Leave the text unchanged, do not delete or add anything, just correct typos, punctuation, and paragraphing! Use [English] and use Markdown for formatting. An electron configuration indicates how many electrons are occupying each subshell for an atom. Electron configurations are typically written with electrons in their ground state - that is, with electrons occupying the lowest available energy levels. The electron configuration for hydrogen is 1s1. This indicates one electron is in the 1s subshell. An orbital diagram represents individual orbitals as boxes and electrons as fishhooks. Orbitals are labeled with their subshell. Orbital diagrams contain more information than electron configurations because they show how electrons occupy degenerate orbitals in the same subshell. Helium has two electrons. The Pauli Exclusion Principle says that no two electrons in an atom can have the same four quantum numbers. As indicated in the electron configuration, we will add the second electron to the 1s subshell since each orbital can contain two electrons. However, as indicated in the orbital diagram, we must give the two electrons opposite spins. These two electrons have the quantum numbers 1, 0, 0, +½ and 1, 0, 0, -½. For hydrogen, all orbitals with the same principle quantum number are degenerate. So in hydrogen, the 4s, 4p, 4d, and 4f orbitals all have the same energy. But in a multi-electron atom, as we fill more and more orbitals, we develop sub-level splitting. This means that for a given shell, for example the fourth shell, the 4s orbital is the lowest energy, then the 4p orbitals, then the 4d orbitals, and the 4f orbitals are the highest energy. This sub-level splitting can be explained by considering three factors: Coulomb's law, the shielding effect, and the penetration effect. According to Coulomb's Law, opposite charges attract and like charges repel. So our potential energy of an electron increases as we move the electron farther from the nucleus. The shielding effect reduces the net positive charge that an electron experiences from the nucleus. Imagine we have a lithium atom. Lithium has two electrons in the 1s subshell and one electron in the 2s subshell. The electron in the 2s subshell is more likely to be found farther from the nucleus than the electrons in the 1s subshell. Since like charges repel, this electron which is farthest from the nucleus experiences the repulsion of the two other electrons which are closer to the nucleus. The overall effect is that the electron experiences a net charge from the nucleus of only about +1 instead of +3. The penetration effect reduces the shielding effect due to the overlap of the orbitals. Although the electron in the 2s subshell is more likely to be found farther away from the nucleus than the electrons in the 1s subshell, the probability overlaps so there are times when the electron in the 2s subshell might be just as close to the nucleus. During this time, the electron is not shielded as effectively and it experiences a higher net charge from the nucleus. We can see from this graph that the two electrons in the 1s subshell are most likely to be found within about 100 picometers of the nucleus. We can also see that an electron in the 2s subshell, which is represented by the red line, has some probability of being found within 100 picometers of the nucleus. This overlap of the 2s subshell with the 1s subshell causes the penetration effect of the 2s subshell. Then the probability for the 2s subshell drops off as it reaches a node and then increases again beyond the node. Notice that the 2p subshell does not have a node or the initial overlap of probability like the 2s subshell does. We can continue this pattern into the third shell. The 3s subshell has two nodes which put a relatively greater amount of the probability closer to the nucleus so it has the greatest penetration effect and subsequently the lowest energy compared to the 3p and the 3d subshells. The 3p subshell has just one node which allows it to also have some probability overlapping with the previous shell, so its energy is the next lowest because of the penetration effect. Whereas the 3d subshell does not have the same overlap so it has the highest energy. The overall effect is sublevel splitting. This becomes important because it dictates the order in which orbitals fill in a multi-electron atom. We'll also see later in this session that the general order of filling can also be determined from the periodic table. After helium comes lithium, which has three electrons. According to the Aufbau principle, we can build up electron configurations for multi-electron atoms by putting their electrons in the lowest-energy available orbitals. Our lowest energy orbital is the 1s orbital, which can hold two electrons. The next lowest energy orbital is the 2s orbital, so our third electron goes there. The electron configuration for lithium becomes 1s2 2s1. We can continue applying the Aufbau principle to get the electron configuration for carbon, which is 1s2 2s2 2p2. To create the orbital diagram for carbon, we need to know how to fill degenerate orbitals, which are orbitals with the same energy. According to Hund's rule, we fill degenerate orbitals singly first with parallel spins before doubling up two electrons in the same orbital. So our orbital diagram for carbon will have those two electrons in two different p orbitals with parallel spins. Since electrons repel each other, it makes sense that we should put a single electron in an empty orbital of the same energy before doubling up electrons in the same orbital. For the spin, it doesn't matter if we give both electrons a +½ spin or a -½ spin; they just have to be the same. This is because a moving charge produces a magnetic field, and the magnetic fields of each of these electrons interact to align their spins. We can see the Aufbau principle at work in our electron configurations for lithium through carbon. And we see Hund's rule at work as we finish out the orbital diagrams for nitrogen through neon. The 2p orbitals fill singly first with parallel spins before we double up two electrons in an orbital. We can also write abbreviated electron configurations using the noble gas that precedes the element. For example, by the time we get to sodium's electron configuration, it's becoming long. It is 1s2 2s2 2p6 3s1. The noble gas that precedes sodium is neon, whose electron configuration is 1s2 2s2 2p6. So we can also represent sodium as [Ne] 3s1, where the neon is the noble gas that precedes sodium and it's in brackets. Down a column of the periodic table, our main group elements end with the same electron configuration. For example, the elements in our column 1A all end with one electron in the highest s subshell, and the elements in our column 7A all end with five electrons in the highest p subshell. In fact, our periodic table is arranged in such a way that there are four distinct blocks. In our first two columns with helium included, we're filling our s subshells. In the last six columns, we're filling our p subshells. In our transition metals, we're filling our d subshells. And in our inner transition metals, we're filling our f subshells. We'll use this pattern to create electron configurations from the periodic table. The group number of the main group element indicates how many electrons are in the outermost shell. The row number of the main group element indicates what the value of n should be for the outermost shell. In the transition metals and inner transition metals, there are many exceptions to the normal pattern in the way electrons fill orbitals because those orbitals are very close in energy. Although there are more exceptions than these, four notable exceptions are for chromium, molybdenum, copper, and silver. We would expect the electron configurations of copper and silver to have two electrons in the s subshell and nine electrons in the d subshell. However, in both copper and silver, one electron from the s subshell is brought in to the d subshell so it is fully filled. So the electron configuration of copper ends in 4s1 3d10 and silver ends in 5s1 4d10. The same phenomenon happens with chromium and molybdenum. They also move one electron from the s subshell to the d subshell, which results in a half-filled d subshell. So the electron configuration for chromium ends in 4s1 3d5 and the electron configuration for molybdenum ends in 5s1 4d5.