STUDY GUIDE - TEST 2:
CHAPTERS 4-6

CH123: General Chemistry - Fall 2006

Dr. Jakubowski

For the second examination, you should be able to:

1. Understand the relationships among mole, molecular or atomic weight, and mass, be able to write, interpret chemical equations, and solve quantitative problems about chemical reactions given descriptions of the reactions.
2. Identify the molecular/ionic species in water when a substance is added to water.
3. Write the molecular, ionic, and net ionic equations for the reactions of strong electrolytes in water.
4. Identify precipitation reactions, and predict the products of precipitation reactions based on solubility rules.
5. Define acid and a base, identify acid/base reactions, and predict the products of the reactions.
6. Know the formula and Lewis structures of the common strong acids (nitric, hydrochloric, sulfuric), the weak acid (acetic acid), the strong bases (oxide, hydroxide) and the weak base ammonia.
7. Rank the relative strength of acids and bases based on writing chemical equations, and based on comparisons of the relative stability of the conjugate base produced.
8. Define oxidation, reduction, oxidizing agent, reducing agent
9. Determine/assign oxidations numbers of all atoms in a substance and understand how to use oxidation numbers in identifying and analyzing redox reactions and in comparing the relative strength of a series of molecules with the same atoms
10. Balance redox equations usingthe half-reaction method in acidic or basic solution .
11. Do stoichiometry problems based on redox titrations.
12. Given tables showing the relative strength of oxidizing/reducing agents and acids/bases, predict likely pairs of stronger oxidizing and reducing agents, and of stronger acids and bases, that would likely react to form weaker pairs.
13. Describe interference (constructive, destructive), and diffraction, and their use in understanding the properties of waves and particles.
14. Explain the photoelectric effect and how Einstein's interpretation of it altered our view of light. 
15. Mathematically interpret the graphs of E_kin vs frequency (υ) in the photoelectric effect
16. Explain, with diagram and words, the Rutherford, Bohr, and Schrodinger's models of the atom, and describe the successes and limitations of each of these theories.  (For example, explain how each of the above theories explains (or cannot explain) atomic absorption and emission spectra.
17. Discuss the meaning and differences in the probability amplitude,ψ, the probability amplitude, ψ², the probability amplitude as depicted in electron cloud maps, and the radial distribution function.
18. Describe the significance of the quantum numbers n, l, m_l, and m_s.
19. Draw  for 1s, 2s  and 2p orbitals:  graphs of electron probability  vs distance from the nucleus, and cartoon pictures representing the electron density and the shapes of the orbitals.
20. Write the quantum number for the electrons in specific orbitals
21. Describe orbitals and the relative energy of the orbitals in hydrogen and polyelectronic atoms. 
22. Explain the meaning of core electrons, valence electrons, shielding, penetration, core charge, and Z_eff.
23. Explain and use the Aufbau principles, Pauli Exclusion Principle and Hund's rule to write the electron configuration of atoms and ions (1s²2s²sp⁶... etc) in the periodic table, and to draw orbital diagrams (placing arrows in boxes representing the orbitals) for these configurations.
24. Explain the term degeneracy and why for the hydrogen atom the s, p, d, and f orbitals in the same shell (same n value) are degenerate, but for other elements, the orbitals in a given shell are not.
25. Explain how the shape of the present Periodic Table can be accounted for by quantum theory.
26. Explain periodic trends in ionization energy, electron affinity, atom and ion size, core charge and chemical reactivity for the elements in the Periodic Table with respect to electron configuration and Zeff.
27. Draw energy diagrams showing the relationships between the energy of electrons in an atom, and the ionization energy. 
28. Define lattice energies and predict which ionic compounds have lower and higher lattice energies.
29. Given energy values like ionization energy, and electron affinities, use a Born-Haber thermodynamic cycle to calculate lattice energy.
30. Explain, and manipulate the following formulas:
    • c =ul
    • p = mv ; p_ang  =  mvr
    • λ= h/p
    • E = hυ
    • E = mc²
    • F=ma
    • ΔpΔx > h/4π
    • E_kin = mv²/2
    • E_pot = -ke²/r (electrostatic)  E_pot = mgh (gravitational potential)
    • F_coul. = kQ₁Q₂/r²
    • E_tot = mv²/2  - ke²/r  = -ke²/2r   => The Bohr equation for H assuming a stable atom with a classical inward Coulombic force balanced by kinetic energy
    • Bohr quantum postulate:  p_ang  =  mvr = nh/2π
    • Bohr equations for hydrogen:   E=-B/n² = -1312/n²  (kJ/mol e) = -2.18 x 10^-18/n² J/e