All atoms except hydrogen are multiple-electron atoms. The physical and chemical properties of elements are directly related to the number of electrons a neutral atom has. The periodic table of the elements groups elements with similar properties into columns. This systematic organization is related to the number of electrons in a neutral atom, called the atomic number, ZZ size 12{n} {}. We shall see in this section that the exclusion principle is key to the underlying explanations, and that it applies far beyond the realm of atomic physics.

In 1925, the Austrian physicist Wolfgang Pauli (see Figure 1) proposed the following rule: No two electrons can have the same set of quantum numbers. That is, no two electrons can be in the same state. This statement is known as the Pauli exclusion principle, because it excludes electrons from being in the same state. The Pauli exclusion principle is extremely powerful and very broadly applicable. It applies to any identical particles with half-integral intrinsic spin—that is, having s=1/2, 3/2, ...s=1/2, 3/2, ... size 12{s=1/2,`3/2, "." "." "." "." } {} Thus no two electrons can have the same set of quantum numbers.

Figure 1: The Austrian physicist Wolfgang Pauli (1900–1958) played a major role in the development of quantum mechanics. He proposed the exclusion principle; hypothesized the existence of an important particle, called the neutrino, before it was directly observed; made fundamental contributions to several areas of theoretical physics; and influenced many students who went on to do important work of their own. (credit: Nobel Foundation, via Wikimedia Commons)

Let us examine how the exclusion principle applies to electrons in atoms. The quantum numbers involved were defined in Quantum Numbers and Rules as n, l,ml, sn, l,ml, s, and msms size 12{m rSub { size 8{s} } } {}. Since ss size 12{s} {} is always 1/21/2 size 12{1/2} {} for electrons, it is redundant to list ss, and so we omit it and specify the state of an electron by a set of four numbers n,l,ml,msn,l,ml,ms. For example, the quantum numbers 2, 1, 0,−1/22, 1, 0,−1/2 size 12{ left (2,` 1,` 0,` - 1/2 right )} {} completely specify the state of an electron in an atom.

Since no two electrons can have the same set of quantum numbers, there are limits to how many of them can be in the same energy state. Note that nn size 12{n} {} determines the energy state in the absence of a magnetic field. So we first choose nn size 12{n} {}, and then we see how many electrons can be in this energy state or energy level. Consider the n=1n=1 size 12{n=1} {} level, for example. The only value ll size 12{l} {} can have is 0 (see (Reference) for a list of possible values once nn size 12{n} {} is known), and thus mlml can only be 0. The spin projection msms can be either +1/2+1/2 or −1/2−1/2, and so there can be two electrons in the n=1n=1 state. One has quantum numbers 1, 0, 0,+1/21, 0, 0,+1/2, and the other has 1, 0, 0,−1/21, 0, 0,−1/2. Figure 2 illustrates that there can be one or two electrons having n=1n=1 size 12{n=1} {}, but not three.

Figure 2: The Pauli exclusion principle explains why some configurations of electrons are allowed while others are not. Since electrons cannot have the same set of quantum numbers, a maximum of two can be in the n=1n=1 size 12{n=1} {} level, and a third electron must reside in the higher-energy n=2n=2 size 12{n=2} {} level. If there are two electrons in the n=1n=1 size 12{n=1} {} level, their spins must be in opposite directions. (More precisely, their spin projections must differ.)

Because of the Pauli exclusion principle, only hydrogen and helium can have all of their electrons in the n=1n=1 size 12{n=1} {} state. Lithium (see the periodic table) has three electrons, and so one must be in the n=2n=2 size 12{n=2} {} level. This leads to the concept of shells and shell filling. As we progress up in the number of electrons, we go from hydrogen to helium, lithium, beryllium, boron, and so on, and we see that there are limits to the number of electrons for each value of nn size 12{n} {}. Higher values of the shell nn size 12{n} {} correspond to higher energies, and they can allow more electrons because of the various combinations of l,mll,ml size 12{l, m rSub { size 8{l} } } {}, and msms size 12{m rSub { size 8{s} } } {} that are possible. Each value of the principal quantum number nn size 12{n} {} thus corresponds to an atomic shell into which a limited number of electrons can go. Shells and the number of electrons in them determine the physical and chemical properties of atoms, since it is the outermost electrons that interact most with anything outside the atom.

The probability clouds of electrons with the lowest value of ll size 12{l} {} are closest to the nucleus and, thus, more tightly bound. Thus when shells fill, they start with l=0l=0 size 12{l=0} {}, progress to l=1l=1 size 12{l=1} {}, and so on. Each value of ll size 12{l} {} thus corresponds to a subshell.

The table given below lists symbols traditionally used to denote shells and subshells.

To denote shells and subshells, we write nlnl size 12{ ital "nl"} {} with a number for nn size 12{n} {} and a letter for ll size 12{l} {}. For example, an electron in the n=1n=1 size 12{n=1} {} state must have l=0l=0 size 12{l=1} {}, and it is denoted as a 1s1s size 12{1s} {} electron. Two electrons in the n=1n=1 size 12{n=1} {} state is denoted as 1s21s2 size 12{1s rSup { size 8{2} } } {}. Another example is an electron in the n=2n=2 size 12{n=2} {} state with l=1l=1 size 12{l=1} {}, written as 2p2p size 12{2p} {}. The case of three electrons with these quantum numbers is written 2p32p3 size 12{2p rSup { size 8{3} } } {}. This notation, called spectroscopic notation, is generalized as shown in Figure 3.

Figure 3

Counting the number of possible combinations of quantum numbers allowed by the exclusion principle, we can determine how many electrons it takes to fill each subshell and shell.

Example 1: How Many Electrons Can Be in This Shell?

List all the possible sets of quantum numbers for the n=2n=2 size 12{n=2} {} shell, and determine the number of electrons that can be in the shell and each of its subshells.

Strategy

Given n=2n=2 size 12{n=2} {} for the shell, the rules for quantum numbers limit ll size 12{l} {} to be 0 or 1. The shell therefore has two subshells, labeled 2s2s size 12{2s} {} and 2p2p size 12{2p} {}. Since the lowest ll size 12{l} {} subshell fills first, we start with the 2s2s size 12{2s} {} subshell possibilities and then proceed with the 2p2p size 12{2p} {} subshell.

Solution

It is convenient to list the possible quantum numbers in a table, as shown below.

Figure 4

Discussion

It is laborious to make a table like this every time we want to know how many electrons can be in a shell or subshell. There exist general rules that are easy to apply, as we shall now see.

The number of electrons that can be in a subshell depends entirely on the value of ll size 12{l} {}. Once ll size 12{l} {} is known, there are a fixed number of values of mlml size 12{m rSub { size 8{l} } } {}, each of which can have two values for msms size 12{m rSub { size 8{s} } } {} First, since mlml size 12{m rSub { size 8{l} } } {} goes from −l−l size 12{ - l} {} to l in steps of 1, there are 2l+12l+1 size 12{2l+1} {} possibilities. This number is multiplied by 2, since each electron can be spin up or spin down. Thus the maximum number of electrons that can be in a subshell is 22l+122l+1 size 12{2 left (2l+1 right )} {}.

For example, the 2s2s size 12{2s} {} subshell in Example 1 has a maximum of 2 electrons in it, since 22l+1=20+1=222l+1=20+1=2 size 12{2 left (2l+1 right )=2 left (0+1 right )=2} {} for this subshell. Similarly, the 2p2p size 12{2p} {} subshell has a maximum of 6 electrons, since 22l+1=22+1=622l+1=22+1=6 size 12{2 left (2l+1 right )=2 left (2+1 right )=6} {}. For a shell, the maximum number is the sum of what can fit in the subshells. Some algebra shows that the maximum number of electrons that can be in a shell is 2n22n2 size 12{2n rSup { size 8{2} } } {}.

For example, for the first shell n=1n=1 size 12{n=1} {}, and so 2n2=22n2=2 size 12{2n rSup { size 8{2} } =2} {}. We have already seen that only two electrons can be in the n=1n=1 size 12{n=1} {} shell. Similarly, for the second shell, n=2n=2 size 12{n=2} {}, and so 2n2=82n2=8 size 12{2n rSup { size 8{2} } =8} {}. As found in Example 1, the total number of electrons in the n=2n=2 size 12{n=2} {} shell is 8.

Table 2 shows electron configurations for the first 20 elements in the periodic table, starting with hydrogen and its single electron and ending with calcium. The Pauli exclusion principle determines the maximum number of electrons allowed in each shell and subshell. But the order in which the shells and subshells are filled is complicated because of the large numbers of interactions between electrons.

Examining the above table, you can see that as the number of electrons in an atom increases from 1 in hydrogen to 2 in helium and so on, the lowest-energy shell gets filled first—that is, the n=1n=1 size 12{n=1} {} shell fills first, and then the n=2n=2 size 12{n=2} {} shell begins to fill. Within a shell, the subshells fill starting with the lowest ll size 12{l} {}, or with the ss size 12{s} {} subshell, then the pp size 12{p} {}, and so on, usually until all subshells are filled. The first exception to this occurs for potassium, where the 4s4s size 12{4s} {} subshell begins to fill before any electrons go into the 3d3d size 12{3d} {} subshell. The next exception is not shown in Table 2; it occurs for rubidium, where the 5s5s size 12{5s} {} subshell starts to fill before the 4d4d size 12{4d} {} subshell. The reason for these exceptions is that l=0l=0 size 12{l=0} {} electrons have probability clouds that penetrate closer to the nucleus and, thus, are more tightly bound (lower in energy).

Figure 5 shows the periodic table of the elements, through element 118. Of special interest are elements in the main groups, namely, those in the columns numbered 1, 2, 13, 14, 15, 16, 17, and 18.

Figure 5: Periodic table of the elements (credit: National Institute of Standards and Technology, U.S. Department of Commerce)

The number of electrons in the outermost subshell determines the atom’s chemical properties, since it is these electrons that are farthest from the nucleus and thus interact most with other atoms. If the outermost subshell can accept or give up an electron easily, then the atom will be highly reactive chemically. Each group in the periodic table is characterized by its outermost electron configuration. Perhaps the most familiar is Group 18 (Group VIII), the noble gases (helium, neon, argon, etc.). These gases are all characterized by a filled outer subshell that is particularly stable. This means that they have large ionization energies and do not readily give up an electron. Furthermore, if they were to accept an extra electron, it would be in a significantly higher level and thus loosely bound. Chemical reactions often involve sharing electrons. Noble gases can be forced into unstable chemical compounds only under high pressure and temperature.

Group 17 (Group VII) contains the halogens, such as fluorine, chlorine, iodine and bromine, each of which has one less electron than a neighboring noble gas. Each halogen has 5 pp size 12{p} {} electrons (a p5p5 size 12{p} {} configuration), while the pp size 12{p} {} subshell can hold 6 electrons. This means the halogens have one vacancy in their outermost subshell. They thus readily accept an extra electron (it becomes tightly bound, closing the shell as in noble gases) and are highly reactive chemically. The halogens are also likely to form singly negative ions, such as C1−C1− size 12{C1 rSup { size 8{ - {}} } } {}, fitting an extra electron into the vacancy in the outer subshell. In contrast, alkali metals, such as sodium and potassium, all have a single ss size 12{s} {} electron in their outermost subshell (an s 1 s 1 size 12{s rSup { size 8{1} } } {} configuration) and are members of Group 1 (Group I). These elements easily give up their extra electron and are thus highly reactive chemically. As you might expect, they also tend to form singly positive ions, such as Na+Na+ size 12{"Na" rSup { size 8{+{}} } } {}, by losing their loosely bound outermost electron. They are metals (conductors), because the loosely bound outer electron can move freely.

Of course, other groups are also of interest. Carbon, silicon, and germanium, for example, have similar chemistries and are in Group 4 (Group IV). Carbon, in particular, is extraordinary in its ability to form many types of bonds and to be part of long chains, such as inorganic molecules. The large group of what are called transitional elements is characterized by the filling of the dd size 12{d} {} subshells and crossing of energy levels. Heavier groups, such as the lanthanide series, are more complex—their shells do not fill in simple order. But the groups recognized by chemists such as Mendeleev have an explanation in the substructure of atoms.