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 P x = n x P T

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

12. Chapter 6: Energy relationships in chemical reactions Calorimetry: find the energy change in a reaction (or process) q cal + q rxn = 0 q rxn = -q cal q = ms  t = C  t

13. Ch 7: Electronic structure of atoms Atomic orbitals –s, p Electron configurations –Quantum numbers –1s 2 2s 2 2p 6 … 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  H = BE(reactants) – BE(products)

17. Ch 10: Molecular Geometry & Hybridization of Atomic Orbitals Geometries (VSEPR model) Hybridization Sigma (  ) vs. pi (  ) 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. K c 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 –K w = 10 -14 Strong vs. weak acids K a & K b Calculate pH, given pK a and concentration of a weak acid Calculate concentration of a weak acid to give a pH (given pK a )

25. Ch 18: Thermodynamics Entropy (S): disorder –Increased S (more disorder) favorable –Decreased H (less thermal energy) favorable  G =  H - T  S