Part C. Electron configurations are a shorthand form of an orbital diagram, describing which orbitals are occupied for a given element. For example, 1s22s22p1 is the electron configuration of boron.
Enter the complete electron configuration for arsenic (As).
Draw orbital diagrams, and use them to derive electron configurations
To understand how to draw orbital diagrams, and how they are used to write electron configurations.
The electron configuration of an element is the arrangment of its electrons in their atomic orbitals. Electron configurations can be used to predict most of the chemical properties of an element
Orbital diagrams are a useful tool to aid in the derivation of the electron configuration of an element. Orbital diagrams are filled using the aufbau principle, the Pauli principle, and Hund's rule.
Aufbau is German for "building up." The aufbau principle simply states that electrons are added to an orbital diagram one at a time to the lowest energy orbital available, and that the orbital diagram is thus "built up." However, due to shielding of the nucleus, the energies of orbitals are not always in order of energy level (n). For example, the 4s orbital is lower in energy than the 3d orbital for elements with more than one electron. To aid in remembering the energy order of orbitals, draw a diagram with the energy levels (1 through 8) down the left of the diagram, and the subshells of each energy level across in rows, with each row offset by one (so 3s is below 2p, 4p is below 3d, etc). To determine the order in which orbitals fill, read the diagram from top to bottom, left to right. This results in the order 1s2s2p3s3p4s3d4p5s, etc. This order is often called the "aufbau order."
The Pauli principle states that no two electrons in an atom can have the same value of all four quantum numbers (n, l, ml , and ms). The first three quantum numbers (n, l, and ml) specify a particular orbital, such as 1s. The fourth quantum number (ms) specifies the spin of the electron. Since there are only two possibly values for ms (+ 1/2 and − 1/2 ), only two electrons can occupy any given orbital. Remember that p subshells consist of three separate orbitals (px , py , and pz), for a total of up to six electrons in a given p subshell. Similarly, d subshells consist of five separate orbitals, and f subshells consists of seven separate orbitals.
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