CH123: General Chemistry - Fall 2006

Dr. Jakubowski

For the first examination, you should be able to:

  1. use dimensional analysis in problem solving.
  2. express numbers in scientific notation to the appropriate number of significant figures and with the correct units.
  3. understand the difference between homogeneous and heterogeneous mixtures and the difference between pure elements and pure compounds, giving examples.
  4. understand the principles underlying methods to separate the components of mixtures into pure substances; explain the principles of chromatography.
  5. understand the difference between chemical and physical reactions.
  6. understand early concepts in the development of chemistry and how the ideas of specific  people  influenced the development of our concept of matter:
    • La Voisier Law of mass conservation
    • Law of Constant Composition
    • Dalton - Law of Multiple Proportions; Atomic Theory
    • Law of Combining Volumes of Gases
    • Avogadro - Molecules; Avogadro's Hypothesis
    • Thompson - discovery of the electron; Plum Pudding model of the atom
    • Millikan - oil drop experiment
    • Rutherford - nuclear model of the atom
  7. explain in words and through geometric symbols for atoms/molecules (a nanoscopic representation) how Dalton's Atomic Theory and Avogadro's Hypothesis can explain the Laws of Mass Conservation, Constant Composition, Multiple Proportions, and Combining Volumes (macroscopic observations)
  8. describe the differences between laws and theories.
  9. identify in an experiment the independent and dependent variables, and those held fixed.
  10. give an example of a scientific theory which must be falsifiable, and a non-scientific theory, one that is not.
  11. explain the difference between atoms, molecules, ions, and molecular (polyatomic) ions. Give examples of each.
  12. name ionic compounds and simple binary compounds of nonmetals.
  13. describe how magnetic and electric fields affects the movement of charged particles.
  14. describe the atomic and molecular mass scales.
  15. define isotopes and allotropes and give examples
  16. use dimensional analysis to convert between moles, molar mass, and number of particles.
  17. explain periodic law embodied in the empirical arrangement of elements in the periodic table
  18. describe how the arrangement of electrons into shells explains the organization of the periodic table
  19. explain a graph showing ionization energies of the element as a function of atomic number is consistent with a shell-like structure for the arrangement of electrons around nuclei
  20. describe the differences between ionic, polar covalent, and nonpolar covalent bonds
  21. Identify ionic, polar covalent, and nonpolar covalent bonds in chemical structures
  22. describe the meaning of electronegativity
  23. write molecular and condensed formulas for molecules
  24. draw Lewis structures for molecules and molecular ions
  25. determine formal charge and partial charges in any Lewis structure
  26. know the formula and Lewis structures of common molecular ions
  27. draw a mechanism with Lewis structures using curved arrows to show electron flow in the forming and breaking of bonds in the formation of products
  28. write formulas for salts containing molecular ions
  29. determine the formal charge on any atom/ion
  30. explain how molecular ions often arise from the reaction of water with oxyacids
  31. describe atoms, molecules and their properties using macroscopic, nanoscopic, symbolic, and mathematical representations
  32. understand the relationships among mole, molecular or atomic mass, and molar mass
  33. write, balance, and interpret simple chemical equations,
  34. define molarity, understand the relationships among moles, volume, mass, and molarity,.
  35. calculate:
  • average atomic weight of an element given abundancy of isotopes
  • % composition of a given element in a compound
  • number of particles, moles, or mass, given two of the three quantities
  • empirical formula of a compound
  • molecular formula given the empirical formula and molecular weight
  • values of molarity, volume, mass given any two.
  • stoichiometric relations (moles, grams, or molarity) under a variety of conditions including excess/limiting reagents and in solution.

Note:  Questions based on the lab might be expected.