Chemistry Lecture Notes

Atoms and Elements

• Atoms: Basic building blocks of matter, including humans.
  • Consist of a core (protons and neutrons) and electrons.
• Elements: Defined by the number of protons.
  • Example: Water consists of Hydrogen and Oxygen.

Atomic Structure

• Electron Shells: Electrons arranged in shells; valence electrons are in the outermost shell.
• Periodic Table: Lists elements with the same number of valence electrons in columns.
  • Main groups: Group number indicates the number of valence electrons (except helium).
  • Transition metals follow different patterns.

Chemical Behavior and Bonds

• Valence Electrons: Determine chemical reactions.
• Alkali Metals: One valence electron, shiny, soft.
• Periods and Groups: Same number of shells in a period, similar chemical behavior in a group.
• Isotopes: Different numbers of neutrons, often unstable.
• Ions: Atoms with a charge; cations (+) and anions (-).
  • Formed by gaining or losing electrons.

Periodic Table Insights

• Element Information: Name, symbol, protons, electrons, atomic mass.
• Categories: Metals, non-metals, and semimetals.
• Molecules and Compounds: Atoms bonded together form molecules; different elements form compounds.

Molecular Structure and Bonds

• Lewis-Dot Structures: Show valence electrons and bonds.
• Electron Shell Stability: Atoms aim for full outer shells (often 8 electrons).
• Types of Bonds:
  • Covalent Bonds: Sharing of electrons.
  • Ionic Bonds: Transfer of electrons; significant difference in electronegativity.
  • Metallic Bonds: Valence electrons are delocalized, found in metals.

Bond Strength and Types

• Electronegativity: Increases from bottom left to top right in the periodic table.
  • Fluorine has the highest.
• Ionic vs Covalent: Ionic bonds are stronger than covalent.
• Polar vs Nonpolar: Based on electronegativity difference.
  • Water is an example of a polar molecule.

Interactions Between Molecules

• Hydrogen Bonds: Strong dipole interactions (e.g., H with F, O, or N).
• Van der Waals Forces: Even nonpolar molecules can have temporary dipoles.
• Solubility: "Like dissolves like." Water dissolves polar substances.
  • Soap uses surfactants with a polar head and nonpolar tail.

States of Matter

• Solid, Liquid, Gas: Defined by particle arrangement and movement.
• Temperature and Entropy:
  • Temperature: Average kinetic energy.
  • Entropy: Amount of disorder.

Chemical Reactions

• Types: Synthesis, decomposition, single replacement, double replacement.
• Stoichiometry: Ratios of reactants and products based on conservation of mass.
• Activation Energy: Necessary for reactions, reduced by catalysts.

Energy in Reactions

• Enthalpy: Internal heat content.
  • Exothermic: Releases heat.
  • Endothermic: Absorbs heat.
• Gibbs Free Energy: Determines spontaneous reactions based on enthalpy and entropy.

Equilibrium and Acids/Bases

• Chemical Equilibrium: Reactions occur at the same rate in both directions.
• Acids and Bases: Defined by proton donation and acceptance.
  • Amphoteric: Acts as both acid and base.
  • pH Scale: Measures acidity; pH + pOH = 14.
  • Redox Reactions: Transfer of electrons, changes in oxidation states.

Quantum Mechanics in Chemistry

• Quantum Numbers: Describe electron positions and behaviors.
  • Principal Quantum Number (n): Indicates shells.
  • Subshells and Orbitals: s, p, d, f; varied electron capacities.
• Pauli Exclusion Principle: No two electrons can have the same set of quantum numbers.
• Aufbau Principle: Order of electron filling in orbitals.

Final Notes

• Understanding electron configuration helps determine valence electrons and chemical reactivity.
• The periodic table assists in predicting element behavior based on its structure.