The Aufbau section discussed how electrons fill the lowest energy orbitals first, and then move up to higher energy orbitals only after the lower energy orbitals are full. However, there is a problem with this rule. Certainly, 1s orbitals should be filled before 2s orbitals, because the 1s orbitals have a lower value of \(n\), and thus a lower energy. What about filling the three different 2p orbitals? In what order should they be filled? The answer to this question involves Hund's rule. Hund's rule states that: When assigning electrons to orbitals, an electron first seeks to fill all the orbitals with similar energy (also referred to as degenerate orbitals) before pairing with another electron in a half-filled orbital. Atoms at ground states tend to have as many unpaired electrons as possible. In visualizing this process, consider how electrons exhibit the same behavior as the same poles on a magnet would if they came into contact; as the negatively charged electrons fill orbitals, they first try to get as far as possible from each other before having to pair up.
Example \(\PageIndex{1}\): Nitrogen Atoms Consider the correct electron configuration of the nitrogen (Z = 7) atom: 1s2 2s2 2p3 The p orbitals are half-filled; there are three electrons and three p orbitals. This is because the three electrons in the 2p subshell will fill all the empty orbitals first before pairing with electrons in them. Keep in mind that elemental nitrogen is found in nature typically as molecular nitrogen, \(\ce{N2}\), which requires molecular orbitals instead of atomic orbitals as demonstrated above.
Example \(\PageIndex{2}\): Oxygen Atoms Next, consider oxygen (Z = 8) atom, the element after nitrogen in the same period; its electron configuration is: 1s2 2s2 2p4  Oxygen has one more electron than nitrogen; as the orbitals are all half-filled, the new electron must pair up. Keep in mind that elemental oxygen is found in nature typically as molecular oxygen, \(\ce{O_2}\), which has molecular orbitals instead of atomic orbitals as demonstrated above.
According to the first rule, electrons always enter an empty orbital before they pair up. Electrons are negatively charged and, as a result, they repel each other. Electrons tend to minimize repulsion by occupying their own orbitals, rather than sharing an orbital with another electron. Furthermore, quantum-mechanical calculations have shown that the electrons in singly occupied orbitals are less effectively screened or shielded from the nucleus. Electron shielding is further discussed in the next section. For the second rule, unpaired electrons in singly occupied orbitals have the same spins. Technically speaking, the first electron in a sublevel could be either "spin-up" or "spin-down." Once the spin of the first electron in a sublevel is chosen, however, the spins of all of the other electrons in that sublevel depend on that first spin. To avoid confusion, scientists typically draw the first electron, and any other unpaired electron, in an orbital as "spin-up."
Example \(\PageIndex{3}\): Carbon and Oxygen Consider the electron configuration for carbon atoms: 1s22s22p2: The two 2s electrons will occupy the same orbital, whereas the two 2p electrons will be in different orbital (and aligned the same direction) in accordance with Hund's rule. Consider also the electron configuration of oxygen. Oxygen has 8 electrons. The electron configuration can be written as 1s22s22p4. To draw the orbital diagram, begin with the following observations: the first two electrons will pair up in the 1s orbital; the next two electrons will pair up in the 2s orbital. That leaves 4 electrons, which must be placed in the 2p orbitals. According to Hund’s rule, all orbitals will be singly occupied before any is doubly occupied. Therefore, two p orbital get one electron and one will have two electrons. Hund's rule also stipulates that all of the unpaired electrons must have the same spin. In keeping with convention, the unpaired electrons are drawn as "spin-up", which gives (Figure 1).
When atoms come into contact with one another, it is the outermost electrons of these atoms, or valence shell, that will interact first. An atom is least stable (and therefore most reactive) when its valence shell is not full. The valence electrons are largely responsible for an element's chemical behavior. Elements that have the same number of valence electrons often have similar chemical properties. Electron configurations can also predict stability. An atom is most stable (and therefore unreactive) when all its orbitals are full. The most stable configurations are the ones that have full energy levels. These configurations occur in the noble gases. The noble gases are very stable elements that do not react easily with any other elements. Electron configurations can assist in making predictions about the ways in which certain elements will react, and the chemical compounds or molecules that different elements will form. Hund's Rules is shared under a CC BY-SA license and was authored, remixed, and/or curated by LibreTexts.
Learning Outcomes
The structure of the atom was discussed in the previous unit and now we will focus on the role that electrons play in the formation of compounds. Regardless of the type of compound or the number of atoms or electrons involved, it is the electrons of those atoms that interact to form a compound.
Electrons are not randomly arranged in an atom and their position within the atom can be described using electron arrangements, which are a simplified version of electron configurations. For each element of interest, we look at the number of electrons in a single atom and then determine how those electrons are arranged based on the atomic model. The main idea behind electron arrangements is that electrons can only exist at certain energy levels. By understanding the energy levels of electrons in an atom, we can predict properties and understand behavior of the atom. As shown in the figure below, there are multiple energy levels where electrons can be found. As the energy level increases, the energy difference between them decreases. A maximum of two electrons can be found in the \(n=1\) level; eight electrons can be in the \(n=2\) level. Although the \(n=3\) and \(n=4\) levels show only eight electrons in this diagram, those energy levels can hold more but not until we start looking at the transition metals. We will only be concerned with the electron arrangements of elements through calcium \(\left( Z = 20 \right)\) so we will put a maximum of eight electrons in the \(n=3\) level and two in the \(n=4\) level.
Example \(\PageIndex{1}\) What is the electron arrangement of oxygen? Solution Oxygen has eight electrons. The first two electrons will go in the \(n=1\) level. Two is the maximum number of electrons for the level so the other electrons will have to go in a higher energy level. The \(n=2\) level can hold up to eight electrons so the remaining six electrons will go in the \(n=2\) level. The electron arrangement of oxygen is (2, 6).
Example \(\PageIndex{2}\) What is the electron arrangement of chlorine? Solution Chlorine has 17 electrons. The first two electrons will go in the \(n=1\) level. Two is the maximum number of electrons for the level so the other electrons will have to go in higher energy levels. The \(n=2\) level can hold up to eight electrons so the next 8 electrons will go in the \(n=2\) level. The remaining 7 electrons can go in the \(n=3\) level since it holds a maximum of 8 electrons. The electron arrangement of chlorine is (2, 8, 7). The electron arrangement also provides information about the number of valence electrons. The valence electrons are the electrons in the highest energy level and the ones involved in ion and bond formation. Knowing the number of valence electrons will allow us to predict how a particular element will interact with other elements. Electrons in lower energy levels are called the core electrons. Let's look at the figure below which shows the electron diagram for magnesium and its 12 electrons. The first two electrons are found in the \(n=1\) energy level, the next eight electrons are found in the \(n=2\) level, and the remaining two electrons are found in the \(n=3\) level. The electrons always fill the lowest energy levels available until that level is filled, then electrons fill the next energy level until it is filled. This continues for all of the electrons in an atom. We can show the electron arrangement as (2, 8, 2) representing the electrons in the \(n=1\), \(n=2\), and \(n=3\) levels, respectively. The electron arrangement also shows the number of valence electrons which is two for magnesium because there are two electrons in the \(n=3\) energy level which is the highest occupied energy level for magnesium. This corresponds to the \(2+\) charge formed when magnesium forms an ion. It is willing to lose 2 electrons so that it has the same electron arrangements as the nearest noble gas, which is neon (2, 8). Atoms will gain or lose electrons to look like the nearest noble gas because the noble gases are unreactive due to the stability of having eight electrons in the highest energy level. This desire of atoms to have eight electrons in their outermost shell is known as the octet rule.
Example \(\PageIndex{3}\) What is the electron arrangement of aluminum? How many valence electrons does it have? Solution Aluminum has 13 electrons so it will have the electron arrangement (2, 8, 3) which represents two electrons in the \(n=1\) energy level, eight electrons in the \(n=2\) level, and three electrons in the \(n=3\) level. Aluminum has three valence electrons (indicated by the three electrons in the \(n=3\) level).
Example \(\PageIndex{4}\) How many valence electrons does chlorine have? How many electrons will chlorine gain or lose to form an ion? Solution Chlorine has 7 electrons in its valence shell. To meet the octet rule, it must either gain one electron or lose seven electrons. Gaining one is easier than losing seven so it will gain one electron to have a total of eight electrons when it forms an ion (i.e. charged particle). |
What is the number of the principal energy level in which the valence electrons are found in oxygen?
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