Chemistry Final Review

  1. Matter, measurement, problem solving
  • The scientific method
  • Classification of matter
  • Chemical and physical changes
  • Metric units
  • Temperature scales
  • Dimensional analysis
  • Accuracy
  • Precision
  • Significant figures
  1. Atoms and elements
  • Dalton’s atomic theory
  • Subatomic particles
  • Isotopes
  • Ions
  • Mass number
  • Atomic number
  • Using the periodic table to predict reactivity and ion formation
  • Weighted average atomic mass
  • The mole
  • Molar mass
  1. Molecules, compounds, and chemical equations

  • Ionic bonds

  • Covalent bonds

  • Ionic compounds

  • Molecular compounds

  • Formulas and names of ionic compounds

  • Formulas and names of covalent compounds

  • Formula and molar mass of compounds

  • Percent composition

  • Determination of empirical and molecular formulas

  • Types of chemical reactions

  • Balancing chemical equations

  1. Chemical quantities and aqueous reactions
  • Reaction stoichiometry
  • Theoretical yield
  • Limiting and excess reactants
  • Percent yield
  • Molarity
  • Dilution
  • Molarity in stoichiometry
  • Electrolytes and nonelectrolytes
  • Strong and weak electrolytes
  • Solubility of ionic compounds
  • Identifying acids and bases using Bronsted-Lowry definition
  • Precipitation and gas formation reactions
  • Molecular, ionic, and net ionic equations
  1. Gases
  • Pressure and pressure measurement
  • Common units for pressure
  • Kinetic molecular theory
  • The gas laws
  • Density of gases
  • Effusion and diffusion
  • Graham’s law
  • Real gases
  1. Thermochemistry
  • Potential and kinetic energy
  • State functions
  • System and surroundings
  • First law of thermodynamics
  • Signs of heat and work
  • Internal energy
  • Calorimetry
  • Enthalpy
  • Hess’ law
  • Formation reactions
  • Heat of formation
  • Heat of reaction
  1. Quantum mechanical model of the atom
  • Electromagnetic radiation
  • Wavelength
  • Frequency
  • Energy
  • velocity of propagation

plane of polarization

  • Diffraction
  • Interference
  • Dispersion,
  • Wave and particle nature of electrons
  • Absorption and emission of radiation
  • Quantum numbers
  • Orbitals
  1. Periodic properties of elements:
  • Orbital diagrams: s,p,d
  • Electron configurations: fill in orbital diagrams
  • Valence electrons: surrounding electrons
  • Core electrons: central electrons
  • Periodic trends in atomic radius: bigger going left and down
  • Ionization affinity:
  • Electron affinity:
  • Ionic charge:
  1. Chemical bonding: the Lewis Model
  • Lattice energy: Lattice energies become less exothermic (less negative) with increasing ionic radius and more exothermic (more negative) with increasing magnitude of ionic charge
  • Resonance structures: Two or more Lewis structures have the same skeletal formula, but different electron arrangements
  • Exceptions to the octet rule: odd-electron species, incomplete octets, expanded octets
  • Odd-Electron species: Molecules with odd number of electrons in their Lewis structures are free radicals
  • Incomplete octets: The most important of these is boron, which forms compounds with only six electrons around B, rather than eight.
  • Expanded octets: Elements in the third row of the periodic table and beyond often exhibit expanded octets of up to 12 (and occasionally 14) electrons
  • Bond energies: Also known as bond enthalpy or bond dissociation. The energy required to break 1 mole of the bonds in the gas phase.
  • Bond energies are always ____ because it takes energy to break a bond (endothermic): positive
  • Bonds break → ____: endothermic (positive bond energy)
  • Bond enthalpy equation: Bonds broken + bonds formed
  1. Chemical bonding: molecular shapes, hybridization, and molecular-orbital theory
  • VSPER: Valence Shell Electron Pair Repulsion Theory (VSEPR): Negatively charged groups are attracted to the nucleus and repel each other. This repulsion determines the molecular geometry (shape) of a compound
  • Five basic molecular shapes: Linear geometry, trigonal planar geometry, tetrahedral geometry, trigonal bipyramidal geometry, octahedral geometry
  • Order of repulsion: lone pair- lone pair is most repulsive → lone pair-bonding pair → Bonding pair-bonding pair is the least repulsive
  • The bond angles get progressively ___ as the number of lone pairs on the central atom increases from zero in CH4 to one in NH3 to two in H2O: smaller
  • Use VSPER theory to predict molecular and electron geometries and bond angles: (slide 3 and 4 on Chapter 11)
  • Covalent Bond Polarity: If the atoms have identical electronegativity, the bond is purely covalent or nonpolar and the electrons are evenly shared.
  • Ionic: If there is a large electronegativity difference between atoms in a bond between a metal and a nonmetal, electrons are transferred and the bond is ionic
  • Polar covalent: If the two nonmetals have different electronegativity, the bond is polar covalent and the electrons are unevenly shared
  • Paramagnetism: When there are unpaired electrons
  • Diamagnetism: All electrons are paired
  • Sigma bond (σ bond): Single bond
  • Double bond: 1 sigma + 1 pi
  • Triple bond: 1 sigma + 2 pi bonds
  • Linear: The dipole moments of two identical polar bonds pointing in opposite directions cancel. The molecule is nonpolar.
  • Bent: The dipole moments of two polar bonds with an angle of less than 180 between them do not cancel. The resultant dipole moment vector is shown in red. The molecule is polar
  • Trigonal Planar: The dipole moments of three identical polar bonds at 120 from each other cancel. The molecule is nonpolar
  • Tetrahedral: The dipole moments of four identical polar bonds in a tetrahedral arrangement (109.5 from each other) cancel. The molecule is nonpolar
  • Trigonal pyramidal: The dipole moments of three polar bonds in a trigonal pyramidal arrangement do not cancel. The resultant dipole moment vector is shown in red. The molecule is polar