AP Chemistry
Chapter 6 Outline for Concepts to Know
6.1 Wave Nature of Light
 Basic anatomy and vocabulary of a wave (wavelength, frequency, amplitude)
 Relationship between speed, wavelength and frequency
 c= 3.00x108 m/s
 order of categories of electromagnetic radiation
 order and approximate range of visible radiation 4 – 8(x10-7) m
6.2 Quantized Energy and Photons
 Concept of smallest unit of light energy as photon – having properties of both particles and waves
 Quantum as smallest possible packets or quantities of energy
 Photoelectric effect – effect of changing frequency? Changing intensity?
 Calculations using Planck’s constant, such as E=hv, will NOT be tested
6.3 Line Spectra and the Bohr Model
 Emission spectrum of hydrogen (see p. 225) is due to energy transitions of the single electron of hydrogen
being excited to higher energy levels and then falling back down, emitting specific wavelengths of light.
 Line spectra for other elements are generally more complex, but are all due to various energy transitions of
electrons moving from excited to ground (or less excited) levels
 Calculations of energy states of hydrogen and other atoms will NOT be tested
 Bohr model will not be specifically tested
6.4Wave Behavior of Matter
 Conceptual understanding that all matter has an energy- and a wave-equivalent, but that this only has
relevance for chemistry if considering very small particles (such as electrons)
 Calculations of matter waves (equation 6.8) will not be tested
 Uncertainty principle as being a limit to which the momentum and position of an object can be determined –
relevance for chemistry? Electrons can be described as existing within a probability cloud around a nucleus (or
between nuclei), but can never be pinpointed.
6.5 Quantum Mechanics and Atomic Orbitals
 Understand orbitals as areas of high probability for where an electron may be found, and that orbitals have
increasing energy in predictable ways
 Know how to identify electrons around atoms with principle energy levels (n levels), sub-levels (s,p,d,f), orbitals
within sublevels (1,3,5,7), and spin (+/- ½) {discussed in 6.7}
 Numeric values of l, ml, ms will not be tested. (i.e. Skip problems like sample problem 6.6)
6.6 Representations of orbitals
 Know basic shapes of s,p,d, orbitals.
 Identify meaning of figure 6.18 Radial Probability function for subsequent occurrences of of an orbital in higher
n levels
6.7 Many-electron atoms
 Note the effect of more than one electron on energy of sublevels (Compare figure 6.17 with 6.24)

Spin and Pauli exclusion principle for placement of electrons
6.8 – 6.9 Electron configurations and the periodic table
 Be able to complete electron configurations and orbital diagrams for ground-state multi-electron atoms or ions
 Degeneracy, electron placement and Hund’s rule
 Be able to write “short-cut” condensed form electron configurations; core vs. valence lectrons
 Dealing with electrons in transition metals and rare earths
 Anomalous electron configurations: note for chromium and copper – be able to offer explanation if oresented
with other similar