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7-1Use the details of modern atomic theory to explain each of the following experimental observations. Please
note that the "octet rule" is not part of modern atomic theory.
(a) The radius of the chlorine atom is smaller than the radius of the chloride ion, Cl-.
(Radii: Cl atom = 0.099 nm; Cl- ion = 0.181 nm)
(b) The first ionization energy of aluminum is lower than the first ionization energy of magnesium. (First
ionization energies: 12Mg = 7.6 ev; 13Al = 6.0 ev)
(c) The second ionization energy of sodium is about three times greater than the second ionization energy of
(d) The first ionization energy of selenium (Z = 34) is less than that of bromine (Z = 35).
(e) The first ionization energy of sodium is less than that of lithium.
The first ionization energy of sodium is less that that of neon.
7-2 The diagram shows the first ionization energies for the elements
from Li to Ne. Briefly (1 - 3 sentences) explain each of the following in
terms of atomic structure.
(a) In general, there is an increase in the first ionization energy
from Li to Ne.
(b) The first I.E. of B is lower than that of Be.
7-3 Properties of the chemical elements often show regular variation with respect to their positions in the
periodic table.
(a) Describe the acid–base character of the oxides of the elements in the third period (Na to Ar). Give
examples of one acidic oxide and one basic oxide and show with equations how these oxides react with
(b) How does the oxidizing strength of the halogens (that is, the tendency to gain an electron) vary down the
group? Account for this trend.
(c) The atomic radius of molybdenum (0.201 nm) is greater than that of chromium (0.185 nm) but almost the
same as that of tungsten (0.202 nm). Why?
(a) Why is it that irradiating a piece of a metal with short wavelength light causes an electron to be ejected,
yet irradiating the same piece of metal with longer wavelength light yields no electrons?
(b) Cite one piece of evidence which shows that particles, other than photons, have wavelengths.
(c) Describe what happens when the double-slit experiment is performed with a light beam. What does it
(d) Draw a p orbital. What does this drawing actually represent?
(e) In the Bohr model of the atom the angular momentum of electrons is quantized. What two other electron
properties are also quantized?
What about the Bohr model is incorrect according to modern quantum theory?
(g) In what way was the Bohr model of the atom an improvement over the planetary model?
7-5 The emission spectrum of hydrogen consists of several series of sharp emission lines in the ultraviolet
(Lyman series), in the visible (Balmer series), and in the infrared (Paschen series, Brackett series, etc.) regions
of the spectrum.
(a) What feature of the electronic energies of the hydrogen atom explains why the emission spectrum consists
of discrete wavelengths rather than a continuum?
(b) Draw an electronic energy level diagram for the hydrogen atom and
indicate on it the transition corresponding to the line of lowest frequency in
the Balmer series.
(c) Account for the existence of several series of lines in the hydrogen
spectrum. What distinguishes one series of lines from another?
(d) Why does Balmer's equation work for hydrogen but not for helium?
(e) What does each symbol represent in the formula
= 𝑅 (𝑚2 −
(a) Write the ground state electron configuration for an arsenic atom, showing the number of electrons in each
subshell. You may start with the previous noble gas.
(b) Give one permissible set of four quantum numbers for each of the electrons in the shell of highest energy in
an As atom when it is in its ground state.