Lecture Notes on Chemistry Topics

1. Atomic Structure and Properties

Periodic Table

• Groups: Alkali metals, alkaline earth metals, transition metals, halogens, noble gases.
• Mass Number: Protons (P) + Neutrons (N).
• Isotopes: Atoms of the same element with different neutron numbers.
• Average Atomic Mass: Calculated from isotope mass and relative abundance.

Moles and Calculations

• Ideal Gas Law: PV = nRT.
• Avogadro’s Number: 6.022 x 10^23.
• Standard Conditions (STP): 1 atm and 273K, 22.4 L/mol.
• Molarity (M): moles of solute per liter of solution.
• Percent Composition: Ratio of mass of each element to total molar mass.
• Formulas: Empirical (simplest ratio) and molecular (actual ratio).

Energy Levels

• Electron Potential Energy: Increases with distance from the nucleus.
• Quantized Energy Levels: Specific allowed energy levels for electrons.
• Coulomb’s Law: F = kq1q2/r^2.
• Electron Transitions: Electrons absorb or emit energy when changing energy levels.

Photoelectron Spectroscopy

• Energy Units: Measured in electronvolts (eV).
• Energy Equation: Incoming radiation energy = binding energy + kinetic energy.
• Photoelectron Spectrum: Represents different energy levels and subshells.

Electron Configuration

• Notation: Spdf, using noble gas shorthand.
• Rules:
  • Aufbau Principle: Electrons fill lowest subshells first.
  • Pauli Exclusion Principle: 2 electrons/orbital, opposite spins.
  • Hund’s Rule: Maximize unpaired electrons in subshells.

Periodic Trends

• Attractions: Increase with proximity to nucleus or more protons.
• Shielding Effect: Inner electrons reduce attraction between nucleus and valence electrons.
• Trends: Atomic radius increases down the group; ionization energy and electronegativity increase across periods.

2. Molecular and Ionic Compound Structure and Properties

Types of Bonds

• Ionic Bonds: Electrostatic attraction in lattice between cations and anions.
• Metallic Bonds: Sea of electrons model, formation of alloys.
• Covalent Bonds: Shared electron pairs between nonmetals, forms molecules.
• Network Covalent Bonds: Strong lattice, high melting/boiling points.

Conductivity

• Ionic Compounds: Conduct in aqueous and liquid forms.
• Molecular Covalent: Generally do not conduct.
• Metallic Bonds: Conduct in solid and liquid phases.

Lewis Structures and Geometry

• Lewis Dot Structures: Represent valence electrons.
• VSEPR Theory: Predicts molecular shapes based on electron pair repulsion.

3. Intermolecular Forces and Properties

Polarity

• Polar Covalent Bonds: Unequal electron sharing; creates dipoles.
• Molecular Polarity: Depends on bond polarity and geometry.

Intermolecular Forces (IMF)

• Dipole-Dipole: Attractions between polar molecules.
• Hydrogen Bonds: Strong dipole-dipole attractions involving F, O, or N.
• London Dispersion Forces: Weak forces due to temporary dipoles.

Vapor Pressure

• Depends on IMF: Stronger IMFs lead to lower vapor pressure.

Solution Separation

• Chromatography: Separates based on polarity and solubility.
• Distillation: Uses boiling points to separate components.

4. Chemical Reactions

Types of Reactions

• Synthesis, Decomposition, Acid-Base, Redox, Combustion, Precipitation.

Solubility Rules

• Always Soluble: Alkali metals and ammonium ions.

Oxidation States and Redox

• Neutral Atoms: Oxidation state of zero.
• Redox Reactions: Identify oxidation and reduction components.

5. Kinetics

Rate Laws

• Rate Equation: Rate = k [A]^x [B]^y.
• Orders of Reaction: Determined experimentally.

Collision Theory

• Factors Affecting Rate: Concentration, temperature, orientation.

6. Thermodynamics

Concepts

• Enthalpy (ΔH): Heat content, indicates exothermic or endothermic reactions.
• Entropy (ΔS): Disorder of the system.
• Gibbs Free Energy (ΔG): ΔG = ΔH - TΔS, determines spontaneity.

7. Equilibrium

Equilibrium Constant (K)

• Expression: Relates concentrations of reactants and products.
• Le Chatelier’s Principle: Predicts shifts in equilibrium with changes in conditions.

8. Acids and Bases

pH and pOH

• Formulas: pH = -log[H+], pOH = -log[OH-].

Strong vs Weak Acids/Bases

• Dissociation: Strong acids/bases dissociate completely, weak ones do not.

Buffers

• Buffer Solutions: Resist changes in pH.

9. Applications of Thermodynamics

Galvanic Cells

• Redox Reactions: Generate electrical current.

Electrolytic Cells

• Nonspontaneous Reactions: Drive reactions using external voltage.

10. Laboratory Techniques

• Measurement Precautions: Avoid errors by considering factors like convection currents and thorough equipment rinsing.