Atoms and Elements

Atoms: The basic unit of matter, made of a core (protons and neutrons) and electrons.
Elements: Determined by the number of protons.
Water: Composed of Hydrogen and Oxygen.

Structure of Atoms

Electron Shells: Atoms have multiple shells; outermost electrons are called valence electrons.
- Most chemistry revolves around the behavior of these electrons.
Periodic Table:
- Elements in the same group have the same number of valence electrons.
- Group numbers indicate the number of valence electrons (1-8), except Helium (2).
- Alkali Metals: Group 1 elements (excluding Hydrogen), common traits include:
  - One valence electron
  - Shiny and soft metals

Isotopes: Atoms of the same element with different numbers of neutrons, often unstable.
Ions: Charged atoms;
- Anions: Negative charge (more electrons).
- Cations: Positive charge (fewer electrons).

Molecules: Two or more atoms bonded together.
Compounds: At least two different elements.
- Example: Explosive metal + toxic gas = table salt (NaCl).
Molecular Formulas: Represent the number of each atom in a molecule.
- Isomers: Molecules with the same formula but different structures (e.g., graphite vs. diamonds).

Covalent Bonds: Atoms share electrons to achieve full outer shells.
- Driven by energy reduction (lower potential energy).
Electronegativity: The ability of an atom to attract electrons.
- Increases from bottom left to top right of the periodic table.
- Example: Fluorine has the strongest pull.
- Ionic Bonds form when electronegativity difference > 1.7 (e.g., Sodium Chloride).
Metallic Bonds: Delocalized electrons in a grid of positively charged nuclei, responsible for conductivity and malleability.
Types of Covalent Bonds:
- Nonpolar Covalent Bonds: Electrons shared equally (difference < 0.5).
- Polar Covalent Bonds: One atom pulls on electrons more (0.5 < difference < 1.7).

Three Main States: Solid, Liquid, Gas.
Temperature: Average kinetic energy of particles.
Entropy: Measure of disorder.
- Solids at low temperatures (ordered) vs. gases at high temperatures (disordered).
Plasma: Ionized gas at extremely high temperatures.

Types of Reactions:
- Synthesis
- Decomposition
- Single Replacement
- Double Replacement
Stoichiometry: Ratios of reactants needed for reactions.
Mole Concept: Measurement of particles (1 mole = atomic mass in grams).
Chemical Changes: Accompanied by observable changes (bubbles, color changes, etc.).

Activation Energy: Energy needed to start reactions.
Catalysts: Lower activation energy, not consumed in reactions.
Enthalpy: Total internal energy of a system.
Exothermic Reactions: Release heat (enthalpy decreases).
Endothermic Reactions: Absorb heat (enthalpy increases).
Gibbs Free Energy: Determines spontaneity of reactions.
- Exergonic: Negative change (spontaneous).
- Endergonic: Positive change (non-spontaneous).

Acids and Bases (Bronsted-Lowry):
- Acids donate protons; bases accept protons.
- Amphoteric: Substances that can act as either an acid or a base.
- pH: Measurement of acidity (negative log of hydronium ion concentration).
- Strong acids dissociate completely; weak acids do not.
Neutralization: Reaction between acids and bases to form water and salt.

Reduction-Oxidation (Redox): Involves changes in oxidation numbers.
- Oxidation: Loss of electrons.
- Reduction: Gain of electrons.

Quantum Numbers: Describe electrons in atoms.
- N (shell), l (subshell), ml (orbital), ms (spin).
Electron Configuration: Filling of electron subshells following the Aufbau principle.
- Example: Electron configuration of Sodium.
- Valence Electrons: Can be determined by electron configuration.

The lecture covered fundamental concepts in chemistry, including atomic structure, bonding, reactions, and the behavior of matter.