Download Final Review

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 fraction
Px = nxPT
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
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
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
DH = BE(reactants) – BE(products)
Ch 10: Molecular Geometry & Hybridization
of Atomic Orbitals
• Geometries (VSEPR model)
• Hybridization
• Sigma (s) vs. pi (p) 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 laws
Rate = 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
DG = DH - TDS