how i would study: look over exams look over review sheets difficulties? work hw problems, examples...
Post on 12-Jan-2016
214 Views
Preview:
TRANSCRIPT
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
top related