Download Chapter 8 Periodic Relationships Among the Elements Section 8.1

Chapter 8
Periodic Relationships Among the Elements
This chapter presents a qualitative view of the periodic (repeating) relationships of the elements in the
periodic table. Upon completion of Chapter 8, the student should be able to:
1. Explain the basis of the periodic table as described by Mendeleev and Meyer and indicate the
shortcomings of their method.
2. Explain the basis of the periodic table as described by Moseley and how it predicted properties of
“missing” elements.
3. Identify elements that correspond to each of the following groups:
• representative elements
• noble gases
• transition metals
• lanthanides
• actinides
4. Describe the electron configuration of cations and anions and identity ions and atoms that are
isoelectronic.
5. Apply the concept of effective nuclear charge and shielding constants (screening constants) to justify
why the first ionization energy is always smaller than the second ionization energy of a given atom.
6. Predict the trends from left to right and top to bottom of the periodic table for each of the following:
• atomic radius
• ionic radius
• ionization energy
• electron affinity
• metallic character
7. Relate why hydrogen could be placed in a class by itself when reviewing its chemical properties.
8. Provide examples of Group 1A elements reacting with oxygen to form oxides, peroxides, and
superoxides.
9. Predict the reaction of alkali metals with water.
10. Describe the reactivity of alkaline earth metals with water.
11. Relate how strontium-90 could lead to human illness.
12. Compare the reactivity of boron, a metalloid, to aluminum.
13. Identify the metals, nonmetal, and metalloids of Group 4A.
14. Recall the reactions that form nitric acid, phosphoric acid and sulfuric acid.
15. List the halides (halogens)
16. Indicate the three hydrohalic acids that are strong acids and the one hydrohalic acid that is a weak
acid.
17. Explain why the name for Group 8A has changed from inert gases to noble gases.
18. List the three “coinage” metals and explain their relative inertness.
19. Rationalize the characteristics of the properties of oxides of the third period elements.
20. Classify oxides as acidic, basic, or amphoteric.
21. Explain why concentrated bases such as NaOH should not be stored in glass containers.
Section 8.1
Development of the Periodic Table
Your author points out that the original periodic table developed by Mendeleev and Meyer based on
atomic mass fell short because such an arrangement places potassium before argon. This would suggest that
potassium should be in the same group as the noble gases and that argon would be classified with the alkali
metals. This dilemma is resolved when atomic number is used to build the periodic table instead of atomic
mass. An interesting assignment might be to have your students examine the rest of the periodic table and
find where increasing atomic mass does not yield a smooth increase in atomic number. Once these
inconsistencies are found, then perhaps by investigating the properties of these elements involved, suggest
why atomic number places the elements correctly while atomic mass would not.
Section 8.2
Periodic Classification of the Elements
A combination of electron configurations (Chapter Seven) and the periodic table can be used to explain
a great deal of reaction chemistry. The similarity of electron configurations for the alkali metals is easy for
students to see. Once this has been pointed out, students find it fairly easy to accept that sodium and
potassium would have similar properties. Also, as the electron configurations change as one moves from left
to right across the periodic table, it is reasonable to expect that chemical properties will also change. The
concept of valence electrons, the outer electrons of an atom, is introduced in this section. We will use this
concept extensively later when we discuss electron dot notation and valence shell electron pair repulsion
theory (VSEPR).
Your author uses this section to introduce your students to the correct method of writing electron
configurations for anions and cations. Electron configurations for representative elements are straight forward
as long as your students understand the concept of isoelectronic (having the same number of electrons). All
ions of representative elements (representative elements only and not transition metals) that are isoelectronic
to an atom will have the same electron configuration as that atom. For example, Na+, F-, O2-, and N3- all are
isoelectronic to Ne (ten electrons), thus the electron configurations of these four ions will be the same as Ne.
This method will not hold for transition metal ion electron configuration however. Let us use iron as an
example. Recall that in Chapter Seven we used the concepts of paramagnetism and diamagnetism.
Experimentally, we find that Fe, Fe2+, and Fe 3+ are all paramagnetic with Fe and Fe2+ both having four
unpaired electrons while Fe3+ has five unpaired electrons. Using only the valence electrons for Fe, we have
the following box diagram (Electron configuration of Fe is [Ar]4s23d6).
↑↓
↑
↑↓
4s
↑
↑
↑
3d
To make Fe2+ from Fe, we must remove two electrons. Our first inclination would be to remove two electrons
from the 3d level to get
↑
↑
↑
3d
↑
4s
↑↓
which would be consistent with Fe2+ having four unpaired electrons; however, this method will not be
successful for Fe3+ (five unpaired electrons) because using the same logic as above would yield the following
for Fe3+
↑
↑
3d
↑
4s
↑↓
which has only three unpaired electrons and not the five that is required. The correct method is to remove the
electrons from the 4s orbital first which will result in the following for Fe2+
↑↓
↑
↑
↑
↑
3d
4s
giving four unpaired electrons as is experimentally determined. This method will result in the following for Fe3+
↑
↑
↑
↑
↑
3d
4s
which gives five unpaired electrons as is required.
Your author points out that the reason the s electrons are removed first, is because electron-electron
and electron-nucleus interactions are different for atoms than for ions. One way to rationalize removing the 4s
electrons first is to recall that 3d orbital electrons have the greatest probability of being found in orbitals that
loop inward toward the nucleus while 4s orbital electrons are more likely to be found further away from the
nucleus. Students may then conclude that 3s electrons should be removed before the 3p because of the
shape of the p orbitals; however, note that both 3s and 3p have the same principle quantum number while 4s
and 3d do not. An interesting question to ask your students is whether Cu+ is paramagnetic or diamagnetic.
This question incorporates one of the exceptions from Chapter Seven on electron configuration and also the
concept of removing 4s electron before 3d when making the copper ion.
Section 8.3
Periodic Variation in Physical Properties
It is easy for students to accept that atomic radii increase as one moves down a column in the periodic
table. It is logical for more electrons to require more space. It is, however, more difficult for them to accept
that atomic radii decrease as one moves from left to right across a period. One way to assist them in
accepting this is to remind them that the principle quantum number does not change as one stays in the same
period; therefore, there is little energy change in the electrons as one progresses across a given period (unlike
going down a group in a given column when n increases). The number of protons increases as one moves
across a period, the effective nuclear charge increases and pulls the electrons in closer, and thus the atomic
radii decrease.
If an ion has more electrons than protons (anion), it is reasonable that it will have a larger radius than
the corresponding atom where the number of electrons equals the number of protons. The same logic is used
to understand that an ion that has more protons than electrons (cation) will have a radius smaller than the
neutral atom. This is illustrated with a grouping of isoelectronic (same number of electrons) ions and atoms
such as Al3+, Mg2+, Na+, Ne, and F-. As listed, they are in order of increasing size. All five have ten electrons,
but the number of protons is varying, thus the size difference.
Section 8.4
Ionization Energy
Ionization energy, the energy required to remove an electron from a gas phase atom, follows once the
concept of atomic radius is understood. It makes sense that an electron that is further away from the nucleus
(larger atomic radius) will require less energy to remove. If it requires less energy, then by definition, it has a
lower ionization energy. Therefore, since the size of atoms increases as one goes down a column in the
periodic table, the ionization energy will decrease. The size of atoms decreases as one goes from left to right
across a period, thus the ionization must increase.
It is interesting to note that since it requires energy for a gaseous atom to ionize, it is an endothermic
reaction. This indicates that isolated gaseous atoms will not spontaneously ionize because that would violate
the law of conservation of energy.
Section 8.5
Electron Affinity
Electron affinity is defined as the negative of the energy change that occurs when an electron is
accepted by an atom in the gaseous phase. Stated in a different way, electron affinity is the amount of energy
required to remove an electron form of a gaseous ion. Therefore, if electron affinity is large (a large amount of
energy is required to remove an electron from the anion), then the anion is very stable.
Section 8.6
Variation in Chemical Properties of the Representative Elements
Your author does an extensive discussion of each group of representative elements in this section. It is
noted that properties are similar within a group, but they do vary when the group contains a metalloid. It may
be worth mentioning that the reaction of sulfur trioxide with water
SO3(g) + H2O(ℓ) H2SO4(aq)
is the fundamental reaction for what we know as acid rain. Also, when reviewing Group 7A, remind your
students that hydrofluoric acid is a weak acid (does not completely ionize), but is reactive enough to etch
glass. This is to once again reinforce the concept that weak acids are not necessarily less dangerous than
strong acids.
Finally, on a practical note, your author indicates that very concentrated bases like NaOH(aq) should not
be stored in Pyrex glassware because of the potential reaction of SiO2(s) and the base. Many of us have
probably experienced ground-glass stoppers stuck in volumetric flasks because of this reaction.