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?

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

Chapter 2: Atomic Theory

• Chemical formulas– Molecular formula vs. empirical formula– Naming compounds

• Ionic (Table 2.3) vs. molecular

– Atomic number

Table 2.3

Chapter 3: Stoichiometry

• Atomic mass, molecular mass

• Molar mass

• Percent composition/determining empirical formulas

• Chemical equations– What do coefficients tell you?

Chapter 3: Stoichiometry

• Limiting reagents– Assume each reagent is limiting, calculate

theoretical yields. Lower result?– Actual, theoretical, percent yields

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

Table 4.2

Chapter 4: Reactions in aqueous solutions

• Molarity

• Gravimetric analysis– Essentially limiting reagent problems

• Acid-base titrations– #mol acid = #mol base

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 fractionPx = nxPT

Chapter 6: Energy relationships in chemical reactions

• Endothermic vs exothermic

• E = q + w– q = heat (thermal energy)– w = work (w = -PV)

• Enthalpy/thermochemical equations– H = H9products) – H(reactants)– H of formation

• Indirect vs. direct methods

Chapter 6: Energy relationships in chemical reactions

• Calorimetry: find the energy change in a reaction (or process)qcal + qrxn = 0

qrxn = -qcal

q = mst = Ct

Ch 7: Electronic structure of atoms

• Atomic orbitals– s, p

• Electron configurations– Quantum numbers– 1s2 2s2 2p6 …

• Pauli exclusion principle

• Hund’s rule

Fig. 7.21

Ch 8: The Periodic Table

• Isoelectronic

• Effective nuclear charge– Atomic/ionic radius– Ionization energy– Electron affinity

Ch 9: The Covalent Bond

• Lewis structures

• Formal charge

• Resonance

• Electronegativities– Covalent/polar covalent/ionic

• Bond energies H = BE(reactants) – BE(products)

Ch 10: Molecular Geometry & Hybridization of Atomic Orbitals

• Geometries (VSEPR model)

• Hybridization

• Sigma () vs. pi () bonds

Table 10.1

No lone pairs

Table 10.2

With one pairs

Table 10.4

Hybridization

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

Ch 14: Chemical Kinetics

• Rate of reaction– Decrease of reactant/increase of product– Depends on coefficients

• Rate lawsRate = k[A]x[B]y

• Half-life (first order)• Rate vs. temperature

– Collision frequency– Activation energy– Arrhenius equation

Ch 15: Chemical Equilibrium

• Equilibrium constant

• Direction of a reaction– Q vs. Kc

• Le Châtlier– Concentration (adding reactant or product)– Pressure– Temperature

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)

Ch 18: Thermodynamics

• Entropy (S): disorder– Increased S (more disorder) favorable– Decreased H (less thermal energy) favorable

G = H - TS