Overview
This lecture covers the early history of chemistry, the development of atomic theory, basic chemical laws, the structure of atoms, and a detailed guide to chemical nomenclature for general chemistry.
Early History and Fundamental Laws
- Early chemistry was mostly alchemy, focusing on chemical changes like fire, brewing, and rust.
- Robert Boyle defined an element as a substance that cannot be broken down further.
- Joseph Priestley discovered oxygen and conducted early gas experiments.
- Antoine Lavoisier developed the law of conservation of mass: mass is neither created nor destroyed in reactions.
- Joseph Proust established the law of definite proportion: a compound always has the same element mass ratios.
- John Dalton created the law of multiple proportions: element mass ratios in compounds are small whole numbers.
Dalton’s Atomic Theory
- Each element is composed of tiny, indivisible particles called atoms.
- Atoms of a given element are identical, and different from other elements.
- Chemical compounds are formed by combinations of different elements’ atoms in fixed ratios.
- Dalton's theory integrated and explained the earlier chemical laws.
- Atomic theory was widely accepted only after experimental evidence and imaging techniques like AFM and STM.
Discovery of Subatomic Particles
- J.J. Thomson discovered the electron, a negatively charged particle much lighter than atoms.
- Millikan determined the charge and mass of the electron using the oil drop experiment.
- Radioactivity was discovered by Becquerel, with three types: alpha (He²⁺), beta (electrons), and gamma (energy).
Structure of the Atom and Isotopes
- Rutherford’s gold foil experiment revealed the nuclear atom: a dense, positively charged nucleus with electrons moving around it.
- The nucleus contains protons (positive) and neutrons (neutral); electrons occupy most of the atom’s space.
- The number of protons (atomic number, Z) defines the element; neutrons can vary (isotopes).
- Mass number (A) = protons + neutrons.
- Isotopic symbols: ( ^A_Z )Element.
Chemical Bonds
- Covalent bonds: atoms share electrons to form molecules.
- Ionic bonds: electrons are transferred, forming cations (positive) and anions (negative) held by electrostatic attraction.
- Example: NaCl is formed from Na⁺ and Cl⁻.
The Periodic Table
- Elements are divided into metals, nonmetals, and metalloids (border elements).
- Vertical columns = groups/families; horizontal rows = periods.
- Alkali metals (Group 1): +1 charge; Alkaline earth metals (Group 2): +2 charge.
- Halogens (Group 17): −1 charge; Noble gases (Group 18): neutral.
Chemical Nomenclature Overview
Type 1 Binary Ionic Compounds (Metal + Nonmetal)
- Name cation (element name), then anion (root + 'ide').
- No prefixes for number; formula reflects charge neutrality.
- Example: NaCl—sodium chloride; MgI₂—magnesium iodide.
Type 2 Binary Ionic Compounds (Transition Metals)
- Use roman numerals to specify cation charge.
- Anion keeps 'ide' ending.
- Example: FeS—iron(II) sulfide; PbO₂—lead(IV) oxide.
Polyatomic Ions and Patterns
- Memorize common polyatomic ions (e.g., OH⁻, SO₄²⁻, NO₃⁻).
- Anions ending in -ate have more oxygens than -ite.
- Naming patterns use prefixes (hypo-, per-) for less/more oxygen.
- Only memorize two polyatomic cations: NH₄⁺ (ammonium), H₃O⁺ (hydronium).
Binary Covalent Compounds (Nonmetal + Nonmetal)
- Use prefixes for both elements (except mono- on the first).
- Second element gets 'ide' ending.
- Example: CO—carbon monoxide; N₂O₄—dinitrogen tetroxide.
Acids
- Binary acids (H + nonmetal anion): hydro- + root + -ic acid.
- Oxyacids: -ate anion → -ic acid; -ite anion → -ous acid.
- Example: HCl—hydrochloric acid; HNO₃—nitric acid; H₂SO₄—sulfuric acid.
Acid Salts and Special Naming
- Acid salts: contain partially neutralized polyatomic ions (e.g., NaHSO₄—sodium hydrogen sulfate or sodium bisulfate).
- Use hydrogen or dihydrogen prefixes for polyatomic anions with H.
Key Terms & Definitions
- Element — A substance that cannot be broken down into simpler substances.
- Law of Conservation of Mass — Mass is neither created nor destroyed in chemical reactions.
- Law of Definite Proportion — A compound always has the same mass ratio of elements.
- Law of Multiple Proportions — Element mass ratios in compounds are small whole numbers.
- Atom — The smallest unit of an element retaining its properties.
- Isotope — Atoms of the same element with different numbers of neutrons.
- Ion — Charged atom or molecule (cation: positive, anion: negative).
- Covalent Bond — A bond formed by the sharing of electrons.
- Ionic Bond — A bond formed by the transfer and electrostatic attraction of electrons.
- Polyatomic Ion — Two or more atoms covalently bonded carrying a charge.
- Oxoanion — Polyatomic anion containing oxygen.
Action Items / Next Steps
- Memorize common polyatomic ions and their charges (see textbook Table 2.5).
- Practice naming and writing formulas for type 1, type 2, and type 3 compounds.
- Review section on chemical nomenclature and the periodic table families.
- Complete any assigned practice problems or readings from Zumdahl Chapter 2.
Joseph Priestley is the correct answer because:
- He was the scientist who first collected oxygen gas by a method called collecting over water, where oxygen gas is captured by displacing water in a container.
- This method involved heating mercuric oxide to release oxygen gas, which Priestley collected in a water-filled apparatus.
- His work was foundational in the discovery and study of oxygen as a distinct gas.
- In honor of his contributions, the American Chemical Society awards the Priestley Medal annually to outstanding experimental chemists.
So, Priestley is directly linked to both the experimental technique of collecting oxygen over water and the prestigious chemistry medal named after him.