General Chemistry

Introduction to Atoms and Elements

  • Atoms: The fundamental building blocks of matter.

    • Core: Contains protons and neutrons.

    • Electrons: Orbit the core.

  • Elements: Different types of atoms based on their number of protons.

    • Example: Water is made of hydrogen and oxygen.

Valence Electrons and the Periodic Table

  • Valence Electrons: Electrons in the outermost shell of an atom.

    • Most chemical reactions involve the behavior of valence electrons.

  • Periodic Table: A table organizing elements based on their properties.

    • Groups: Columns that represent elements with the same number of valence electrons.

    • Periods: Rows that represent elements with the same number of electron shells.

Isotopes, Ions, and Reading the Periodic Table

  • Isotopes: Atoms of the same element with different numbers of neutrons.

  • Ions: Charged atoms.

    • Cations: Positive ions.

    • Anions: Negative ions.

  • Reading the Periodic Table:

    • Symbol: Abbreviation for the element.

    • Atomic Number: Number of protons.

    • Atomic Mass: Average mass of atoms of that element.

    • Metals: Left side of the periodic table.

    • Non-metals: Right side of the periodic table.

    • Semimetals: Between metals and non-metals.

Molecules and Compounds

  • Molecules: Two or more atoms bonded together.

  • Compounds: Molecules composed of at least two different elements.

    • Example: Water (H2O) is a compound.

  • Molecular Formula: Shows the number of each atom in a molecule.

  • Isomers: Molecules with the same molecular formula but different structures.

Chemical Bonding

  • Covalent Bonds: Sharing of electrons between atoms.

    • Electronegativity: The ability of an atom to attract electrons.

  • Ionic Bonds: Transfer of electrons between atoms, forming ions.

  • Metallic Bonds: Sharing of electrons between metal atoms.

  • Intermolecular Forces: Forces between molecules.

    • Hydrogen Bonds: Strong dipole-dipole interactions.

    • Van der Waals Forces: Weak temporary attractions.

States of Matter

  • Solid: Fixed shape and volume.

  • Liquid: Fixed volume, but variable shape.

  • Gas: No fixed shape or volume.

  • Plasma: Highly ionized gas.

  • Temperature: Average kinetic energy of particles.

  • Entropy: Measure of disorder.

Chemical Reactions

  • Types of Chemical Reactions: Synthesis, decomposition, single replacement, double replacement.

  • Stoichiometry: The quantitative relationship between reactants and products.

  • Balancing Equations: Ensuring the same number of atoms of each element on both sides of a chemical equation.

  • The Mole: A unit of measurement for the amount of a substance.

  • Physical vs. Chemical Change: Physical changes alter the appearance without changing the substance, while chemical changes create new substances.

  • Activation Energy: The energy required to start a chemical reaction.

  • Catalysts: Substances that speed up reactions without being consumed.

  • Enthalpy: The heat content of a system.

    • Exothermic: Releases heat.

    • Endothermic: Absorbs heat.

  • Gibbs Free Energy: Determines the spontaneity of a reaction.

  • Chemical Equilibrium: A state where the rates of the forward and reverse reactions are equal.

Acids and Bases

  • Brønsted-Lowry Theory: Acids donate protons, bases accept protons.

  • pH: Measure of acidity or basicity.

  • Neutralization Reactions: Reactions between acids and bases that produce water and a salt.

  • Redox Reactions: Reactions involving the transfer of electrons.

    • Oxidation Numbers: Indicate the degree of oxidation or reduction of an atom.

Quantum Chemistry

  • Quantum Numbers: Describe the properties of electrons in an atom.

    • n: Principal quantum number (shell).

    • l: Azimuthal quantum number (subshell).

    • ml: Magnetic quantum number (orbital).

    • ms: Spin quantum number.

  • Electron Configuration: The arrangement of electrons in an atom's orbitals.

  • Aufbau Principle: The order in which orbitals are filled.