Chapter 5: Nuclear Chemistry

Atomic Structure and Nuclear Stability

Nuclear chemistry focuses on the structure and behavior of atomic nuclei, including the processes by which unstable nuclei change.

• Mass Number (A): The total number of protons and neutrons in the nucleus of an atom.

• Atomic Number (Z): The number of protons in the nucleus, which defines the element.

• Nuclear Forces: The strong nuclear force holds protons and neutrons together in the nucleus, overcoming the repulsive electromagnetic force between protons.

Radioactivity and Types of Radiation

Some nuclei are unstable and emit radiation to become more stable. This process is called radioactive decay.

• Radiation: The emission of energy as particles or waves from unstable nuclei.

• Radioisotope: An isotope with an unstable nucleus that emits radiation.

• Types of Radiation:

  • Alpha (α) particles: Helium nuclei (), low penetration, stopped by paper.

  • Beta (β) particles: High-speed electrons (), moderate penetration, stopped by aluminum.

  • Gamma (γ) rays: High-energy electromagnetic waves, high penetration, require lead or concrete for shielding.

  • Positrons: Positive electrons ().

  • Protons () and Neutrons (): Subatomic particles involved in nuclear reactions.

• Discovery: Becquerel discovered radioactivity; Marie and Pierre Curie isolated radioactive elements; Rutherford discovered alpha and beta particles; Villard discovered gamma rays.

Radioactive Decay and Nuclear Reactions

• Radioactive Decay: The spontaneous emission of radiation from an unstable nucleus.

• Nuclear Fission: Splitting of a heavy nucleus into lighter nuclei, releasing energy.

• Nuclear Fusion: Combining light nuclei to form a heavier nucleus, releasing energy.

• Transmutation: Conversion of one element into another via nuclear reactions. First artificial transmutation by Rutherford.

• Balanced Nuclear Equations: Both mass number and atomic number must be conserved. Example:

Detection and Measurement of Radioactivity

• Detection Methods: Geiger-Müller counter (detects ionizing particles), Scintillation counter (detects light flashes from radiation).

• Units:

  • Activity: Becquerel (Bq), Curie (Ci)

  • Absorbed Dose: Gray (Gy), rad

  • Biological Damage: Sievert (Sv), rem

• Ionizing Radiation: Can remove electrons from atoms, causing somatic (body) or genetic (heritable) damage.

• Penetrating Power: Gamma > Beta > Alpha

Half-Life and Radioisotope Applications

• Half-Life (t1/2): The time required for half of a radioactive sample to decay. Formula:

• Calculations: Determine remaining mass, elapsed time, or original mass using half-life formula.

• Medical Uses: Radioisotopes in imaging, cancer therapy, and tracers. Factors: half-life, type of radiation, biological compatibility.

• Radiation Therapy: External beam, internal (brachytherapy), systemic. Chemotherapy uses chemicals, not radiation.

• Background Radiation: Largest source is radon gas; average American exposure is about 3-6 mSv/year.

• Nuclear Power: About 20% of US electricity; uses nuclear fission of uranium-235.

Chapter 6: Ionic and Molecular Compounds

Atomic Bonding and the Octet Rule

Atoms react to achieve a stable electron configuration, often an octet (8 valence electrons).

• Octet Rule: Atoms gain, lose, or share electrons to achieve 8 valence electrons.

• Ionic Bonds: Formed by transfer of electrons from metals to nonmetals.

• Covalent Bonds: Formed by sharing electrons between nonmetals.

Ionic Compounds

• Formation: Metal atoms lose electrons to form cations; nonmetals gain electrons to form anions.

• Monatomic Ions: Single-atom ions (e.g., Na+, Cl-).

• Polyatomic Ions: Ions composed of multiple atoms (e.g., SO42-).

• Predicting Formulas: Combine ions in ratios that yield a neutral compound.

• Naming: Use element name for cation; anion ends in -ide. For metals with variable charge, indicate charge with Roman numerals (e.g., Fe2+ is iron(II)).

• Common Anions and Polyatomic Ions: (See Table 6.2 and course resources.)

Molecular (Covalent) Compounds

• Formation: Nonmetals share electrons to form molecules.

• Naming: Use prefixes to indicate number of each atom (e.g., CO2 is carbon dioxide).

Lewis Structures and Bonding

• Lewis Structures: Diagrams showing valence electrons and bonds. Can include single, double, or triple bonds.

• Dipole: A separation of charge in a bond or molecule. Dipole moment is shown with an arrow pointing toward the more electronegative atom.

Electronegativity and Bond Polarity

• Electronegativity (EN): The ability of an atom to attract electrons in a bond. Increases across a period, decreases down a group.

• Bond Polarity: Determined by difference in EN ():

  • Nonpolar covalent:

  • Polar covalent:

  • Ionic:

VSEPR Theory and Molecular Shape

• VSEPR: Valence Shell Electron Pair Repulsion theory predicts molecular shapes based on electron pair repulsion.

• Common Shapes:

  • Linear: 180°

  • Trigonal planar: 120°

  • Tetrahedral: 109.5°

Molecular Polarity and Intermolecular Forces

• Polar Molecule: Has an uneven distribution of charge due to bond polarity and molecular shape.

• Attractive Forces:

  • Dispersion Forces: Weak, present in all molecules.

  • Dipole-Dipole: Between polar molecules.

  • Hydrogen Bonding: Strongest, occurs when H is bonded to N, O, or F.

• Effect on Properties: Stronger forces lead to higher melting and boiling points.

Chapter 7: Chemical Quantities and Reactions

The Mole and Avogadro's Number

The mole is a counting unit for atoms, molecules, or ions, allowing chemists to relate mass to number of particles.

• Definition: 1 mole = particles (Avogadro's number).

• Conversions: Moles to particles and vice versa using Avogadro's number.

• Subscripts in Formulas: Indicate the number of each atom in a molecule.

Molar Mass and Stoichiometry

• Molar Mass: Mass of 1 mole of a substance (g/mol). Calculated by summing atomic masses from the periodic table.

• Conversions: Grams ↔ Moles ↔ Particles. Example:

• Stoichiometry: Quantitative relationships in chemical reactions. Use balanced equations to relate moles of reactants and products.

Chemical Equations and Reaction Types

• Components: Coefficients, physical states, arrows.

• Balancing: Ensure equal numbers of each atom on both sides.

• Types of Reactions:

  • Combination (Synthesis)

  • Decomposition

  • Single Replacement

  • Double Replacement

  • Combustion

Oxidation-Reduction (Redox) Reactions

• Oxidation: Loss of electrons, gain of oxygen, or loss of hydrogen.

• Reduction: Gain of electrons, loss of oxygen, or gain of hydrogen.

• OIL RIG: Oxidation Is Loss, Reduction Is Gain (of electrons).

• LEO goes GER: Lose Electrons = Oxidation, Gain Electrons = Reduction.

Law of Conservation of Mass and Mole Ratios

• Law of Conservation of Mass: Mass is neither created nor destroyed in a chemical reaction.

• Mole-Mole Factor: Ratio from balanced equation used to convert between moles of reactants and products.

• Stoichiometric Calculations: Use mole ratios to determine amounts of reactants or products. Example:

Reaction Conditions and Energy Changes

• Reaction Conditions: Sufficient energy and proper orientation are required for reactions to occur.

• Exothermic Reactions: Release heat; heat is a product.

• Endothermic Reactions: Absorb heat; heat is a reactant.

• Energy Diagrams: Show energy changes during reactions; exothermic reactions have products lower in energy than reactants.

Reaction Rates and Catalysts

• Factors Affecting Rate: Concentration, temperature, surface area, presence of a catalyst.

• Catalyst: Substance that increases reaction rate without being consumed. Biological catalysts are called enzymes.

Table: Types of Radiation and Their Properties

Table: Common Polyatomic Ions (Partial List)

Additional info: For full mastery, students should review all assigned homework, lecture slides, and textbook sample problems, as well as memorize key ions and practice balancing equations and stoichiometry problems.