Chem 2 AP Homework #8-2: Problems pg. 339 #49, 50, 51, 53, 54, 55, 59, 60, 61, 64, 127
49
Define ionization energy. Why are ionization energy measurements usually made when atoms
are in the gaseous state? Why is the second ionization energy always greater than the first for
any element?
Ionization energy is the minimum energy required to remove an electron from a gaseous atom in its ground
state. The gas phase is used to eliminate influences of neighboring atoms such as intermolecular forces.
The second ionization energy is always larger than the first because the repulsion between electrons is
less while the nuclear charge remains constant, increasing the attraction of the electron to the nucleus.
50
Sketch the outline of the periodic table and show group and
period trends in the first ionization energy of the elements.
What types of elements have the highest ionization energies
and what types the lowest?
The noble gases have the highest ionization energies and the alkali
metals have the lowest activation energies.
51
Arrange the following in order of increasing first ionization energy:
Ionization energy increases across a row of the periodic table and increases up a column or group (note:
the trend is not an explanation). The correct order of increasing ionization energy is:
Cs < Na < Al < S < Cl
53
Use the third period of the periodic table as an example to illustrate the change in first
ionization energies of the elements as we move from left to right. Explain the trend.
Apart from the small irregularities, the ionization energies of elements in a period increase with
increasing atomic number. We can explain this trend by referring to the increase in effective nuclear
charge from left to right. A larger effective nuclear charge means a more tightly held outer electron, and
hence a higher first ionization energy. Thus, in the third period, sodium has the lowest and neon has the
highest first ionization energy.
54
In general, ionization energy increases from left to right across a period. Aluminum, however,
has a lower ionization energy than magnesium. Explain.
The Group 3A elements (such as Al) all have a single electron in the outermost p subshell, which is well
shielded from the nuclear charge by the inner electrons and the ns2 electrons. Therefore, less energy is
needed to remove a single p electron than to remove a paired s electron from the same principal energy
level (such as for Mg).
55
The first and second ionization energies of K are 419 kJ/mol and 3052 kJ/mol, and those of Ca
are 590 kJ/mol and 1145 kJ/mol. Compare their values and comment on the differences.
To form the +2 ion of calcium, it is only necessary to remove two valence electrons. For potassium,
however, the second electron must come from the atom's noble gas core which accounts for the much
higher second ionization energy. Would you expect a similar effect if you tried to form the +3 ion of
calcium?
59
(a) Define electron affinity. (b) Why are electron affinity measurements are made with gaseous
atoms? (c) Ionization energy is always a positive quantity, whereas electron affinity may be
either positive or negative. Explain.
(a) Electron affinity is the negative of the energy change that occurs when an electron is accepted by an
atom in the gaseous state to form an ion.
(b) As with ionization energy, the gaseous state is used to eliminate influences from neighboring atoms.
(c) The values can be positive, indicating a favorable process with a more stable product, or negative,
which indicates an unfavorable process with a less stable product.
2
HOMEWORK #8-3 ANSWER KEY
60
Explain the trends in electron affinity from aluminum to chlorine (see Table 8.3).
There is a trend to increase the electron affinity moving left to right in a period since the alkali metals
have little tendency to accept electrons, but the halogens are much more stable as a single anion. Al has
a slightly higher EA than Mg since Mg’s electron would enter an empty, higher-energy p orbitals, slightly
destabilizing Mg; Si– is slightly more stable than Al–, but P– is less stable since the additional electron
must overcome electron repulsion from entering a half-filled p orbital (recall that P’s electron
configuration is 3p3). O– is more stable, and Cl– is the most stable since the additional electron
experiences an increased coulombic attraction to the nucleus without substantial increase in electronelectron repulsion or shielding by core electrons.
61
Arrange the elements in each of the following groups in increasing order of the most positive
electron affinity:
(a) K < Na < Li
(b) I < Br < F < Cl
64
Explain why alkali metals have a greater affinity for electrons than alkaline earth metals.
Alkali metals have a valence electron configuration of ns1 so they can accept another electron in the ns
orbital. On the other hand, alkaline earth metals have a valence electron configuration of ns2. Alkaline
earth metals have little tendency to accept another electron, as it would have to go into a higher energy p
orbital.
127
Why do elements that have high ionization energies also have more positive electron affinities?
Which group of elements would be an exception to this generalization?
Both ionization energy and electron affinity are affected by atomic size − the smaller the atom, the
greater the attraction between the electrons and the nucleus. If it is difficult to remove an electron from
an atom (that is, high ionization energy), then it follows that it would also be relatively favorable to add
an electron to the atom (large electron affinity).
Noble gases would be an exception to this generalization since putting an electron into a Noble gas
would require putting it into a well shielded orbital at much higher energy, an unfavorable process.