Hydrogen (atomic number 1), with the simplest of all atomic structures, has just one electron on principal energy level 1, so, in effect, its valence electron is also a core electron. The valence configuration for hydrogen is thus written as 1 s 1 . It should be noted, as described in the Electrons essay, that if a hydrogen atom (or any other atom) is in an excited state, it may reach energy levels beyond its normal, or ground, state.
Moving straight down the periodic table to francium (atomic number 87), which is in the same column as hydrogen, one finds that it has a valence electron configuration of 7 s 1 . Thus, although francium is vastly more complex and energy-filled than hydrogen, the two elements have the same valence shell configuration; only the number of the principal energy level is different. All the elements listed below hydrogen in Group 1 are therefore classified together as alkali metals. Obviously, hydrogen—a gas—is not part of the alkali metal family, nor does it clearly belong to any other family: it is the "lone wolf" of the periodic table.
Now look at two elements in Group 2, with beryllium (atomic number 4) and radium (88) at the top and bottom respectively. Beryllium has a valence shell configuration of 2 s 2 . This means its valence shell is at principal energy level 2, where there are two electrons on an s orbital pattern. Radium, though it is on period 7, nonetheless has the same valence shell configuration: 7 s 2 . This defines the alkaline earth metals family in terms of valence shell configuration.
For now, let us ignore groups 3 through 6—not to mention the columns between groups 2 and 3, unnumbered in the North American system—and skip over to Group 7. All the elements in this column, known as halogens, have valence shell configurations of n s 2 n p 5 . Beyond Group 7 is Group 8, the noble gases, all but one of whom have valence shell configurations of n s 2 n p 6 . The exception is helium, which has an s 2 valence shell. This seems to put it with the alkaline earth metals, but of course helium is not a metal. In terms of its actual behavior, it clearly belongs to the noble gases family.
The configurations of these valence shells have implications with regard to the ways in which elements bond, a subject developed at some length in the Chemical Bonding essay. Here we will consider it only in passing, to clarify the fact that electron configuration produces observable results. This is most obvious with the noble gases, which tend to resist bonding with most other elements because they already have eight electrons in their valence shell—the same number of valence electrons that most other atoms achieve only after they have bonded.
Groups 3 through 6, along with hydrogen and the four families so far identified, constitute the 44 representative or main-group elements. In 43 of these 44, the number of valence shell electrons is the same as the group number in the North American system. (Helium, which is in Group 8 but has two valence electrons, is the lone exception.) By contrast, the 40 elements listed in the "dip" at the middle of the chart—the transition metals—follow a less easily defined pattern. This is part of the reason why the North American system does not list them by group number, and also why neither system lists the two other families within the transition elements—the lanthanides and actinides.
Before addressing the transition metals, however, let us consider patterns of orbital filling, which also differentiate the representative elements from the transition elements. Each successive representative element fills all the orbitals of the elements that precede it (with some exceptions that will be explained), then goes on to add one more possible electron configuration. The total number of electrons—not just valence shell electrons—is the same as the atomic number. Thus fluorine, with an atomic number of 9, has a complete configuration of 1 s 2 2 s 2 2 p 5 . Neon, directly following it with an atomic number of 10, has a total configuration of 1 s 2 2 s 2 2 p 6 . (Again, this is not the same as the valence shell configuration, which is contained in the last two sub-levels represented: for example, 2 s 2 2 p 6 for neon.)
The chart that follows shows the pattern by which orbitals are filled. Note that in several places, the pattern of filling becomes "out of order," something that will be explained below.
Orbital Filling by Principal Energy Level
Generally, the 44 representative elements follow a regular pattern of orbital filling, and this is particularly so for the first 18 elements. Imagine a small amphitheater, shaped like a cone, with smaller rows of seats at the front. These rows are also designated by section, with the section number being the same as the number of rows in that section.
The two seats in the front row comprise a section labeled 1 or 1 s, and this is completely filled after helium (atomic number 2) enters the auditorium. Now the elements start filling section 2, which contains two rows. The first row of section 2, labeled 2 s, also has two seats, and after beryllium (4), it too is filled. Row 2 p has 6 seats, and it is finally filled with the entrance of neon (10). Now, all of section 2 has been filled; therefore, the eleventh element, sodium, starts filling section 3 at the first of its three rows. This row is 3 s —which, like all s rows, has only two seats. Thus, when element 13, aluminum, enters the theatre, it takes a seat in row 3 p, and eventually argon (18), completes that six-seat row.
By the pattern so far established, element 19 (potassium) should begin filling row 3 d by taking the first of its 10 seats. Instead, it moves on to section 4, which has four rows, and it takes the first seat in the first of those rows, 4 s. Calcium (20) follows it, filling the 4 s row. But when the next element, scandium (21), comes into the theatre, it goes to row 3 d, where potassium "should have" gone, if it had continued filling sections in order. Scandium is followed by nine companions (the first row of transition elements) before another representative element, gallium (31), comes into the theatre. (For reasons that will not be discussed here, chromium and copper, elements 24 and 29, respectively, have valence electrons in 4 s —which puts them slightly off the transition metal pattern.)
According to the "proper" order of filling seats, now that 3 d (and hence all of section 3) is filled, gallium should take a seat in 4 s. But those seats have already been taken by the two preceding representative elements, so gallium takes the first of six seats in 4 p. After that row fills up at krypton (36), it is again "proper" for the next representative element, rubidium (37), to take a seat in 4 d. Instead, just as potassium skipped 3 d, rubidium skips 4 d and opens up section 5 by taking the first of two seats in 5 s.
Just as before, the next transition element—yttrium (39)—begins filling up section 4 d, and is followed by nine more transition elements until cadmium (48) fills up that section. Then, the representative elements resume with indium (49), which, like gallium, skips ahead to section 5 p. And so it goes through the remainder of the periodic table, which ends with two representative elements followed by the last 10 transition metals.
Given the fact that it is actually the representative elements that skip the d sublevels, and the transition metals that go back and fill them, one might wonder if the names "representative" and "transition" (implying an interruption) should be reversed. However, remember the correlation between the number of valence shell electrons and group number for the representative elements. Furthermore, the transition metals are the only elements that fill the d orbitals.
This brings us to the reason why the lanthanides and actinides are set apart even from the transition metals. In most versions of the periodic table, lanthanum (57) is followed by hafnium (72) in the transition metals section of the chart. Similarly, actinium (89) is followed by rutherfordium (104). The "missing" metals—lanthanides and actinides, respectively—are listed at the bottom of the chart. There are reasons for this, as well as for the names of these groups.
After the 6 s orbital fills with the representative element barium (56), lanthanum does what a transition metal does—it begins filling the 5 d orbital. But after lanthanum, something strange happens: cerium (58) quits filling 5 d, and moves to fill the 4 f orbital. The filling of that orbital continues throughout the entire lanthanide series, all the way to lutetium (71). Thus, lanthanides can be defined as those metals that fill the 4 f orbital; however, because lanthanum exhibits similar properties, it is usually included with the lanthanides. Sometimes the term "lanthanide series" is used to distinguish the other 14 lanthanides from lanthanum itself.
A similar pattern occurs for the actinides. The 7 s orbital fills with radium (88), after which actinium (89) begins filling the 6 d orbital. Next comes thorium, first of the actinides, which begins the filling of the 5 f orbital. This is completed with element 103, lawrencium. Actinides can thus be defined as those metals that fill the 5 f orbital; but again, because actinium exhibits similar properties, it is usually included with the actinides.
The reader will note that for the seven families so far identified, we have generally not discussed them in terms of properties that can more easily be discerned—such as color, phase of matter, bonding characteristics, and so on. Instead, they have been examined primarily from the standpoint of orbital filling, which provides a solid chemical foundation for identifying families. Macroscopic characteristics, as well as the ways that the various elements find application in daily life, are discussed within essays devoted to the various groups.
Note, also, that the families so far identified account for only 92 elements out of a total of 112 listed on the periodic table: hydrogen; six alkali metals; six alkaline earth metals; five halogens; six noble gases; 40 transition metals; 14 lanthanides; and 14 actinides. What about the other 20? Some discussions of element families assign these elements, all of which are in groups 3 through 6, to families of their own, which will be mentioned briefly. However, because these "families" are not recognized by all chemists, in this book the 20 elements of groups 3 through 6 are described generally as metals, nonmetals, and metalloids.
Metals are lustrous or shiny in appearance, and malleable, meaning that they can be molded into different shapes without breaking. They are excellent conductors of heat and electricity, and tend to form positive ions by losing electrons. On the periodic table, metals fill the left, center, and part of the right-hand side of the chart. Thus it should not come as a surprise that most elements (87, in fact) are metals. This list includes alkali metals, alkaline earth metals, transition metals, lanthanides, and actinides, as well as seven elements in groups 3 through 6—aluminum, gallium, indium, thallium, tin, lead, and bismuth.
Nonmetals have a dull appearance; are not malleable; are poor conductors of heat and electricity; and tend to gain electrons to form negative ions. They are thus the opposite of metals in most regards, as befits their name. Nonmetals, which occupy the upper right-hand side of the periodic table, include the noble gases, halogens, and seven elements in groups 3 through 5. These nonmetal "orphans" are boron, carbon, nitrogen, oxygen, phosphorus, sulfur, and selenium. To these seven orphans could be added an eighth, from Group 1: hydrogen. As with the metals, a separate essay—with a special focus on the "orphans"—is devoted to nonmetals.
Occupying a diagonal region between the metals and nonmetals are metalloids, elements which exhibit characteristics of both metals and nonmetals. They are all solids, but are not lustrous, and conduct heat and electricity moderately well. The six metalloids are silicon, germanium, arsenic, antimony, tellurium, and polonium. Astatine is sometimes identified as a seventh metalloid; however, in this book, it is treated as a member of the halogen family.
Some sources list "families" rather than collections of "orphan" metals, metalloids, and nonmetals, in groups 3 through 6. These designations are not used in this book; however, they should be mentioned briefly. Group 3 is sometimes called the boron family; Group 4, the carbon family; Group 5, the nitrogen family; and Group 6, the oxygen family. Sometimes Group 5 is designated as the pnictogens, and Group 6 as the chalcogens.
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