Chemistry Lecture Notes

Atoms and Elements

• Everything is made of atoms.
• Atoms consist of:
  • Core (protons and neutrons)
  • Electrons (in multiple shells)
• Valence Electrons: Electrons in the outermost shell.
• Elements are identified by their number of protons.

The Periodic Table

• Elements with the same number of valence electrons have similar chemical behavior.
• Groups: Elements in the same column (same valence electrons).
  • Group # corresponds to number of valence electrons (1-8, except Helium).
• Periods: Elements in the same row have the same number of shells.
  • Mass increases from left to right.

Isotopes and Ions

• Different isotopes have different numbers of neutrons.
• Ions: Charged atoms.
  • Anions: Negative charge (more electrons).
  • Cations: Positive charge (fewer electrons).

Molecular Structure

• Molecules: Two or more atoms bonded together.
• Compounds: Molecules with at least two different elements.
• Molecular formulas count atoms of each type.
• Isomers: Same molecular formula but different structures (e.g., graphite vs. diamonds).

Bonding

• Covalent Bond: Atoms share electrons to achieve full outer shells (lower energy state).
• Electronegativity: Measure of an atom's ability to attract electrons.
  • Increases from bottom left to top right in the periodic table.
  • Example: Fluorine has the strongest electronegativity.
• Ionic Bonds: Formed when electronegativity difference is > 1.7 (e.g., Sodium Chloride).
• Metallic Bonds: Found in metals; electrons are delocalized and free to move.

Intermolecular Forces (IMFs)

• Hydrogen Bonds: Special dipole-dipole interactions involving hydrogen and highly electronegative atoms (e.g., O, N, F).
• Van der Waals Forces: Temporary dipoles in nonpolar molecules.

States of Matter

• Solid: Tightly packed particles, fixed structure.
• Liquid: Free-moving particles, fixed volume.
• Gas: Particles have high energy and fill available volume.
• Plasma: Ionized gas at very high temperatures.

Chemical Reactions

• Types of reactions: synthesis, decomposition, single replacement, double replacement.
• Stoichiometry: Ratio of reactants in a reaction, based on conservation of mass.
• Moles: Measure of quantity in reactions (1 mole = atomic mass in grams).
• Activation energy is required to start a chemical reaction.

Energy Changes

• Enthalpy: Internal energy of a system; exothermic (heat released) vs. endothermic (heat absorbed).
• Gibbs Free Energy: Determines spontaneity of a reaction.
  • Negative ΔG: spontaneous.
  • Positive ΔG: non-spontaneous.

Acid-Base Chemistry

• Bronsted-Lowry Theory: Acids donate protons, bases accept protons.
• pH: Measure of hydronium ion concentration; scale from 0-14.
• Strong acids/bases dissociate completely; weak acids/bases do not.

Redox Reactions

• Involves transfer of electrons; oxidation (loss of electrons) vs. reduction (gain of electrons).
• Oxidation numbers help identify electron flow in reactions.

Quantum Chemistry

• Electrons defined by four quantum numbers: n, l, ml, ms.
• Electron configurations follow the Aufbau principle.
• Valence Electrons: Determined by electron configurations.

Conclusion

• Key concepts in chemistry: atoms, bonding, reactions, and energy changes.
• Understanding these foundations is critical for chemistry studies.