the periodic table and valence electrons

When we get to nitrogen ( Z = 7, with seven electrons), Hund’s rule tells us that the lowest-energy arrangement is Experimentally, it is found that the ground state of a neutral carbon atom does indeed contain two unpaired electrons. By Hund’s rule, the electron configuration of carbon, which is 1 s 22 s 22 p 2, is understood to correspond to the orbital diagram shown in c. Hund, 1896–1997), which today says that the lowest-energy electron configuration for an atom is the one that has the maximum number of electrons with parallel spins in degenerate orbitals. Choice c illustrates Hund’s rule (named after the German physicist Friedrich H. Similarly, experiments have shown that choice b is slightly higher in energy (less stable) than choice c because electrons in degenerate orbitals prefer to line up with their spins parallel thus, we can eliminate choice b. Figure 6.29 tells us that the next lowest energy orbital is 2 s, so the orbital diagram for lithium isīecause of electron-electron repulsions, it is more favorable energetically for an electron to be in an unoccupied orbital than in one that is already occupied hence we can eliminate choice a. We know that the 1 s orbital can hold two of the electrons with their spins paired. The next element is lithium, with Z = 3 and three electrons in the neutral atom. Otherwise, our configuration would violate the Pauli principle. Written as 1 s 2, where the superscript 2 implies the pairing of spins. The orbital diagram for the helium atom is therefore From the Pauli exclusion principle, we know that an orbital can contain two electrons with opposite spin, so we place the second electron in the same orbital as the first but pointing down, so that the electrons are paired. We place one electron in the orbital that is lowest in energy, the 1 s orbital. For hydrogen, therefore, the single electron is placed in the 1 s orbital, and the electron configuration (also known a spdf notation) is written as 1 s 1 and read as “one-s-one.”Ī neutral helium atom, with an atomic number of 2 ( Z = 2), has two electrons. Here is a schematic orbital box diagram for a hydrogen atom in its ground state:įrom the orbital diagram, we can write the electron configuration in an abbreviated form in which the occupied orbitals are identified by their principal quantum number n and their value of l ( s, p, d, or f), with the number of electrons in the subshell indicated by a superscript. A filled orbital is indicated by ↑↓, in which the electron spins are said to be paired. That is, recognizing that each orbital can hold two electrons, one with spin up ↑, corresponding to m s = +½, which is arbitrarily written first, and one with spin down ↓, corresponding to m s = −½. First we determine the number of electrons in the atom then we add electrons one at a time to the lowest-energy orbital available without violating the Pauli Exclusion Principle. We write electronic configurations by following the aufbau principle (from German, meaning “building up”). The periodic table is an incredibly helpful tool in writing electron configurations. Using the periodic table to determine the electron configurations of atoms is key, but also keep in mind that there are certain rules to follow when assigning electrons to different orbitals. Recall, we can use the periodic table to rank the energy levels of various orbitals. The valence electrons, electrons in the outermost shell, are the determining factor for the unique chemistry of the element. Many of the physical and chemical properties of elements can be correlated to their unique electron configurations. Commonly, the electron configuration is used to describe the orbitals of an atom in its ground state, but it can also be used to represent an atom that has ionized into a cation or anion by compensating for the loss of or gain of electrons in their subsequent orbitals (we will examine those in the next section). The electron configuration of an atom is the representation of the arrangement of electrons distributed among the orbital shells and subshells.

The periodic table and valence electrons