Chemistry Third Edition Julia Burdge Lecture Power Points

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Chemistry Third Edition Julia Burdge Lecture Power. Points Chapter 7 Electron Configuration and the

Chemistry Third Edition Julia Burdge Lecture Power. Points Chapter 7 Electron Configuration and the Periodic Table Copyright © 2012, The Mc. Graw-Hill Compaies, Inc. Permission required for reproduction or display.

CHAPTER 7. 1 7. 2 7. 3 7. 4 7. 5 7. 6 7.

CHAPTER 7. 1 7. 2 7. 3 7. 4 7. 5 7. 6 7. 7 7 Electron Configuration and the Periodic Table Development of the Periodic Table The Modern Periodic Table Effective Nuclear Charge Periodic Trends in Properties of Elements Electron Configuration of Ions Ionic Radius Periodic Trends in Chemical Properties of the Main Group Elements 2

7. 1 Development of the Periodic Table Topics Development of the Periodic Table 3

7. 1 Development of the Periodic Table Topics Development of the Periodic Table 3

7. 1 Development of the Periodic Table In 1864 the English chemist John Newlands

7. 1 Development of the Periodic Table In 1864 the English chemist John Newlands noticed that when the elements were arranged in order of atomic mass, every eighth element had similar properties. Newlands referred to this peculiar relationship as the law of octaves. However, this “law” turned out to be inadequate for elements beyond calcium, and Newlands’s work was not accepted by the scientific community. 4

7. 1 Development of the Periodic Table In 1869 the Russian chemist Dmitri Mendeleev

7. 1 Development of the Periodic Table In 1869 the Russian chemist Dmitri Mendeleev and the German chemist Lothar Meyer independently proposed a much more extensive tabulation of the elements based on the regular, periodic recurrence of properties—a phenomenon known as periodicity. 5

7. 1 Development of the Periodic Table Mendeleev ordered the elements based on atomic

7. 1 Development of the Periodic Table Mendeleev ordered the elements based on atomic mass, which led to inconsistencies. For example, argon and potassium would switch places on the periodic table if it were ordered by atomic mass. In such a case, potassium (very reactive) would be grouped with neon and helium (unreactive) and argon (unreactive) would be grouped with sodium and lithium (very reactive). This seriously upsets the idea of periodicity! 6

7. 1 Development of the Periodic Table The key to periodicity turns out to

7. 1 Development of the Periodic Table The key to periodicity turns out to be atomic number instead of atomic mass. Modern periodic table orders the elements by atomic number. Elements in a column of the modern periodic table have very similar chemical properties. 7

SAMPLE PROBLEM 7. 1 What elements would you expect to exhibit properties most similar

SAMPLE PROBLEM 7. 1 What elements would you expect to exhibit properties most similar to those of chlorine? Strategy Because elements in the same group tend to have similar properties, you should identify elements in the same group as chlorine. Setup Chlorine is a member of Group 7 A. 8

SAMPLE PROBLEM 7. 1 Solution Fluorine, bromine, and iodine, the other nonmetals in Group

SAMPLE PROBLEM 7. 1 Solution Fluorine, bromine, and iodine, the other nonmetals in Group 7 A, should have properties most similar to those of chlorine. 9

7. 2 The Modern Periodic Table Topics Classification of Elements Representing Free Elements in

7. 2 The Modern Periodic Table Topics Classification of Elements Representing Free Elements in Chemical Equations 10

7. 2 The Modern Periodic Table Classification of Elements Based on the type of

7. 2 The Modern Periodic Table Classification of Elements Based on the type of subshell containing the outermost electrons, the elements can be divided into categories—the main group elements, the noble gases, the transition elements (or transition metals), the lanthanides, and the actinides. The main group elements (also called the representative elements) are the elements in Groups 1 A through 7 A. 11

7. 2 The Modern Periodic Table 12

7. 2 The Modern Periodic Table 12

7. 2 The Modern Periodic Table Classification of Elements The outermost electrons of an

7. 2 The Modern Periodic Table Classification of Elements The outermost electrons of an atom are called valence electrons, which are the ones involved in the formation of chemical bonds between atoms. The similarity of the valence electron configurations (i. e. , they have the same number and type of valence electrons) is what makes the elements in the same group resemble one another chemically. For instance, the halogens (Group 7 A) all have outer electron configurations of ns 2 np 5, and they have similar properties. 13

7. 2 The Modern Periodic Table Classification of Elements 14

7. 2 The Modern Periodic Table Classification of Elements 14

7. 2 The Modern Periodic Table Representing Free Elements in Chemical Equations Metals Because

7. 2 The Modern Periodic Table Representing Free Elements in Chemical Equations Metals Because metals typically do not exist in discrete molecular units but rather in complex, three-dimensional networks of atoms, we always use their empirical formulas in chemical equations. The empirical formulas are the same as the symbols that represent the elements. For example, the empirical formula for iron is Fe, the same as the symbol for the element. 15

7. 2 The Modern Periodic Table Representing Free Elements in Chemical Equations Nonmetals For

7. 2 The Modern Periodic Table Representing Free Elements in Chemical Equations Nonmetals For nonmetals that exist as polyatomic molecules, we generally use the molecular formula in equations: H 2, N 2, O 2, F 2, Cl 2, Br 2, I 2, and P 4, for example. In the case of sulfur, however, we usually use the empirical formula S rather than the molecular formula S 8. 16

7. 2 The Modern Periodic Table Representing Free Elements in Chemical Equations Noble Gases

7. 2 The Modern Periodic Table Representing Free Elements in Chemical Equations Noble Gases All the noble gases exist as isolated atoms, so we use their symbols: He, Ne, Ar, Kr, Xe, and Rn. Metalloids The metalloids, like the metals, all have complex threedimensional networks, so we also represent them with their empirical formulas—that is, their symbols: B, Si, Ge, and so on. 17

7. 3 Effective Nuclear Charge Topics Effective Nuclear Charge 18

7. 3 Effective Nuclear Charge Topics Effective Nuclear Charge 18

7. 3 Effective Nuclear Charge Nuclear charge (Z) is simply the number of protons

7. 3 Effective Nuclear Charge Nuclear charge (Z) is simply the number of protons in the nucleus of an atom. An electron in a many-electron atom is partially shielded from the positive charge of the nucleus by the other electrons in the atom. Effective nuclear charge (Zeff) is the actual magnitude of positive charge that is “experienced” by a shielded electron in the atom. Because of shielding, the effective nuclear charge is less than the actual nuclear charge. 19

7. 3 Effective Nuclear Charge Although all the electrons in an atom shield one

7. 3 Effective Nuclear Charge Although all the electrons in an atom shield one another to some extent, those that are most effective at shielding are the core electrons. As a result, the value of Zeff increases steadily from left to right across a period of the periodic table because the number of core electrons remains the same (only the number of protons, Z, and the number of valence electrons increases). 20

7. 3 Effective Nuclear Charge In general, the effective nuclear charge is given by

7. 3 Effective Nuclear Charge In general, the effective nuclear charge is given by where is the shielding constant. The shielding constant is greater than zero but smaller than Z. 21

7. 3 Effective Nuclear Charge The change in Zeff as we move from the

7. 3 Effective Nuclear Charge The change in Zeff as we move from the top of a group to the bottom is generally less significant than the change as we move across a period. Although each step down a group represents a large increase in the nuclear charge, there is also an additional shell of core electrons to shield the valence electrons from the nucleus. Consequently, the effective nuclear charge changes less than the nuclear charge as we move down a column of the periodic table. 22

7. 4 Periodic Trends in Properties of Elements Topics Atomic Radius Ionization Energy Electron

7. 4 Periodic Trends in Properties of Elements Topics Atomic Radius Ionization Energy Electron Affinity Metallic Character Explaining Periodic Trends 23

7. 4 Periodic Trends in Properties of Elements Atomic Radius There are two ways

7. 4 Periodic Trends in Properties of Elements Atomic Radius There are two ways in which the atomic radius is commonly defined. One is the metallic radius, which is half the distance between the nuclei of two adjacent, identical metal atoms. 24

7. 4 Periodic Trends in Properties of Elements Atomic Radius The other is the

7. 4 Periodic Trends in Properties of Elements Atomic Radius The other is the covalent radius, which is half the distance between adjacent, identical nuclei in a molecule. 25

7. 4 Periodic Trends in Properties of Elements Atomic Radius The atomic radius decreases

7. 4 Periodic Trends in Properties of Elements Atomic Radius The atomic radius decreases as we move from left to right across a period and increases from top to bottom as we move down within a group. 26

7. 4 Periodic Trends in Properties of Elements 27

7. 4 Periodic Trends in Properties of Elements 27

7. 4 Periodic Trends in Properties of Elements Atomic Radius The increase down a

7. 4 Periodic Trends in Properties of Elements Atomic Radius The increase down a group is fairly easily explained. As we step down a column, the outermost occupied shell has an ever-increasing value of n, so it lies farther from the nucleus, making the radius bigger. 28

7. 4 Periodic Trends in Properties of Elements Atomic Radius As we move from

7. 4 Periodic Trends in Properties of Elements Atomic Radius As we move from left to right across a period, the effective nuclear charge increases and each step to the right adds another electron to the valence shell. So, there will be a more powerful attraction between the nucleus and the valence shell when the magnitudes of both charges increase. The result is that as we step across a period the valence shell is drawn closer to the nucleus, making the atomic radius smaller. 29

SAMPLE PROBLEM 7. 3 Referring only to a periodic table, arrange the elements P,

SAMPLE PROBLEM 7. 3 Referring only to a periodic table, arrange the elements P, S, and O in order of increasing atomic radius. Strategy Use the left-to-right (decreasing) and top-to-bottom (increasing) trends to compare the atomic radii of two of the three elements at a time. Setup Sulfur is to the right of phosphorus in the third row, so sulfur should be smaller than phosphorus. Oxygen is above sulfur in Group 6 A, so oxygen should be smaller than sulfur. 30

SAMPLE PROBLEM Solution 7. 3 O < S < P. 31

SAMPLE PROBLEM Solution 7. 3 O < S < P. 31

7. 4 Periodic Trends in Properties of Elements Ionization Energy Ionization energy (IE) is

7. 4 Periodic Trends in Properties of Elements Ionization Energy Ionization energy (IE) is the minimum energy required to remove an electron from an atom in the gas phase. Sodium, for example, has an ionization energy of 495. 8 k. J/mol, which is the energy input required to drive the process Specifically, this is the first ionization energy of sodium, IE 1(Na), which corresponds to the removal of the most loosely held electron. 32

7. 4 Periodic Trends in Properties of Elements Ionization Energy 33

7. 4 Periodic Trends in Properties of Elements Ionization Energy 33

7. 4 Periodic Trends in Properties of Elements Ionization Energy 34

7. 4 Periodic Trends in Properties of Elements Ionization Energy 34

7. 4 Periodic Trends in Properties of Elements Ionization Energy In general, as effective

7. 4 Periodic Trends in Properties of Elements Ionization Energy In general, as effective nuclear charge increases, ionization energy also increases. Thus, IE 1 increases from left to right across a period. Despite this trend, IE 1 for a Group 3 A element is smaller than that for the corresponding Group 2 A element. 35

7. 4 Periodic Trends in Properties of Elements Ionization Energy Likewise, IE 1 for

7. 4 Periodic Trends in Properties of Elements Ionization Energy Likewise, IE 1 for a Group 6 A element is smaller than that for the corresponding Group 5 A element. Both of these interruptions of the upward trend in IE 1 can be explained by using electron configuration. 36

7. 4 Periodic Trends in Properties of Elements Ionization Energy 37

7. 4 Periodic Trends in Properties of Elements Ionization Energy 37

7. 4 Periodic Trends in Properties of Elements Ionization Energy 38

7. 4 Periodic Trends in Properties of Elements Ionization Energy 38

SAMPLE PROBLEM 7. 4 Would you expect Na or Mg to have the greater

SAMPLE PROBLEM 7. 4 Would you expect Na or Mg to have the greater first ionization energy (IE 1)? Which should have the greater second ionization energy (IE 2)? Strategy Consider effective nuclear charge and electron configuration to compare the ionization energies. Effective nuclear charge increases from left to right in a period (thus increasing IE), and it is more difficult to remove a paired core electron than an unpaired valence electron. 39

SAMPLE PROBLEM 7. 4 Setup Na is in Group 1 A, and Mg is

SAMPLE PROBLEM 7. 4 Setup Na is in Group 1 A, and Mg is beside it in Group 2 A. Na has one valence electron, and Mg has two valence electrons. Solution IE 1(Mg) > IE 1(Na) because Mg is to the right of Na in the periodic table (i. e. , Mg has the greater effective nuclear charge, so it is more difficult to remove its electron). IE 2(Na) > IE 2(Mg) because the second ionization of Mg removes a valence electron, whereas the second ionization of Na removes a core electron. 40

7. 4 Periodic Trends in Properties of Elements Electron Affinity Electron affinity (EA) is

7. 4 Periodic Trends in Properties of Elements Electron Affinity Electron affinity (EA) is the energy released (the negative of the enthalpy change H) when an atom in the gas phase accepts an electron. 41

7. 4 Periodic Trends in Properties of Elements Electron Affinity 42

7. 4 Periodic Trends in Properties of Elements Electron Affinity 42

7. 4 Periodic Trends in Properties of Elements 43

7. 4 Periodic Trends in Properties of Elements 43

7. 4 Periodic Trends in Properties of Elements Electron Affinity Like ionization energy, electron

7. 4 Periodic Trends in Properties of Elements Electron Affinity Like ionization energy, electron affinity increases from left to right across a period. This trend in EA is due to the increase in effective nuclear charge from left to right (i. e. , it becomes progressively easier to add a negatively charged electron as the positive charge of the element’s nucleus increases). 44

7. 4 Periodic Trends in Properties of Elements Electron Affinity There also periodic interruptions

7. 4 Periodic Trends in Properties of Elements Electron Affinity There also periodic interruptions of the upward trend of EA from left to right, similar to those observed for IE 1, although they do not occur for the same elements. 45

7. 4 Periodic Trends in Properties of Elements Electron Affinity 46

7. 4 Periodic Trends in Properties of Elements Electron Affinity 46

SAMPLE PROBLEM 7. 5 For each pair of elements, indicate which one you would

SAMPLE PROBLEM 7. 5 For each pair of elements, indicate which one you would expect to have the greater first electron affinity, EA 1: (a) Al or Si, (b) Si or P. 47

SAMPLE PROBLEM 7. 5 Strategy Consider the effective nuclear charge and electron configuration to

SAMPLE PROBLEM 7. 5 Strategy Consider the effective nuclear charge and electron configuration to compare the electron affinities. The effective nuclear charge increases from left to right in a period (thus generally increasing EA), and it is more difficult to add an electron to a partially occupied orbital than to an empty one. Writing out orbital diagrams for the valence electrons is helpful for this type of problem. 48

SAMPLE PROBLEM 7. 5 Setup 49

SAMPLE PROBLEM 7. 5 Setup 49

SAMPLE PROBLEM 7. 5 Solution (a) EA 1(Si)> EA 1(Al) because Si is to

SAMPLE PROBLEM 7. 5 Solution (a) EA 1(Si)> EA 1(Al) because Si is to the right of Al and therefore has a greater effective nuclear charge. (b) EA 1(Si) > EA 1(P) because although P is to the right of Si in the third period of the periodic table (giving P the larger Zeff), adding an electron to a P atom requires placing it in a 3 p orbital that is partially occupied. The energy cost of pairing electrons outweighs the energy advantage of adding an electron to an atom with a larger effective nuclear charge. 50

7. 4 Periodic Trends in Properties of Elements Metallic Character Metals tend to •

7. 4 Periodic Trends in Properties of Elements Metallic Character Metals tend to • Be shiny, lustrous, and malleable • Be good conductors of both heat and electricity • Have low ionization energies (so they commonly form cations) • Form ionic compounds with chlorine (metal chlorides) • Form basic, ionic compounds with oxygen (metal oxides) 51

7. 4 Periodic Trends in Properties of Elements Metallic Character Nonmetals tend to •

7. 4 Periodic Trends in Properties of Elements Metallic Character Nonmetals tend to • Vary in color and lack the shiny appearance associated with metals • Be brittle, rather than malleable • Be poor conductors of both heat and electricity • Form acidic, molecular compounds with oxygen • Have high electron affinities (so they commonly form anions) 52

7. 4 Periodic Trends in Properties of Elements Metallic Character Metallic character increases from

7. 4 Periodic Trends in Properties of Elements Metallic Character Metallic character increases from top to bottom in a group and decreases from left to right within a period. Metalloids are elements with properties intermediate between those of metals and nonmetals. 53

7. 5 Electron Configuration of Ions Topics Ions of Main Group Elements Ions of

7. 5 Electron Configuration of Ions Topics Ions of Main Group Elements Ions of d-Block Elements 54

7. 5 Electron Configuration of Ions of Main Group Elements The 1 s 2

7. 5 Electron Configuration of Ions of Main Group Elements The 1 s 2 configuration of He and the ns 2 np 6 (n 2) valence electron configurations of the other noble gases are extraordinarily stable. Other main group elements tend to either lose or gain the number of electrons needed to achieve the same number of electrons as the nearest noble gas. Species with identical electron configurations are called isoelectronic. 55

7. 5 Electron Configuration of Ions of Main Group Elements To write the electron

7. 5 Electron Configuration of Ions of Main Group Elements To write the electron configuration of an ion formed by a main group element, we first write the configuration for the atom and either add or remove the appropriate number of electrons. Electron configurations for the sodium and chloride ions are 56

7. 5 Electron Configuration of Ions of Main Group Elements We can also write

7. 5 Electron Configuration of Ions of Main Group Elements We can also write electron configurations for ions using the noble gas core. 57

SAMPLE PROBLEM 7. 7 Write electron configurations for the following ions of main group

SAMPLE PROBLEM 7. 7 Write electron configurations for the following ions of main group elements: (a) N 3–, (b) Ba 2+, and (c) Be 2+. Strategy First write electron configurations for the atoms. Then add electrons (for anions) or remove electrons (for cations) to account for the charge. 58

SAMPLE PROBLEM 7. 7 Setup (a) N 3– forms when N ([He]2 s 22

SAMPLE PROBLEM 7. 7 Setup (a) N 3– forms when N ([He]2 s 22 p 3), a main group nonmetal, gains three electrons. (b) Ba 2+ forms when Ba ([Xe]6 s 2) loses two electrons. (c) Be 2+ forms when Be ([He]2 s 2) loses two electrons. Solution (a) [Ne] (b) [Xe] (c) [He] 59

7. 5 Electron Configuration of Ions of d-Block Elements We might expect the two

7. 5 Electron Configuration of Ions of d-Block Elements We might expect the two electrons lost in the formation of the Fe 2+ ion to come from the 3 d subshell. It turns out, though, that an atom always loses electrons first from the shell with the highest value of n. In the case of Fe, that would be the 4 s subshell. 60

7. 6 Ionic Radius Topics Comparing Ionic Radius with Atomic Radius Isoelectronic Series 61

7. 6 Ionic Radius Topics Comparing Ionic Radius with Atomic Radius Isoelectronic Series 61

7. 6 Ionic Radius Comparing Ionic Radius with Atomic Radius When an atom gains

7. 6 Ionic Radius Comparing Ionic Radius with Atomic Radius When an atom gains or loses one or more electrons to become an ion, its radius changes. The ionic radius, the radius of a cation or an anion, affects the physical and chemical properties of an ionic compound. The three-dimensional structure of an ionic compound, for example, depends on the relative sizes of its cations and anions. 62

7. 6 Ionic Radius Comparing Ionic Radius with Atomic Radius When an atom loses

7. 6 Ionic Radius Comparing Ionic Radius with Atomic Radius When an atom loses an electron and becomes a cation, its radius decreases due in part to a reduction in electron repulsions (and consequently a reduction in shielding) in the valence shell. A significant decrease in radius occurs when all of an atom’s valence electrons are removed. This is the case with ions of most main group elements, which are isoelectronic with the noble gases preceding them. 63

7. 6 Ionic Radius Comparing Ionic Radius with Atomic Radius 64

7. 6 Ionic Radius Comparing Ionic Radius with Atomic Radius 64

7. 6 Ionic Radius Isoelectronic Series An isoelectronic series is a series of two

7. 6 Ionic Radius Isoelectronic Series An isoelectronic series is a series of two or more species that have identical electron configurations, but different nuclear charges. For example, O 2–, F–, and Ne constitute an isoelectronic series. Although these three species have identical electron configurations, they have different radii. 65

7. 6 Ionic Radius Isoelectronic Series In an isoelectronic series, the species with the

7. 6 Ionic Radius Isoelectronic Series In an isoelectronic series, the species with the smallest nuclear charge (i. e. , the smallest atomic number Z) will have the largest radius. The species with the largest nuclear charge (i. e. , the largest Z) will have the smallest radius. 66

7. 6 Ionic Radius 67

7. 6 Ionic Radius 67

SAMPLE PROBLEM 7. 9 Identify the isoelectronic series in the following group of species,

SAMPLE PROBLEM 7. 9 Identify the isoelectronic series in the following group of species, and arrange the ions in order of increasing radius: K+, Ne, Ar, Kr, P 3–, S 2–, and Cl–. 68

SAMPLE PROBLEM 7. 9 Setup The number of electrons in each species is as

SAMPLE PROBLEM 7. 9 Setup The number of electrons in each species is as follows: 18 (K+), 10 (Ne), 18 (Ar), 36 (Kr), 18 (P 3–), 18 (S 2–), and 18 (Cl–). The species with 18 electrons constitute the isoelectronic series. The nuclear charges of the ions with 18 electrons are +19 (K+), +15 (P 3–), +16 (S 2–), and +17 (Cl–). 69

SAMPLE PROBLEM 7. 9 Solution The isoelectronic series includes K+, Ar, P 3–, S

SAMPLE PROBLEM 7. 9 Solution The isoelectronic series includes K+, Ar, P 3–, S 2–, and Cl–. The ions, in order of increasing radius, are: K+ < Cl– < S 2– < P 3– 70

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Topics General

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Topics General Trends in Chemical Properties of the Active Metals Properties of Other Main Group Elements Comparison of Group 1 A and Group 1 B Elements Variation in Properties of Oxides Within a Period 71

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements General Trends

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements General Trends in Chemical Properties We have said that elements in the same group resemble one another in chemical behavior because they have similar valence electron configurations. This statement, although correct in the general sense, must be applied with caution. Chemists have long known that the properties of the first member of each group (Li, Be, B, C, N, O, and F) are different from those of the rest of the members of the same group. 72

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements General Trends

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements General Trends in Chemical Properties Another trend in the chemical behavior of main group elements is the diagonal relationship. Diagonal relationships refer to similarities between pairs of elements in different groups and periods of the periodic table. 73

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements General Trends

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements General Trends in Chemical Properties The reason for this phenomenon is the similarity of charge densities of their cations. 74

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements General Trends

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements General Trends in Chemical Properties Hydrogen (1 s 1) There is no completely suitable position for hydrogen in the periodic table (it really belongs in a group by itself). 75

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Properties of

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Properties of the Active Metals Group 1 A Elements (ns 1, n ≥ 2) These elements all have low ionization energies, making it easy for them to become M+ ions. In fact, these metals are so reactive that they are never found in nature in the pure elemental state. They react with water to produce hydrogen gas and the corresponding metal hydroxide: 76

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements © The

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements © The Mc. Graw-Hill Companies, Inc. /Charles D. Winters, photographer The Mc. Graw-Hill Companies, Inc. /Stephen Frisch, photographer 77

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Properties of

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Properties of the Active Metals Group 1 A Elements (ns 1, n ≥ 2) Lithium forms lithium oxide (containing the oxide ion, O 2–): The other alkali metals all form oxides or peroxides containing the peroxide ion, O 22–): 78

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Properties of

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Properties of the Active Metals Group 1 A Elements (ns 1, n ≥ 2) Potassium, rubidium, and cesium also form superoxides (containing the superoxide ion, O 2– ): 79

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Properties of

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Properties of the Active Metals Group 2 A Elements (ns 2, n ≥ 2) As a group, the alkaline earth metals are somewhat less reactive than the alkali metals. Both the first and the second ionization energies decrease (and metallic character increases) from beryllium to barium. Group 2 A elements tend to form M 2+ ions, where M denotes an alkaline earth metal atom. 80

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements The Mc.

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements The Mc. Graw-Hill Companies, Inc. /Stephen Frisch, photographer © The Mc. Graw-Hill Companies, Inc. /Charles D. Winters, photographer © Ed R. Degginger/Color-Pic, Inc. The Mc. Graw-Hill Companies, Inc. /Stephen Frisch, photographer 81

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Properties of

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Properties of the Active Metals Group 2 A Elements (ns 2, n ≥ 2) The reactions of alkaline earth metals with water vary considerably. Beryllium does not react with water; magnesium reacts slowly with steam; and calcium, strontium, and barium react vigorously with cold water. 82

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Properties of

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Properties of the Active Metals Group 2 A Elements (ns 2, n ≥ 2) The reactivity of the alkaline earth metals toward oxygen also increases from Be to Ba. Beryllium and magnesium form oxides (Be. O and Mg. O) only at elevated temperatures, whereas Ca. O, Sr. O, and Ba. O form at room temperature. 83

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Properties of

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Properties of Other Main Group Elements Group 3 A Elements (ns 2 np 1, n ≥ 2) Boron, the first member of the group, is a metalloid; the others (Al, Ga, In, and Tl) are metals. Boron does not form binary ionic compounds and is unreactive toward both oxygen and water. Aluminum, the next element in the group, readily forms aluminum oxide when exposed to air: 84

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Properties of

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Properties of Other Main Group Elements Group 3 A Elements (ns 2 np 1, n ≥ 2) Aluminum forms the Al 3+ ion. It reacts with hydrochloric acid according to the equation: The other Group 3 A metals (Ga, In, and Tl) can form both M+ and M 3+ ions. As we move down the group, the M+ ion becomes the more stable of the two. 85

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements The Mc.

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements The Mc. Graw-Hill Companies, Inc. /Stephen Frisch, photographer © The Mc. Graw-Hill Companies, Inc. /Charles D. Winters, photographer The Mc. Graw-Hill Companies, Inc. /Stephen Frisch, photographer 86

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Properties of

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Properties of Other Main Group Elements Group 4 A Elements (ns 2 np 2, n ≥ 2) Carbon, the first member of the group, is a nonmetal, whereas silicon and germanium, the next two members, are metalloids. Tin and lead, the last two members of the group, are metals. They do not react with water, but they do react with aqueous acid to produce hydrogen gas: 87

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Properties of

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Properties of Other Main Group Elements Group 4 A Elements (ns 2 np 2, n ≥ 2) The Group 4 A elements form compounds in both the +2 and +4 oxidation states. For carbon and silicon, the +4 oxidation state is the more stable one. In tin compounds the +4 oxidation state is only slightly more stable than the +2 oxidation state. In lead compounds the +2 oxidation state is the more stable one. 88

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Properties of

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Properties of Other Main Group Elements Group 4 A Elements (ns 2 np 2, n ≥ 2) The outer electron configuration of lead is 6 s 26 p 2, and lead tends to lose only the 6 p electrons to form Pb 2+ rather than both the 6 p and 6 s electrons to form Pb 4+. 89

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements © The

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements © The Mc. Graw-Hill Companies, Inc. /Charles D. Winters, photographer 90

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Properties of

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Properties of Other Main Group Elements Group 5 A Elements (ns 2 np 3, n ≥ 2) Nitrogen and phosphorus are nonmetals, arsenic and antimony are metalloids, and bismuth is a metal. Because Group 5 A contains elements in all three categories, we expect greater variation in their chemical properties. 91

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Properties of

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Properties of Other Main Group Elements Group 5 A Elements (ns 2 np 3, n ≥ 2) Elemental nitrogen is a diatomic gas (N 2). It forms a variety of oxides (NO, N 2 O, NO 2, N 2 O 4, and N 2 O 5), all of which are gases except for N 2 O 5, which is a solid at room temperature. Nitrogen has a tendency to accept three electrons to form the nitride ion (N 3–). Most metal nitrides, such as Li 3 N and Mg 3 N 2, are ionic compounds. 92

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Properties of

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Properties of Other Main Group Elements Group 5 A Elements (ns 2 np 3, n ≥ 2) Phosphorus exists as individual P 4 molecules (white phosphorus) or chains of P 4 molecules (red phosphorus). It forms two solid oxides with the formulas P 4 O 6 and P 4 O 10. 93

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements 94

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements 94

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Properties of

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Properties of Other Main Group Elements Group 6 A Elements (ns 2 np 4, n ≥ 2) The first three members of the group (oxygen, sulfur, and selenium) are nonmetals, whereas the last two (tellurium and polonium) are metalloids. Oxygen is a colorless, odorless, diatomic gas; elemental sulfur and selenium exist as the molecules S 8 and Se 8, respectively; and tellurium and polonium have more extensive threedimensional structures. 95

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Properties of

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Properties of Other Main Group Elements Group 6 A Elements (ns 2 np 4, n ≥ 2) Oxygen has a tendency to accept two electrons to form the oxide ion (O 2–) in many compounds. Sulfur, selenium, and tellurium also form ions by accepting two electrons: S 2–, Se 2–, and Te 2–. The elements in Group 6 A (especially oxygen) form a large number of molecular compounds with nonmetals. 96

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements © David

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements © David Tietz/Editorial Image, LLC © The Mc. Graw-Hill Companies, Inc. /Charles D. Winters, photographer © Richard Treptow/Photo Researchers The Mc. Graw-Hill Companies, Inc. /Stephen Frisch, photographer 97

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Properties of

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Properties of Other Main Group Elements Group 7 A Elements (ns 2 np 5, n ≥ 2) All the halogens are nonmetals with the general formula X 2, where X denotes a halogen element. Like the Group 1 A metals, the Group 7 A nonmetals are too reactive to be found in nature in the elemental form. The halogens have high ionization energies and large, energetically favorable electron affinities. Anions derived from the halogens (F–, Cl–, Br–, and I–) are called halides. 98

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Properties of

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Properties of Other Main Group Elements Group 7 A Elements (ns 2 np 5, n ≥ 2) The vast majority of alkali metal halides are ionic compounds. The halogens also form many molecular compounds among themselves, such as ICl and Br. F 3, and with nonmetals in other groups, such as NF 3, PCl 5, and SF 6. The halogens react with hydrogen to form hydrogen halides: 99

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements © The

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements © The Mc. Graw-Hill Companies, Inc. /Stephen Frisch, photographer © Charles D. Winters/Photo Researchers © The Mc. Graw-Hill Companies, Inc. /Charles D. Winters, photographer 100

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Properties of

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Properties of Other Main Group Elements Group 8 A Elements (ns 2 np 6, n ≥ 2) All the noble gases exist as monatomic species. With the exception of helium, which has the electron configuration 1 s 2, their atoms have completely filled outer ns and np subshells. Their electron configurations give the noble gases their great stability. 101

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Properties of

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Properties of Other Main Group Elements Group 8 A Elements (ns 2 np 6, n ≥ 2) The Group 8 A ionization energies are among the highest of all the elements. Their electron affinities are all less than zero, so they have no tendency to accept extra electrons. 102

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Properties of

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Properties of Other Main Group Elements Group 8 A Elements (ns 2 np 6, n ≥ 2) Beginning in 1963, compounds were prepared from the heavier members of the group by exposing them to very strong oxidizing agents such as fluorine and oxygen. Some of the compounds that have been prepared are Xe. F 4, Xe. O 3, Xe. OF 4, Kr. F 2, and most recently, HAr. F. 103

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements The Mc.

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements The Mc. Graw-Hill Companies, Inc. /Stephen Frisch, photographer 104

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Comparison of

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Comparison of Group 1 A and Group 1 B Elements Although the outer electron configurations of Groups 1 A and 1 B are similar (members of both groups have a single valence electron in an s orbital), their chemical properties are very different. The first ionization energies of Cu, Ag, and Au are 745, 731, and 890 k. J/mol, respectively. Because these values are considerably larger than those of the alkali metals, the Group 1 B elements are much less reactive. 105

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Comparison of

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Comparison of Group 1 A and Group 1 B Elements The higher ionization energies of the Group 1 B elements result from incomplete shielding of the nucleus by the inner d electrons (compared with the more effective shielding by the completely filled noble gas cores). Consequently, the outer s electrons of the Group 1 B elements are more strongly attracted by the nucleus. In fact, copper, silver, and gold are so unreactive that they are usually found in the uncombined state in nature. 106

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Variation in

7. 7 Periodic Trends in Chemical Properties of the Main Group Elements Variation in Properties of Oxides Within a Period Most oxides can be classified as acidic or basic depending on whether they produce acidic or basic solutions when dissolved in water (or whether they react as acids or bases). Some oxides are amphoteric, which means that they display both acidic and basic properties. 107