Lecture on Atoms and Chemistry

Overview of Atoms

• Everything is made of atoms, including humans.
• Components of an Atom: Core (protons and neutrons) and electrons.
• Number of protons defines the element.

Atomic Structure

• Traditional View: Atom with multiple electron shells.
• Quantum Mechanics View: Atoms look different; more complex.
• Valence Electrons: Electrons in the outermost shell; crucial for understanding chemistry.

Periodic Table

• Groups: Same number of valence electrons.
  • Exception: Helium (only 2 valence electrons but behaves like noble gas).
• Periods: Same number of shells; increases top to bottom.
• Element Properties: Similar behavior in chemical reactions within the same group.
• Alkali Metals: One valence electron, shiny, soft, and reactive.
• Isotopes: Same element, different number of neutrons, often unstable and radioactive.

Ions

• Neutral Atom: Equal number of protons and electrons.
• Charged Atoms: Ions (positive = cations, negative = anions).

Periodic Table Information

• Provides element name, symbol, number of protons/electrons, and atomic mass.
• Categories: Metals, non-metals, and semimetals.

Molecules and Compounds

• Molecule: Two or more atoms bonded together.
• Compound: Molecule with at least two different elements.
• Molecular Formula: Shows the number of each atom in a molecule.
• Isomers: Same molecular formula, different structures (e.g., graphite vs diamonds).
• Lewis-Dot Structure: Represents valence electrons and bonds.

Chemical Bonds

• Covalent Bonds: Sharing of electrons; atoms achieve lower energy state.
  • Electronegativity: Strength of nucleus' pull on electrons; increases bottom left to top right in periodic table.
  • Ionic Bonds: Significant electronegativity difference (>1.7); involves electron transfer (e.g., NaCl).
  • Metallic Bonds: Electrons are delocalized in a metal grid.
  • Nonpolar Covalent Bonds: Electrons shared equally (difference in electronegativity < 0.5).
  • Polar Covalent Bonds: Unequal sharing; creates electric dipoles.
• Intermolecular Forces (IMFs): Forces acting between molecules (e.g., hydrogen bonds, Van der Waals forces).

States of Matter

• Solid: Tightly packed particles, fixed structure.
• Liquid: Particles move freely, fixed volume.
• Gas: Particles fill available volume.
• Plasma: Ionized gas at high temperatures or electric potential.

Temperature and Entropy

• Temperature: Average kinetic energy of particles.
• Entropy: Measure of disorder; increases with higher temperature and lower pressure.
• Phase Changes: Solids at low temperature/high pressure; gases at high temperature/low pressure.
• Enthalpy: Internal energy/heat content of a system.
• Gibbs Free Energy: Determines if a reaction is spontaneous (exergonic) or not (endergonic).

Chemical Reactions

• Types: Synthesis, decomposition, single replacement, double replacement.
• Stoichiometry: Ratios based on conservation of mass.
• Activation Energy: Required for reactions to occur; catalysts reduce this energy.
• Exothermic vs. Endothermic: Exothermic releases heat; endothermic absorbs heat.
• Chemical Equilibrium: Reversible reactions occur at the same rate; concentrations remain constant.

Acids and Bases

• Brondsted-Lowry Definition: Acids donate protons; bases accept protons.
• pH Scale: Measures acidity; lower pH = more acidic.
• Amphoteric Molecules: Can act as both acid and base.
• Redox Reactions: Transfer of electrons, changing oxidation states.

Quantum Mechanics and Electron Configuration

• Quantum Numbers: Describe electron states (n, l, ml, ms).
• Aufbau Principle: Order of filling electron subshells.
• Electron Configuration: Determines element's chemical behavior.