CHEMISTRY THE CENTRAL SCIENCE
7 PERIODIC PROPERTIES OF THE ELEMENTS
CHAPTER SUMMARY AND KEY TERMS
INTRODUCTION AND SECTION 7.1 The periodic table was first developed by Mendeleev and Meyer on the basis of the similarity in chemical and physical properties exhibited by certain elements. Mose-ley established that each element has a unique atomic number, which added more order to the periodic table. We now recognize that elements in the same column of the periodic table have the same number of electrons in their valence orbitals. This similarity in valence electronic structure leads to the similarities among elements in the same group. The differences among elements in the same group arise because their valence orbitals are in different shells.
SECTION 7.2 Many properties of atoms are due to the average distance of the outer electrons from the nucleus and to the effective nuclear charge experienced by these electrons. The core electrons are very effective in screening the outer electrons from the full charge of the nucleus, whereas electrons in the same shell do not screen each other effectively. As a result, the effective nuclear charge experienced by valence electrons increases as we move left to right across a period.
SECTION 7.3 The size of an atom can be gauged by its bonding atomic radius, based on measurements of the distances separating atoms in their chemical compounds. In general, atomic radii increase as we go down a column in the periodic table and decrease as we proceed left to right across a row.
Cations are smaller than their parent atoms; anions are larger than their parent atoms. For ions of the same charge, size increases going down a column of the periodic table. An isoelectronic series is a series of ions that has the same number of electrons. For such a series, size decreases with increasing nuclear charge as the electrons are attracted more strongly to the nucleus.
SECTION 7.4 The first ionization energy of an atom is the minimum energy needed to remove an electron from the atom in the gas phase, forming a cation. The second ionization energy is the energy needed to remove a second electron, and so forth. Ionization energies show a sharp increase after all the valence electrons have been removed because of the much higher effective nuclear charge experienced by the core electrons. The first ionization energies of the elements show periodic trends that are opposite those seen for atomic radii, with smaller atoms having higher first ionization energies. Thus, first ion-ization energies decrease as we go down a column and increase as we proceed left to right across a row.
We can write electron configurations for ions by first writing the electron configuration of the neutral atom and then removing or adding the appropriate number of electrons. Electrons are removed first from the orbitals with the largest value of n. If there are two valence orbitals with the same value of n (such as 4s and 4p), then the electrons are lost first from the orbital with a higher value of l (in this case, 4p). Electrons are added to orbitals in the reverse order.
SECTION 7.5 The electron affinity of an element is the energy change upon adding an electron to an atom in the gas phase, forming an anion. A negative electron affinity means that the anion is stable; a positive electron affinity means that the anion is not stable relative to the separated atom and electron, in which case its exact value cannot be measured. In general, electron affinities become more negative as we proceed from left to right across the periodic table. The halogens have the most-negative electron affinities. The electron affinities of the noble gases are positive because the added electron would have to occupy a new, higher-energy subshell.
SECTION 7.6 The elements can be categorized as metals, non-metals, and metalloids. Most elements are metals; they occupy the left side and the middle of the periodic table. Nonmetals appear in the upper-right section of the table. Metalloids occupy a narrow band between the metals and nonmetals. The tendency of an element to exhibit the properties of metals, called the metallic character, increases as we proceed down a column and decreases as we proceed from left to right across a row.
Metals have a characteristic luster, and they are good conductors of heat and electricity. When metals react with nonmetals, the metal atoms are oxidized to cations and ionic substances are generally formed. Most metal oxides are basic; they react with acids to form salts and water.
Nonmetals lack metallic luster and are generally poor conductors of heat and electricity. Several are gases at room temperature. Compounds composed entirely of nonmetals are generally molecular. Nonmetals usually form anions in their reactions with metals. Nonmetal oxides are acidic; they react with bases to form salts and water. Metalloids have properties that are intermediate between those of metals and nonmetals.
SECTION 7.7 The periodic properties of the elements can help us understand the properties of groups of the representative elements. The alkali metals (group 1A) are soft metals with low densities and low melting points. They have the lowest ionization energies of the elements. As a result, they are very reactive toward nonmetals, easily losing their outer s electron to form 1+ ions. The alkaline earth metals (group 2A) are harder and more dense and have higher melting points than the alkali metals. They are also very reactive toward nonmetals, although not as reactive as the alkali metals. The alkaline earth metals readily lose their two outer s electrons to form 2+ ions. Both alkali and alkaline earth metals react with hydrogen to form ionic substances that contain the hydride ion, H−.
SECTION 7.8 Hydrogen is a nonmetal with properties that are distinct from any of the groups of the periodic table. It forms molecular compounds with other nonmetals, such as oxygen and the halogens.
Oxygen and sulfur are the most important elements in group 6A. Oxygen is usually found as a diatomic molecule, O2. Ozone, O3, is an important allotrope of oxygen. Oxygen has a strong tendency to gain electrons from other elements, thus oxidizing them. In combination with metals, oxygen is usually found as the oxide ion, O2−, although salts of the peroxide ion, O22−, and superoxide ion, O2−, are sometimes formed. Elemental sulfur is most commonly found as S8 molecules. In combination with metals, it is most often found as the sulfide ion, S2−.
The halogens (group 7A) are nonmetals that exist as diatomic molecules. The halogens have the most negative electron affinities of the elements. Thus, their chemistry is dominated by a tendency to form 1- ions, especially in reactions with metals.
The noble gases (group 8A) are nonmetals that exist as monatomic gases. They are very unreactive because they have completely filled s and p subshells. Only the heaviest noble gases are known to form compounds, and they do so only with very active nonmetals, such as fluorine.
• Understand the meaning of effective nuclear charge, Zeff, and how Zeff depends on nuclear charge and electron configuration. (Section 7.2)
• Use the periodic table to predict the trends in atomic radii, ionic radii, ionization energy, and electron affinity. (Sections 7.2, 7.3, 7.4, and 7.5)
• Explain how the radius of an atom changes upon losing electrons to form a cation or gaining electrons to form an anion. (Section 7.3)
• Be able to write the electron configurations of ions. (Section 7.3)
• Explain how the ionization energy changes as we remove successive electrons. Recognize the jump in ionization energy that occurs when the ionization corresponds to removing a core electron. (Section 7.4)
• Understand how irregularities in the periodic trends for electron affinity can be related to electron configuration. (Section 7.5)
• Recognize the differences in chemical and physical properties of metals and nonmetals, including the basicity of metal oxides and the acidity of nonmetal oxides. (Section 7.6)
• Understand how the atomic properties, such as ionization energy and electron configuration, are related to the chemical reactivity and physical properties of the alkali and alkaline earth metals (groups 1A and 2A). (Section 7.7)
• Be able to write balanced equations for the reactions of the group 1A and 2A metals with water, oxygen, hydrogen, and the halogens. (Sections 7.7 and 7.8)
• Understand and recognize the unique characteristics of hydrogen. (Section 7.7)
• Understand how the atomic properties (such as ionization energy, electron configuration, and electron affinity) of group 6A, 7A, and 8A elements are related to their chemical reactivity and physical properties. (Section 7.8)
Estimating effective nuclear charge