Lecture Notes on Atomic Structure and Basic Chemistry

Composition of Matter

• Everything is made of atoms.
• Atoms consist of a core (protons and neutrons) and electrons.
• Different elements are defined by the number of protons.
• Example: Water is composed of Hydrogen and Oxygen.

Quantum Mechanics and Atomic Models

• Quantum mechanics presents a more complex view of atoms than simple models.
• Atoms have multiple electron shells; outermost electrons are called valence electrons.
• Valence electrons play a critical role in chemical reactions.

Periodic Table

• Elements are listed in the periodic table.
• Elements in the same group have the same number of valence electrons.
• Elements in the same period have the same number of shells.
• Atomic mass increases from left to right in a period.

Isotopes and Ions

• Different isotopes have varying numbers of neutrons.
• Unstable isotopes can release ionizing radiation.
• Atoms with equal protons and electrons are neutral; unequal numbers create ions (anions and cations).

Periodic Table Information

• Each cell details the element's name, symbol, proton count, and atomic mass.
• Rough categorization: metals, non-metals, and semi-metals.

Molecules and Compounds

• Molecules: Two or more bonded atoms.
• Compounds: Molecules consisting of at least two different elements.
• Isomers: Different structures with the same molecular formula.
• Lewis-Dot Structure: Represents valence electrons and bonds.

Chemical Bonds

• Covalent Bonds: Sharing of electrons between atoms.
• Electronegativity: The measure of an atom’s ability to tug on electrons.
• Ionic Bonds: Formed when the electronegativity difference is large (>1.7).
• Metallic Bonds: In metals, involve delocalized electrons.
• Polar and Nonpolar Covalent Bonds: Depend on the difference in electronegativity.
• Hydrogen Bonds: Special type of dipole interaction.
• Van der Waals Forces: Momentary dipoles creating weak attractions.

States of Matter

• Solid: Fixed structure, low entropy, particles only wiggle.
• Liquid: Particles move freely but constrained to a volume.
• Gas: High energy, fills any volume, high entropy.
• Plasma: Ionized gas, very high temperatures or electric potential.

Thermodynamics and Reactions

• Temperature: Average kinetic energy of particles.
• Entropy: Measure of disorder.
• Reactions tend toward lower energy states; exothermic and endothermic processes.
• Activation Energy: Energy needed to start a reaction; catalysts lower this energy.
• Enthalpy: Internal energy or heat content of a system.
• Gibbs Free Energy: Determines reaction spontaneity.

Chemical Equilibrium

• Occurs when reversible reactions happen at the same rate in both directions.
• Equilibriums are common in phase changes and acid-base chemistry.
• Acids: Proton donors.
• Bases: Proton acceptors.
• pH: Measure of hydronium ion concentration.
• pOH: Measure of hydroxide ion concentration.
• Equilibrium: pH + pOH = 14.
• Acid-Base Neutralization forms water and salt.

Redox Reactions

• Reduction-Oxidation Reactions: Involve the transfer of electrons.
• Oxidation Numbers: Imaginary charges to track electron flow.

Quantum Mechanics and Electron Configuration

• Electrons described by four quantum numbers (n, l, ml, ms).
• Orbitals: Regions where electrons are likely found.
• Pauli Exclusion Principle: No two electrons can have the same four quantum numbers.
• Aufbau Principle: Order for filling up electron subshells.

Conclusion

• Elements and compounds exhibit behavior based on their atomic structure and chemical bonds.

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