Niels Bohr first proposed his concept for energy levels in 1913, which depicts electrons circling about the nucleus like planets orbiting a star at various distances. If fire or electricity promoted an electron to a higher level, then it fell back to its original orbit. Such a model helps to explain the existence of line emission spectra.
Niels Bohr and Erwin Schrödinger
Until 1920s, all electrons within a shell were thought to behave identically to one another. In other words, all electrons in the second energy level of an oxygen atom were assumed to be the same. This way of thinking about the atom is a feature of what is commonly called the Bohr model, for Danish physicist Niels Bohr.
But in 1926, Erwin Schrödinger took things a step further and proved that electrons, being so small and moving so rapidly, don’t behave merely like particles, orbiting a nucleus. It’s not enough to think of them like planets orbiting a star. Instead, Schrödinger realized that electrons also behave like waves, and their wave motions could not be envisioned as small points of mass moving in a classical circular orbit about the nucleus.
In a now famous set of equations, Schrödinger showed that there was much-much more to the structure of each energy level in the electron cloud.
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Schrödinger’s Concept for Energy Level
To describe what he learned, Schrödinger proposed that electrons in the same energy level, or shell, could occupy different regions of space, each with their own distinct shape. We call these regions of space ‘subshells’. Within each subshell were individual regions of space that were equal in energy but had varying shape and orientation, called ‘orbitals’. Each orbital can be occupied by one or two electrons, but never more than a pair.
The number of subshells in each main shell increases with each additional shell. The first main shell has only one subshell, but the second main shell has two subshells, the third has three subshells, the fourth has four subshells, and so on.
The emission spectrum lines for a given element can be associated with specific subshell transitions, so the subshells were ultimately labeled ‘s’ for ‘sharp’, ‘p’ for ‘principal’, ‘d’ for ‘diffuse’, and ‘f’ for ‘fundamental’.
The Energy Level of Each Subshell
The first subshell of each main shell contains one orbital, which can hold two electrons, and is labeled ‘s’. The second subshell, when needed, is ‘p’, and it contains three orbitals, with room for six electrons. The third subshell, when that’s needed, is ‘d’, which has five orbitals and room for ten electrons. And the fourth subshell, which is only needed late in the table, is ‘f’. This is the biggest subshell, with seven orbitals, which have a capacity for 14 electrons.
Notice the pattern: 2,6,10,14… These four numbers, for the four types of subshell, are writ large across the overall shape of the periodic table. As we progress through the table from one element to the next, we see elements filling their s-subshell on the left of the table, a p-subshell to the right, a d-subshell in the wide expanse in the middle of the table, and an f-subshell on the landing strip at the bottom.
Because we can neatly carve the table into rectangular segments that correspond to the filling of discrete subshells, we often refer to these regions as ‘blocks’: the s-block, the p-block, the d-block, and the f-block.
The Aufbau Order
Because it’s the energies of these shells and subshells that is going to determine the order in which they fill, let’s understand how they are plotted on an axis of energy.
Our first energy level, of course, has just one subshell, but our second has two, and they’re a varying energy; the 2s is lower in energy than the 2p. This progression continues with our third energy level having 3s, 3p, and 3d subshells, all of varying energy.
What’s important to notice right here is that the 3d subshell is actually just a bit higher in energy than the 4s. So, as we increase through our energy axis, just looking at main shells, of course, the fourth is lower than the fifth is lower than the sixth overall, but we’ve got that little bit of overlap that we have to deal with. And that little bit of overlap is going to change the order in which they fill. So, if you adhere strictly to the energy axis, simply climbing in energy as you go, you’ll notice a drop back from a four to a three, from a six to a five, etc.
Plotting these as a function of their energy in a straight line, gives what’s called the Aufbau order. And the Aufbau order is the predicted order of filling of these shells and subshells with electrons, with two electrons for each orbital.
Fortunately, we don’t have to memorize the Aufbau order because the periodic table is actually a mnemonic that will guide us through this energy level of filling.
Common Questions about the Concept of Energy Levels and the Aufbau Order
In 1913, Niels Bohr introduced the concept of energy levels. According to this concept, Bohr pointed out that electrons orbiting the nucleus are like planets orbiting a star at different distances.
In 1926, Erwin Schrödinger proved that electrons, being very small and moving very rapidly, don’t behave merely like particles, orbiting a nucleus. It’s not enough to think of them like planets orbiting a star. Instead, Schrödinger realized that electrons also behave like waves, and their wave motions could not be envisioned as small points of mass moving in a classical circular orbit about the nucleus.
Each subshell around the nucleus has a specific energy level. The first subshell is labeled ‘s’ and has only one orbital; the second subshell labeled ‘p’ has three orbitals; the third subshell is labeled ‘d’ and has five orbitals; and the fourth subshell, which is labeled ‘f,’ has seven orbitals. Each of the orbitals can hold two electrons.