Honors Chemistry

1. Honors Chemistry Chapter 2: Atoms and Molecules

2. 2.1 Atomic Theory • Ancient Greeks • Democritus – matter is discontinuous (atomos) • Aristotle – matter is continuous (hyle) • Joseph Proust (1799) • Law of Definite Proportions – samples of a compound are always composed of the same proportion by mass • Antoine Lavoisier • Law of Conservation of Mass

3. 2.1 Atomic Theory • John Dalton (1803) • Elements are made of indivisible particles (atoms) which are identical • Compounds are composed of atoms of more than one element combined in whole-number ratios • Chemical reaction involves rearranging atoms, not creating or destroying them • Law of Multiple Proportions • Elements can form different compounds by combining in many whole-number ratios

4. 2.2 Structure of the Atom • J. J. Thomson (1897) • Cathode Ray Tube Experiments • Discovered electron • e/m = -1.76 x 1011 C/kg • Plum Pudding Model • Robert Millikan (1908) • Resolved e/m ratio • Oil Drop Experiment • ee = -1.6022 x 10-19 C • me = 9.11 x 10-31 kg

5. 2.2 Structure of the Atom • Radioactivity • Wilhelm Roentgen – discovered x-rays • Antoine Becquerel – discovered radioactivity • a, b, and g radiation • The Proton • Discovered in modified CRT (canal rays) • ep = +1.6022 x 10-19 C • mp = 1.67262 x 10-27 kg • The Neutron • Discovered by James Chadwick • mn = 1.6794 x 10-27 kg

6. 2.2 Structure of the Atom • Ernst Rutherford (1910) • Attempt to find better model of the atom • Gold Foil Experiment • Performed by Geiger and Marsden • Discovered a nucleus in the atom • The Planetary Model • Nucleus composed of p+ and n • e- orbit nucleus • Held by electrostatic force

7. 2.3 Particles and Isotopes • Atomic Number (Z) • Number of protons • Mass Number (A) • Number of protons + neutrons

8. 2.3 Particles and Isotopes • Nuclear Symbols • 7← mass number Li3← atomic number • Protons = 3 • Electrons = 3 (same as p+ in neutral atom) • Neutrons = 7 – 3 = 4 • Listing Z is redundant and often not done • Isotopes • Same element, different mass number • Example: 35Cl and 37Cl

9. 2.4 Periodic Table

10. 2.4 Periodic Table • Navigating the Table • Period – row across the table • Group (or Family) – column down the table • Metals, nonmetals, and metalloids • Family Names • IA – Alkalai Metals • IIA – Alkaline Earth Metals • “B groups” – Transition Metals • VIA – Chalkogens • VIIA – Halogens • VIIIA – Noble Gases • Lanthanides and Actinides

11. 2.5 Molecules and Ions • Molecules • Neutral atoms bonded together • Diatomic molecule – contains 2 atoms • H2, N2, O2, F2, Cl2, Br2, I2 • Ions • Atom or group of atoms with a charge • Cation – positive charge • Anion – negative charge • Monatomic ion – single atom • Polyatomic ion – group of atoms

12. 2.5 Molecules and Ions • Monatomic Cations (metals) • Group IA: 1+ • Group IIA: 2+ • Group IIIA: 3+ (B and Al only) • Transition and Post-transition metals: • Ag: + Zn, Cd: 2+ • Most others can form multiple charges • Cu+ = copper (I) Cu2+ = copper (II) • Cuprous Cupric

13. 2.5 Molecules and Ions • Monatomic Anions (nonmetals) • Group VIIIA: no charge (noble gases) • Group VIIA: 1- • Group VIA: 2- • Group VA: 3- • Group IVA: 4- • H: 1- • No Roman numerals! • Change ending to -ide

14. 2.5 Molecules and Ions • Polyatomic ions • Oxoanions (element with oxygen): • First one discovered: change ending to –ate • -ite, per-, and hypo- used for other oxoanions • ClO4- perchlorateClO3- chlorateClO2- chloriteClO- hypochlorite • not all exist for every element • e.g., for N, nitrate and nitrite exist, but no others

15. 2.5 Molecules and Ions • 1+ • NH4+ ammonium • Hg22+ mercury (I) • 1- • OH- hydroxide • CN- cyanide • MnO4- permanganate • SCN- thiocyanate • HCOO- formate • CH3COO- acetate • NO3-, FO3-, ClO3-, BrO3-, IO3-

16. 2.5 Molecules and Ions • 2- • O22- peroxide • C2O42- oxalate • S2O32- thiosulfate • Cr2O72- dichromate • CO32-, SiO32-, CrO42-, SO42- • 3- • PO43-, AsO43-

17. 2.6 Chemical Formulas • Molecular Formula • Molecular Formula – true formula for a molecule • Exact number of atoms in the molecule • E. g., O2, O3, NH3 • Structural formula – shows how atoms are attached to each other • H |H – N – H • Empirical Formula – shows simplest ratio of atoms • C6H12O6 CH2O

18. 2.6 Chemical Formulas • Ionic Compounds • Two or more ions stuck together • Charges must neutralize • Sodium chloride = Na+ and Cl- • Equal charges, so NaCl • Magnesium chloride = Mg2+ and Cl- • Need 2 Cl-’s to cancel out each Mg2+ • MgCl2 • Simplify when needed, e. g., lead (IV) oxide • Try these: aluminum sulfide, iron (III) nitrate

19. 2.7 Naming Compounds • Ionic Nomenclature • Name the ions involved! • K2SO4 = potassium sulfate • Try these: K2O, Ba(ClO3)2, CaCO3, FePO4 • Molecular Nomenclature • -ide ending • Greek prefixes for numbers: • Mono-, di-, tri-, tetra- penta-, hexa-, hepta-, octa-, nona-, deca- • Try these: CO2, CCl4, H2O, SO3, UF6

20. 2.7 Naming Compounds • Acids • Produce H+ ions in water solution • Binary acids = hydro – ic acid • HCl = hydrochloric acid • Oxoacids = change ending of anion • -ate  -ic HClO3 = chloric acid • -ite  -ous HClO2 = chlorous acid • Some polyatomic anions are treated as binary • Try these: • HBr, HNO3, H2SO3, HClO4, CH3COOH

21. 2.7 Naming Compounds • Bases • Produce OH- ions in water solution • Named as any other ionic compound • NaOH = sodium hydroxide • Ammonia (NH3) also produces OH- ions in water • Hydrates • Compounds with water molecules attached • BaCl2∙ 2 H2O barium chloride dihydrate • Try this: CuSO4∙ 5 H2O