Chemistry work book

1. Atomic Structure

  • Definition of Atom: The smallest particle of an element.

  • Structure of Atom:

    • Nucleus at the center composed of:

      • Protons (positive charge, mass = 1)

      • Neutrons (no charge, mass = 1)

    • Shells surrounding nucleus containing:

      • Electrons (negative charge, negligible mass)

  • Sub-atomic Particles:

    • Protons: +1 charge, mass = 1.

    • Neutrons: 0 charge, mass = 1.

    • Electrons: -1 charge, negligible mass.

  • Isotopes:

    • Atoms of the same element with the same number of protons but different numbers of neutrons.

    • Example: Oxygen isotopes 16O, 17O, and 18O.

2. Electronic Structure

  • Electronic Configuration: Arrangement of electrons in atomic orbitals.

  • Orbitals: Regions where electrons are likely to be found, characterized as:

    • s-orbitals (2 electrons, spherical)

    • p-orbitals (6 electrons, dumbbell-shaped)

    • d-orbitals (10 electrons, complex shapes)

3. The Periodic Table

  • Periodic Trends:

    • Atomic Radius: Decreases across a period, increases down a group.

    • Ionic Radius: Cations are smaller than their parent atoms; anions are larger.

    • Ionization Energy: Energy required to remove an electron.

    • Electronegativity: Tendency of an atom to attract electrons.

4. Chemical Bonding

  • Types of Bonds:

    • Covalent Bonding: Sharing of electron pairs between atoms.

    • Ionic Bonding: Transfer of electrons resulting in ion formation.

    • Metallic Bonding: Delocalized electrons in a metal lattice.

5. Structure of Matter

  • Simple Atomic and Molecular Structures:

    • Simple atomic (noble gases) existing as individual atoms.

    • Simple molecular (like H2O, CO2) composed of molecules held together by covalent bonds.

6. Intramolecular Bonds

  • Electronegativity: Measure of the ability of an atom to attract shared electrons.

  • Polarity: Distribution of electron density in molecules affecting intermolecular interactions.

  • Types of Intermolecular Forces:

    • London Forces (temporary dipoles)

    • Dipole-Dipole Forces

    • Hydrogen Bonds

7. Quantitative Chemistry

  • Chemical Formula: Representation of a compound using symbols.

  • Mole Concept: A mole is 6.02 x 10^23 particles (Avogadro's number).

  • Methods of Calculating Concentration: Concentration = moles/volume.

8. Energy Changes and Rates of Reactions

  • Factors Affecting Reaction Rates: Concentration, temperature, particle size, and catalysts.

  • Exothermic and endothermic reactions, with energy profiles comparing activation energies.

9. Acids and Bases

  • Properties of Acids: Proton donors, sour taste, and turn blue litmus red.

  • Properties of Bases: Proton acceptors, bitter taste, slippery feel, and turn red litmus blue.

10. Organic Chemistry

  • Functional Groups: Specific groups of atoms within molecules that determine the characteristics of those molecules.

  • Isomerism: Occurs when compounds have the same molecular formula but different structures.

11. Electrochemistry

  • Oxidation and Reduction Reactions: Identifying oxidizing and reducing agents in half-equations.

  • Galvanic and Electrolytic Cells: Converting energy forms, with calculations based on Faraday's laws.

12. Chemical Equilibrium

  • Le Chatelier's Principle: Predicts how the system will adjust if external conditions (temperature, pressure, concentration) change.

  • Equilibrium Constant: Relates the concentrations of products and reactants at equilibrium.