How I would study:

1. How I would study: • Look over exams • Look over review sheets • Difficulties? Work HW problems, examples from the text • Start early: where are your problem spots?

2. Chapter 1: Introduction • Dimensional analysis • Change among units (eg. feet vs. meters) • Prefixes (1 kilogram/1000 grams) • Density (d = m/v) • Scientific notation • Don’t worry about sig. figs

3. Chapter 2: Atomic Theory • Chemical formulas • Molecular formula vs. empirical formula • Naming compounds • Ionic (Table 2.3) vs. molecular • Atomic number

4. Table 2.3

5. Chapter 3: Stoichiometry • Atomic mass, molecular mass • Molar mass • Percent composition/determining empirical formulas • Chemical equations • What do coefficients tell you?

6. Chapter 3: Stoichiometry • Limiting reagents • Assume each reagent is limiting, calculate theoretical yields. Lower result? • Actual, theoretical, percent yields

7. Chapter 4: Reactions in aqueous solutions • Electrolytes • Precipitation reactions • Solubility • Molecular/ionic/net ionic equations • Acid/base reactions • Oxidation-reduction reactions • Writing half-reactions • Oxidation numbers

8. Table 4.2

9. Chapter 4: Reactions in aqueous solutions • Molarity • Gravimetric analysis • Essentially limiting reagent problems • Acid-base titrations • #mol acid = #mol base

10. Chapter 5: Gases • Ideal gas equation (PV = nRT) • Partial pressures • eg. if a gas is collected “over water,” the total pressure comes from the gas and water’s vapor pressure • Mole fraction Px = nxPT

11. Chapter 6: Energy relationships in chemical reactions • Endothermic vs exothermic • DE = q + w • q = heat (thermal energy) • w = work (w = -PDV) • Enthalpy/thermochemical equations • DH = H9products) – H(reactants) • DH of formation • Indirect vs. direct methods

12. Chapter 6: Energy relationships in chemical reactions • Calorimetry: find the energy change in a reaction (or process) qcal + qrxn = 0 qrxn = -qcal q = msDt = CDt

13. Ch 7: Electronic structure of atoms • Atomic orbitals • s, p • Electron configurations • Quantum numbers • 1s2 2s2 2p6 … • Pauli exclusion principle • Hund’s rule

14. Fig. 7.21

15. Ch 8: The Periodic Table • Isoelectronic • Effective nuclear charge • Atomic/ionic radius • Ionization energy • Electron affinity

16. Ch 9: The Covalent Bond • Lewis structures • Formal charge • Resonance • Electronegativities • Covalent/polar covalent/ionic • Bond energies DH = BE(reactants) – BE(products)

17. Ch 10: Molecular Geometry & Hybridization of Atomic Orbitals • Geometries (VSEPR model) • Hybridization • Sigma (s) vs. pi (p) bonds

18. Table 10.1 No lone pairs

19. Table 10.2 With one pairs

20. Table 10.4 Hybridization

21. Ch 12: Intermolecular forces • Boiling, melting points • Dipole: molecule must be polar • Electronegativity AND geometry • Ionic • Ion/dipole • Dipole/dipole • Hydrogen bond • Induced dipole • Dispersion

22. Ch 14: Chemical Kinetics • Rate of reaction • Decrease of reactant/increase of product • Depends on coefficients • Rate laws Rate = k[A]x[B]y • Half-life (first order) • Rate vs. temperature • Collision frequency • Activation energy • Arrhenius equation

23. Ch 15: Chemical Equilibrium • Equilibrium constant • Direction of a reaction • Q vs. Kc • Le Châtlier • Concentration (adding reactant or product) • Pressure • Temperature

24. Ch 16: Acids and Bases Ch 17: Buffers • Conjugate acid/base pairs • Water: both an acid and a base • Kw = 10-14 • Strong vs. weak acids • Ka & Kb • Calculate pH, given pKa and concentration of a weak acid • Calculate concentration of a weak acid to give a pH (given pKa)

25. Ch 18: Thermodynamics • Entropy (S): disorder • Increased S (more disorder) favorable • Decreased H (less thermal energy) favorable DG = DH - TDS