Chapter 6 and 7 Reading Guide
Electronic Structure of Atoms and Periodicity
Section 6.1
Define the following terms:
wavelength
frequency
What is the relationship between wavelength and frequency?
Look at Figure 6.4…. What wavelengths (short or long) and what frequencies (high or low) are characteristic of highenergy radiation?
What is the range of wavelengths of visible light from RED to VIOLET? Which color has greater energy?
In the equation c=λν, what does each variable represent? Since the speed of light is 3.00 x 108 m/s, what units must
be used for wavelength and for frequency?
What is a Hertz?
Section 6.2
Define the following terms:
blackbody radiation:
photoelectric effect:
Emission Spectrum:
quantum:
photon:
How is the work done by Max Planck used by Albert Einstein in his explanation of the photoelectric effect?
Section 6.3
WARNING: Do NOT get scared by all the equations!!
Define the following terms:
continuous spectrum:
line spectrum:
ground State:
excited State:
What is the major flaw in the Rutherford/Bohr model of the atom (“microscopic solar system”)?
How do spectral lines support the three postulates of Bohr’s model?
What is the major limitation of the Bohr model?
Section 6.4
Define the following terms:
momentum:
matter waves:
uncertainty Principle:
Why does the uncertainty principle not apply to macroscopic objects?
Section 6.5
Define the following terms:
wave mechanics (quantum mechanics):
orbital:
principal quantum number (n):
What are its possible values?
angular momentum quantum number (aka azimuthal quantum number) (l):
What are its possible values?
magnetic quantum number (ml):
What are its possible values?
How many subshells can exist in any PEL?
What value will l have for s subshells?
How many values can ml have for s subshells?
for p?
for d?
for p?
for f?
for d?
Section 6.6
Define the following terms:
Radial probability function:
Node:
What is the shape of an s orbital? What is the difference between a 1s, 2s or 3s orbital?
What is the shape of a p orbital? What are the designations for the three p orbitals?
Section 6.7
Define the following terms:
degenerate:
electron spin:
spin magnetic quantum number:
What are its possible values?
for f?
Pauli exclusion principle:
What does the statement “In a many-electron atom, for a given value of n, the energy of an orbital increases with
increasing value of l” mean?
What is the effect of the spin difference between two electrons in the same orbital?
Section 6.8
Define the following terms:
electron configuration:
Hund’s rule:
condensed electron configuration:
core electrons:
valence electrons:
What limitation does Hund’s rule place on where electrons are located in an electron configuration?
Section 6.9
Define the following term:
representative/main-group elements:
Look at Figure 6.30. For each period , define where electrons are added in each block (s, p, d and f) in terms of the
principle quantum number “n”.
What determines the number of elements in each period of a block?
What two subshells can contain valence electrons?
Draw an orbital filling diagram for the 3d and 4s sublevels in Cr and in Cu, using information in Figure 6.31. Why are
these anomalous configurations more stable for the atom than the expected configuration?
Section 7.1
What two scientists are credited with the development of the modern periodic table? Why is one more widely
recognized or his work than the other?
What important contribution did Henry Moseley make to the periodic table?
Section 7.2
Define the following terms:
effective nuclear charge(Zeff):
screening constant:
Why is effective nuclear charge smaller than actual nuclear charge?
Why does the effective nuclear charge increase as the atomic number increases across a period?
Compare the effective nuclear charge in an s sublevel and a p sublevel of the same PEL.
Why is there little change in effective nuclear charge as atomic number increases within a group?
Be sure to read “A Closer Look” on page 260. Follow the diagram for a clear understanding of effective nuclear
charge. This concept is critical to understanding periodic trends.
Section 7.3
Define the following terms:
non-bonding atomic radius (van der Waal’s radius):
bonding atomic radius (covalent radius):
isoelectronic:
State the trend and justification for atomic radii within a group:
State the trend and justification for atomic radii within a period:
Why are cations smaller than their parent atoms? (DON’T just re-write what the text says… interpret it in simple terms!)
Why are anions larger than their parent atoms?
What determines the trend in radius of isoelectronic ions?
Section 7.4
Define the following terms:
ionization energy:
first ionization energy:
second ionization energy:
Why does each successive ionization energy for a particular atom increase?
Account for the sudden large increases in ionization energies in blue in Table 7.
State the trend and justification for first ionization energy within a group:
State the trend and justification for first ionization energy within a period:
Account for the decrease in first ionization energy from Be to B.
Account for the decrease in first ionization energy from N to O.
In the transition metals, which electrons are removed first when forming an ion? Why?
Section 7.5
Define the following term:
electron affinity:
What is the difference between ionization energy and electron affinity?
Why do the halogens have the most negative electron affinity?
Why do the noble gases have such highly positive electron affinities?
What general factors determine the sign and magnitude of electron affinity for a particular atom?
Section7.6
Define the following term:
metallic character:
List some characteristics of metals:
List some characteristics of non-metals:
List some characteristics of metalloids:
The rest of the chapter covers information on specific elements that you may find interesting, but is not a major
emphasis for the AP test!