General Chemistry 1 - Chapter 2: Atoms, Molecules, and Ions
Lecture Overview
Introduction to fundamental chemical laws
Dalton's atomic theory
Basic and modern atomic structure
Terminology: atomic number, mass number, atomic weight, etc.
Molecules, ions, covalent and ionic bonding
Introduction to the periodic table and trends
Naming simple compounds
Fundamental Chemical Laws
Law of Conservation of Mass
Mass is neither created nor destroyed in a chemical reaction.
Example: Methane + Oxygen -> Carbon Dioxide + Water
Total mass before = total mass after reaction.
Law of Definite Proportions
Compounds always contain the same proportion of elements by mass.
Law of Multiple Proportions
When two elements form more than one compound, the masses of one element that combine with a fixed mass of the other are in a ratio of small whole numbers.
Dalton's Atomic Theory (1808)
Elements are made of tiny particles called atoms.
Atoms of the same element are identical; different elements have different atoms.
Compounds are formed by the combination of atoms in fixed ratios.
Chemical reactions involve rearrangement of atoms.
Atomic Structure
Early Models and Discoveries
J.J. Thomson (1898): Discovered electrons using cathode ray tubes.
Ernest Rutherford (1911): Gold foil experiment led to discovery of nucleus.
Niels Bohr: Developed model where electrons orbit the nucleus.
Modern View
Atoms consist of a nucleus (protons and neutrons) with electrons in probability clouds.
Protons (+), neutrons (0), electrons (-).
Key Terminology
Atomic Number (Z): Number of protons in the nucleus.
Mass Number (A): Total number of protons and neutrons.
Atomic Weight: Weighted average of the masses of all naturally occurring isotopes.
Molecules and Ions
Ions
Cations: Positively charged ions (e.g., Na⁺).
Anions: Negatively charged ions (e.g., Cl⁻).
Covalent vs. Ionic Bonding
Covalent Bonds: Atoms share electrons (can be polar or nonpolar).
Ionic Bonds: Electrons are transferred between atoms, forming ions that attract each other.
Electronegativity: Ability of an atom to attract electrons in a bond.
The Periodic Table
History and Structure
Dmitri Mendeleev: Original periodic table based on atomic weight.
Henry Mosley: Revised table based on atomic number.
Periodic Trends
Atomic Radius: Decreases left to right; decreases bottom to top.
Effective Nuclear Charge: Increases left to right; increases bottom to top.
Ionization Energy: Increases left to right; increases bottom to top.
Electronegativity: Increases left to right; increases bottom to top.
Electron Affinity: Increases left to right; increases bottom to top.
Naming Compounds
Type 1 Binary Ionic Compounds
Name the cation first, then the anion.
Cation takes element name; anion takes root + "-ide".
Type 2 Binary Ionic Compounds
Transition metals capable of forming multiple cations.
Use Roman numerals to indicate cation charge.
Binary Covalent Compounds (Type 3)
Formed between two nonmetals.
Use prefixes to denote number of atoms (mono, di, tri, etc.).
Naming with Polyatomic Ions
Recognize and memorize common polyatomic ions.
Acids
Without Oxygen: Prefix "hydro-" and suffix "-ic".
With Oxygen: Based on anion name:
"-ate" becomes "-ic"
"-ite" becomes "-ous"
Conclusion
Comprehensive coverage of atomic theory, chemical laws, periodic trends, and compound naming.
Emphasized importance of understanding foundational concepts and terminology for future chemistry studies.
Encouragement to practice problems and engage with course materials.