderson-Hasselbalch Equation: $pH = pK_a + ext{log} rac{[ ext{A}^-]}{[ ext{HA}]}$

Additional info: Some context and explanations have been expanded for clarity and completeness, including definitions, examples, and academic context for formulas and tables.