Now let’s discuss how to determine the electron configuration for an atom—in other words, how electrons are arranged in an atom. The first and most important rule to remember when attempting to determine how electrons will be arranged in the atom is Hund’s rule, which states that the most stable arrangement of electrons is that which allows the maximum number of unpaired electrons. This arrangement minimizes electron-electron repulsions. Here’s an analogy. In large families with several children, it is a luxury for each child to have his or her own room. There is far less fussing and fighting if siblings are not forced to share living quarters: the entire household experiences a lower-intensity, less-frazzled energy state. Likewise, electrons will go into available orbitals singly before beginning to pair up. All the single–occupant electrons of orbitals have parallel spins, are designated with an upward-pointing arrow, and have a magnetic spin quantum number of +1/2.
As we mentioned earlier, each principal energy level, n, has n sublevels. This means the first has one sublevel, the second has two, the third has three, etc. The sublevels are named s, p, d, and f.
| Energy level principal quantum number, n |
Number of sublevels |
Names of sublevels |
| 1 |
1 |
s |
| 2 |
2 |
s, p |
| 3 |
3 |
s, p, d |
| 4 |
4 |
s, p, d, f |
At each additional sublevel, the number of available orbitals is increased by two: s = 1, p = 3, d = 5, f = 7, and as we stated above, each orbital can hold only two electrons, which must be of opposite spin. So s holds 2, p holds 6 (2 electrons times the number of orbitals, which for the p sublevel is equal to 3), d holds 10, and f holds 14.
| Sublevel |
s |
p |
d |
f |
| Number of orbitals |
1 |
3 |
5 |
7 |
| Maximum number of electrons |
2 |
6 |
10 |
14 |
| Quantum number, l |
0 |
1 |
2 |
3 |
We can use the periodic table to make this task easier.
Notice there are only two elements in the first period (the first row of the periodic table); their electrons are in the first principal energy level: n = 1. The second period (row) contains a total of eight elements, which all have two sublevels: s and p; s sublevels contain two electrons when full, while p sublevels contain six electrons when full (because p sublevels each contain three orbitals).
The third period looks a lot like the second because of electron-electron interference. It takes less energy for an electron to be placed in 4s than in 3d, so 4s fills before 3d. Notice that the middle of the periodic table contains a square of 10 columns: these are the elements in which the d orbitals are being filled (these elements are called the transition metals). Now look at the two rows of 14 elements at the bottom of the table. In these rare earth elements, the f orbitals are being filled.
One final note about electron configurations. You can use the periodic table to quickly determine the valence electron configuration of each element. The valence electrons are the outermost electrons in an atom—the ones that are involved in bonding. The day of the test, as soon as you get your periodic table (which comes in the test booklet), label the rows as shown in the art above. The number at the top of each of the rows (i.e., 1A, 2A, etc.) will tell you how many valence electrons each element in that particular row has, which will be very helpful in determining Lewis dot structures. More on this later.
