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The Periodic Table

In a previous section, the periodic table was introduced as a list of the elements. We also pointed out that the design of the periodic table separates the metals from the nonmetals. In this section we will show how the various features of the table relate to the electron configuration of the different elements and to their position in the table. First let us point out those features using the complete periodic table shown in Figure 5.10. In the table, the elements are placed in rows and columns of varying length. Seven rows are used to show all of the elements now known. These rows are called periods and each period is numbered. Notice that the display of elements labeled "lanthanides" and placed below the table belongs in period 6 between element 57 (lanthanum) and element 72 (hafnium). In some periodic tables, lanthanum is the first member of the lanthanide series. Similarly, the display labeled "actinides" belongs in period 7 between element 89 (actinium) and element 104 (rutherfordium). Again, in some tables actinium is the first member of the actinide series. These two displays are customarily put below the table so that the table will fit into a reasonable space. The columns of the periodic table vary in length. Some are numbered. The short columns, those in the middle of the table, have not been numbered.

FIGURE 5.10 Periodic table of the elements.

The elements in a column make up a family of elements. A family is also known as a group. Thus the elements in column 8 are known as the family or group of noble gases.

A. Electron Configuration and the Periodic Table
Figure 5.11 again shows the periodic table but without the symbols of the elements. Instead it shows the last sublevels filled in describing the electron configurations of the elements in each section. We will use Figure 5.11 and Figure 5.8 to relate the electron configuration of an element to its position in the periodic table.


FIGURE 5.11 The periodic table and the energy level subshells.

FIGURE 5.8 The principal energy levels of an atom and the sublevels and orbitals each contains. The arrows show the order in which the sublevels fill.


In period 1, there are two boxes. In the usual table, these boxes would contain the symbols for hydrogen and helium, the elements in this period. In Figure 5.11 we show instead the letter s indicating that the last added electron for the elements in these boxes is in the 1s sublevel. In period 2, there are eight boxes. Instead of symbols for eight elements, Figure 5.11 shows s in the first two boxes and p in the last six boxes, showing that the 2s and 2p sublevels are being filled as the electron configurations of the elements in these boxes are completed. Period 3 also has eight boxes, which would correspond to the electrons needed to fill the 3s and 3p sublevels.

Look back now to Figure 5.8, which shows the order in which the sublevels fill. Notice that the 4s sublevel is filled immediately after the 3p sublevel. Figure 5.11 shows that elements whose last added electron goes into an s sublevel are in columns 1 and 2. So we must start here a new period, period 4, and put boxes for the elements formed by filling the 4s sublevel in those columns. Figure 5.8 shows that the next sublevel to fill is the 3d sublevel. These are the first d electrons added, so we start new columns for the elements formed by their addition. Ten electrons are needed to fill the five d orbitals, so we start ten columns in this fourth period, placing the columns next to column 2 and between it and column 3. The 4p sublevel is filled next, after the 3d sublevel. The boxes for the elements formed by filling the p orbitals are in place under the boxes for elements formed by adding the 3p electrons.

By consulting Figure 5.8, we see that the next sublevels filled are in the order: 5s, 4d, and 5p. Boxes for the elements formed by filling the orbitals of these sublevels are arranged as were those in period 4. Just as period 4 contains more elements than period 3, period 6 contains more elements than period 5. Period 6 starts with elements whose last added electron is in the 6s sublevel. The next step is where period 6 differs from period 5. Look again at Figure 5.8 and note that the 4f sublevel is filled after the 6s sublevel and before the 5d sublevel. We will need 14 boxes to contain the electrons needed to fill the seven f orbitals. These are the boxes of the lanthanide series, shown below the table. There is some evidence that these orbitals do not fill before one electron is in a 5d orbital, so we have shown in Figure 5.11 the lanthanide series coming after the first d column. After the 4f orbitals are filled, boxes are shown for the rest of the elements formed by adding 5d and 6p electrons. The seventh period contains boxes for the elements formed by filling the 7s, the 5f (the actinide series shown below the table), and finally the 6d sublevels.

Figure 5.11 thus shows the close relationship that exists between the electron configuration of an element and its location in the periodic table. This relationship is further expressed by the following names sometimes given to parts of the table:

columns 1 and 2 s block
columns 3-8 p block
short columns d block
lanthanides and actinides f block

The groups of elements found in these blocks are also known by other names.


B. Categories of Elements in the Periodic Table

1. The representative elements
Elements in the s and p blocks are known as representative elements or main group elements. The term representative dates from early times, when chemists believed that the chemistry of these elements was representative of all elements. Group 8 is not always included in the representative elements because the chemistry of the noble gases is unique to them. In period 7 there are no elements in the p block.

The p block of period 7 would contain elements with atomic numbers greater than 112; such elements have not yet been found in the Earth's crust nor have they been prepared by nuclear reaction.

In the s and p blocks, the period in which the element occurs has the same number as the highest energy level that contains electrons in a ground-state atom. The number of the column in which an element is found is the same as the number of s and p electrons in that level. Sodium is a representative element with 11 electrons. Its electron configuration is:


Sodium is in column 1 of the third period. In a sodium atom, the highest-energy principal energy level containing electrons is the third energy level, and that energy level contains one electron.

2. The transition elements
The transition elements (or transition metals, for they are all metals) are those elements found in the short columns of the d block. Many of these elements are probably familiar to you. The coinage metals--gold, silver, and copper--are here. So is iron, the principal ingredient of steel, as well as those elements that are added to iron to make particular kinds of steel: chromium, nickel, and manganese. In period 7, the d block is not filled. The reason is the same as the reason why the p section of period 7 is empty: these elements do not occur naturally and have not yet been found as the product of a nuclear reaction. Many of the properties of the transition elements are related to the fact that, in their electron structures, the occupied s and d sublevels of highest energy are very close in energy.

3. The inner transition elements
The inner transition elements are those found in the f block of the periodic table (in the two rows below the main body of the table). The elements in this block are chemically very much alike, which will seem reasonable when you consider that they have the same electron configurations in the two outermost energy levels. The differences occur in the next further-in energy level. For example, the electron configuration of cerium (Ce, #58) is:


and that of praseodymium (Pr, #59) is:


The only difference between these two configurations is in the number of 4f electrons. Both the fifth and sixth energy levels contain electrons.

The elements in the lanthanide series are also known as the rare earths. They are used extensively in producing monitors for color television.

The elements in the actinide series are all radioactive, and only three are found in appreciable concentration in the Earth's crust. Of the others, only some have been found in trace amounts in the Earth or in the stars. All have been produced in laboratories as products of nuclear reactions.

C. The Electron Configuration of the Noble Gases; Core Notation
We have established a relationship between the electron configuration of an element and its location in the periodic table. Let us look closer now at the electron configurations of the noble gases, those elements in Group 8 of the periodic table. The electron configurations of these elements are shown in Table 5.3.

TABLE 5.3 Electron configurations of the noble gases (Group 8 elements)
Element Atomic number Electron configuration
He 2 1s2
Ne 10 1s22s22p6
Ar 18 1s22s22p63s23p6
Kr 36 1s22s22p63s23p63d104s24p6
Xe 54 1s22s22p63s23p63d104s24p64d105s25p6
Rn 86 1s22s22p63s23p63d104s24p64d104f145s25p65d106s26p6

A careful examination of these configurations shows that none has any partially filled sublevels. The symbol of a noble gas enclosed in brackets is used to represent those filled sublevels. As an example, consider the electron configuration of bromine:

Br: 1s22s22p63s23p63d104s24p5

The first 18 electrons are in the same orbitals as those of an atom of argon (see Table 5.3). If we use the symbol [Ar] to represent those 18 electrons, we can write the electron configuration of bromine as

Br: [Ar]3d104s24p5

This device is useful because we can write electron configurations more quickly. More importantly, this notation emphasizes the electron configurations in the higher energy levels, where the differences are important in determining the chemistry of an element. This use of the noble gases to represent certain configurations is known as core notation. The symbol of a noble gas enclosed in brackets represents the inner, filled orbitals of an element. Additional electrons are shown outside the brackets in the standard way. Note that only the noble gases can be used in core notation. When using this method, remember that, even though the inner configuration of an element may be written the same as that of a noble gas, the energies of these inner electrons are slightly different.

Table 5.4 shows, in core notation, the electron configurations of the elements in Groups 1 and 6 of the periodic table. Notice how this method emphasizes the similar structure of the elements in a single column.

TABLE 5.4 Electron configurations of elements in Groups I and VI, using core notations
Group 1 Group 6
H 1S1
Li [He]2s1
Na [Ne]3s1
K [Ar]4s1
Rb [Kr]5s1
Cs [Xe]6s1
Fr [Rn]7s1
O [He]2s22p4
S [Ne]3s23p4
Se [Ar]4s23d104p4
Te [Kr]5s24d105p4
Po [Xe]6s24f145d106p4

D. Valence ElectronsTable
In discussing the chemical properties of an element, we often focus on electrons in the outermost occupied energy level. These outer-shell electrons are called valence electrons, and the energy level they occupy is called the valence shell. Valence electrons participate in chemical bonding and chemical reactions. The valence electrons of an element are shown by using a representation of the element called an electron-dot structure or Lewis structure<, named after G. N. Lewis, the twentieth-century American chemist who first pointed out the importance of outer-shell electrons. A Lewis structure shows the symbol of the element surrounded by a number of dots equal to the number of electrons in the outer energy level of the atoms of the element.

You may have noticed in writing electron configurations that the s sublevel of a principal energy level n is always occupied before d electrons are added to the principal energy level numbered n - 1. Immediately after filling the d sublevel of principal level n - 1, the p sublevel of principal level n is filled, and the next sublevel filled will be the s sublevel of the n + 1 principal energy level. This order of filling is illustrated in the configurations of krypton, xenon, and radon in Table 5.3 and of selenium, tellurium, and polonium in Table 5.4. The significance of these observations is that, in the electron configuration of any atom, the principal energy level with the highest number that contains any electrons cannot contain more than eight electrons. This also means that the valence electrons of an atom are the s and p electrons in the occupied principal energy level of highest number. Consequently, no atom can have more than eight valence electrons.

In drawing the Lewis structure of an atom, we imagine a four-sided box around the symbol of the atom and consider that each side of that box corresponds to an orbital. We represent each valence electron as a dot. The first two valence electrons will be s electrons; they would be represented by two dots on a side (it doesn't matter which side) of the symbol. The valence electrons that are in the p subshell are placed first, one on each of the remaining sides of the symbol, and then a second one is added to each side. This method of filling is similar to the one used in drawing box diagrams of electron configurations. As an example, consider the Lewis structure of sodium.

Looking back at Table 5.4, we see that the core notation for sodium is [Ne]3s1. This tells us that a sodium atom has one electron in its outer shell, so its Lewis structure is . The core notation for selenium is [Ar]3d104s24p4. Its Lewis structure is . The ten 3d electrons of selenium are not shown because they are not in the outer shell, which is the principal energy level 4. Lewis structures for the elements in the first three periods and Group 2 of the periodic table are shown in Table 5.5.


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