CHAPTER 7 Problems: 5, 6, 12, 18, 27, 28, 30, 34, 42, 46, 51, 58 a-e, 60 a-e, 65, 76, 82,
84, 96, 116
5) Classify each of the following elements as a metal, a nonmetal, or a metalloid.
As
metalloid
Cl
nonmetal (halogen)
Xe
nonmetal (noble gas)
Ba
metal (alkaline earth metal)
Fe
metal (transition metal)
P
nonmetal
Li
metal (alkali metal)
I
nonmetal (halogen)
B
metalloid
Si
metalloid
6) Compare the physical and chemical properties of metals and nonmetals.
Metals are good conductors of heat and electricity; nonmetals are poor conductors
of heat and electricity
Metals have a shiny and metallic appearance; solid nonmetals are usually dull in
appearance
Metals tend to form cations in ionic compounds; nonmetals tend to form anions in
ionic compounds
Metals are ductile and malleable; nonmetals are not (in fact, solid nonmetals are
often brittle)
Metals tend to form basic oxides; nonmetals tend to form acidic oxides
12) What are valence electrons? For main group elements the number of valence
electrons is equal to its group number. Show that this is true for the following elements.
Valence electrons are the outermost electrons in an atom. For the main group
elements the valence electrons are all of the electrons found in orbitals with the highest
value for n. For transition metals (where the concept of valence electrons has less
meaning) the valence electrons are all electrons beyond the noble gas configuration
a) Al
b) Sr
c) K
d) Br
e) P
f) S
g) C
[Ne] 3s2 3p1
[Kr] 5s2
[Ar] 4s1
[Ar] 4s2 3d10 4p5
[Ne] 3s2 3p3
[Ne] 3s2 3p4
1s2 2s2 2p2
3 valence electrons (group 3)
2 valence electrons (group 2)
1 valence electron (group 1)
7 valence electrons (group7)
5 valence electrons (group 5)
6 valence electrons (group 6)
4 valence electrons (group 4)
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18) Group the following electron configurations in pairs that would represent elements
with similar chemical properties:
a) 1s2 2s2 2p6 3s2
b) 1s2 2s2 2p3
c) 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6
d) 1s2 2s2
e) 1s2 2s2 2p6
f) 1s2 2s2 2p6 3s2 3p3
Elements with the same valence electron configuration are expected to have
similar chemical properties, and so
a+d
c+e
2 valence electrons
8 valence electrons
b+f
5 valence electrons
27) How does atomic radius change
a) From left to right across a period?
Decreases (because Zeff increases, and so electrons are pulled closer to the
nucleus.)
b) From top to bottom in a group?
Increases (because all atoms have approximately the same value for Zeff, and so
the larger the value for n for the valence electrons the larger the atom.)
28) Define ionization energy. Explain why ionization energy measurements are usually
made when atoms are in the gas state. Why is the second ionization energy always
greater than the first ionization energy for any element?
Ionization energy is the minimum energy required to remove an electron from an
atom, molecule, or ion in the gas phase. Ionization energy is measured in the gas state so
that there are no surrounding atoms or ions that might affect the value for the ionization
energy. Since the easiest electron to remove will always be removed first (since
ionization energy is the minimum energy required to remove an electron) it follows that
IE1 < IE2 < IE3 < … where IE1 is the first ionization energy, and so forth.
30) Define electron affinity. Explain why electron affinity measurements are made with
gaseous atoms. Ionization energy is always a positive quantity, whereas electron affinity
may be either positive or negative. Explain.
Electron affinity is the minimum (in magnitude) energy required to add an
electron to an atom, molecule, or ion. Like ionization energy, electron affinity is
measured in the gas phase to make sure that the experimental value for ionization energy
is not affected but surrounding atoms or ions. For atoms that are more stable adding an
electron (like the halogens, which have space in their valence p orbitals for one more
electron) electron affinity is negative, but for atoms which do not have a desire to add an
electron (like the noble gases, which have a filled valence shell) the electron affinity will
be positive.
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34) On the basis of their position in the periodic table, select the atom with the larger
atomic radius in each of the following pairs:
a) Na, Si
Na > Si (size decreases from left to right)
b) Ba, Be
Ba > Be (size increases from top to bottom)
c) N, F
N > F (size decreases from left to right)
d) Br, Cl
Br > Cl (size increases from top to bottom)
e) Ne, Kr
Kr > Ne (size increases from top to bottom)
42) Arrange the following in order of increasing first ionization energy: Na, Cl, Al, S, Cs.
First ionization energy increases from left to right, and decreases from top to
bottom. Therefore
Na < Al < S < Cl
elements in same row
Cs < Na
elements in same group
We can combine these to get Cs < Na < Al < S < Cl
46) The first and second ionization energies for K are 419. kJ/mol and 3052. kJ/mol, and
for Ca are 590. kJ/mol and 1145. kJ/mol. Compare their values and comment on the
differences.
There is a much larger jump in second ionization energy for K than for Ca. We
can use the electron configurations involved to see why this occurs
K
K+
1s2 2s2 2p6 3s2 3p6 4s1
1s2 2s2 2p6 3s2 3p6
Ca
Ca+
1s2 2s2 2p6 3s2 3p6 4s2
1s2 2s2 2p6 3s2 3p6 4s1
For Ca and Ca+ the electron being removed sees an effective nuclear charge of
approximately +2, and so there is not a large change in ionization energy. For K, the
electron is being removed from the 4s orbital (Zeff = +1) while for K+ the electron is being
removed from a 3p orbital (Zeff = +9). Because Zeff is so much larger for removal of the
second electron from potassium, there is a big jump in the second ionization energy
relative to the first ionization energy.
51) Explain why alkali metals have a greater affinity for electrons than alkaline earth
metals.
For the alkali metals there is space in the valence s orbital to hold one more
electron. For the alkaline earth metals the valence s orbital is filled, and so an additional
electron would have to go into a p or d orbital, which are higher in energy. So the
electron affinity for the alkaline earth metals is expected to be positive.
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58) Write the ground state electron configurations of the following ions:
a) Li+ 1s2 = [He]
b) H- 1s2 = [He]
c) N3- 1s2 2s2 2p6 = [Ne]
d) F- 1s2 2s2 2p6 = [Ne]
e) S2- 1s2 2s2 2p6 3s2 3p6 = [Ar]
60) Write the ground state electron configurations of the following transition metal ions.
Recall that for transition metal cations we start with the configur4ation for the
atom, and then remove electrons from the orbital with the highest n (or highest  for
orbitals with the same value for n)
a) Sc3+
b) Ti4+
c) V5+
d) Cr3+
e) Mn2+
Sc
Ti
V
Cr
Mn
[Ar] 4s2 3d1
[Ar] 4s2 3d2
[Ar] 4s2 3d3
[Ar] 4s1 3d5
[Ar] 4s2 3d5
so Sc3+
so Ti4+
so V5+
so Cr3+
so Mn2+
[Ar]
[Ar]
[Ar]
[Ar] 3d3
[Ar] 3d5
65) Indicate which one of the two species in each of the following pairs is smaller:
a) Cl or ClCl is smaller (same nuclear charge + smaller number of
electrons)
+
b) Na or Na
Na+ is smaller (same nuclear charge + smaller number of
electrons)
c) O2- or S2O2- is smaller (same group, same valence electron configuration, and n is larger for the valence electrons in S2-)
2+
3+
d) Mg or Al
Al3+ is smaller (ions have the same number of electrons,
so the larger the charge of the nucleus the smaller the ion)
e) Au+ or Au3+
Au3+ is smaller (same nuclear charge + smaller number of
electrons)
76) As a group the noble gases are very stable chemically. Why?
The effective nuclear charge for the valence electrons for a noble gas is Zeff =
+8(except for helium, where it is +2) and so the noble gases do not want to give up
electrons. Since adding an electron would place it in a higher energy s orbital (which
would have Zeff approximately equal to zero) the noble gases also have no tendency to
add electrons. It is therefore difficult to form chemical substances with the noble gases.
Or, in short - the noble gases have a stable electron configuration and so do not
want to change it.
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82) State whether each of the properties of the main group elements generally increases
or decreases a) from left to right across the period
b) from top to bottom within a group
metallic character - generally decreases from left to right; generally increases from top to
bottom.
atomic size - generally decreases from left to right; generally increases from top to
bottom
ionization energy - generally increases from left to right; generally decreases from top to
bottom
acidity of oxides - generally increases from left to right; generally decreases from top to
bottom
84) Write equations representing the following processes:
a) The electron affinity of SS-(g) + e-  S2-(g)
b) The third ionization energy for titanium
Ti2+(g)  Ti3+(g) + ec) The electron affinity of Mg2+
Mg2+(g) + e-  Mg+(g)
d) The ionization energy for O2O2-(g)  O-(g) + e-
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96) Predict the products of the reaction of the following oxides with water, write the
balanced reaction, and specify whether the oxides are acidic, basic, or amphoteric.
a) Na2O
Na2O(s) + H2O(l)  2 NaOH(aq)
basic
b) BaO
BaO(s) + H2O(l)  Ba(OH)2(aq)
basic
c) CO2
CO2(g) + H2O(l)  H2CO3(aq)
acidic
d) N2O5
N2O5(g) + H2O(l)  2 HNO3(aq)
acidic
e) P4O10
P4O10(s) + 6 H2O(l)  4 H3PO4(aq)
acidic
f) SO3
SO3(g) + H2O(l)  H2SO4(aq)
acidic
116) Consider the halogens chlorine, bromine, and iodine. The melting and boiling
points for chlorine are - 101.5 C and - 34.0 C, and are 113.7 C and 184.3 C for
iodine, respectively. Thus chlorine is a gas and iodine is a solid at room temperature and
pressure.
Predict the melting and boiling points for bromine. Compare your results with the
actual values (found, for example, at www.webelements.com).
Being young and naïve I would predict that the normal melting and boiling point
temperatures for Br should be the average of the values for Cl and I. Therefore
predicted
actual (from the CRC Handbook)
normal melting pt.
6.1 C
- 7.2 C
normal boiling pt.
75.2 C
58.8 C
Our predicted values are close, but slightly higher, than the actual values. We do
correctly predict that bromine is a liquid at room temperature.
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