In the process, a great amount of light and heat is released. The resulting salt is mostly unreactive — it is stable. It will not undergo any explosive reactions, unlike the sodium and chlorine that it is made of. Referring to the octet rule, atoms attempt to get a noble gas electron configuration, which is eight valence electrons. Sodium has one valence electron, so giving it up would result in the same electron configuration as neon.
Chlorine has seven valence electrons, so if it takes one it will have eight an octet. Chlorine has the electron configuration of argon when it gains an electron.
The octet rule could have been satisfied if chlorine gave up all seven of its valence electrons and sodium took them. In that case, both would have the electron configurations of noble gasses, with a full valence shell. Hund's Rule states that when electrons occupy degenerate orbitals i. Furthermore, the most stable configuration results when the spins are parallel i.
Nitrogen, for example, has 3 electrons occupying the 2p orbital. According to Hund's Rule, they must first occupy each of the three degenerate p orbitals, namely the 2p x orbital, 2p y orbital, and the 2p z orbital, and with parallel spins Figure 2. The configuration below is incorrect because the third electron occupies does not occupy the empty 2p z orbital.
Instead, it occupies the half-filled 2p x orbital. This, therefore, is a violation of Hund's Rule Figure 2. Figure 2. A visual representation of the Aufbau Principle and Hund's Rule.
Note that the filling of electrons in each orbital p x , p y and p z is arbitrary as long as the electrons are singly filled before having two electrons occupy the same orbital. Wolfgang Pauli postulated that each electron can be described with a unique set of four quantum numbers. Therefore, if two electrons occupy the same orbital, such as the 3s orbital, their spins must be paired.
The way we designate electronic configurations for cations and anions is essentially similar to that for neutral atoms in their ground state. The electronic configuration of cations is assigned by removing electrons first in the outermost p orbital, followed by the s orbital and finally the d orbitals if any more electrons need to be removed. In this case, all the 4p subshells are empty; hence, we start by removing from the s orbital, which is the 4s orbital.
Hence, we can say that both are isoelectronic. The electronic configuration of anions is assigned by adding electrons according to Aufbau Principle. We add electrons to fill the outermost orbital that is occupied, and then add more electrons to the next higher orbital.
Therefore, its ground state electronic configuration can be written as 1s 2 2s 2 2p 6 3s 2 3p 5. The chloride ion Cl - , on the other hand, has an additional electron for a total of 18 electrons.
Following Aufbau Principle, the electron occupies the partially filled 3p subshell first, making the 3p orbital completely filled. The electronic configuration for Cl - can, therefore, be designated as 1s 2 2s 2 2p 6 3s 2 3p 6. Hence, they are all isoelectronic to each other. Which of the princples explained above tells us that electrons that are paired cannot have the same spin value? What is a possible combination for the quantum numbers of the 5d orbital?
Give an example of an element which has the 5d orbital as it's most outer orbital. Introduction The electron configuration is the standard notation used to describe the electronic structure of an atom. Notation To help describe the appropriate notation for electron configuration, it is best to do so through example.
There are two ways in which electron configuration can be written: I: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 5 or I: [Kr]5s 2 4d 10 5p 5 In both of these types of notations, the order of the energy levels must be written by increased energy, showing the number of electrons in each subshell as an exponent. How can atoms achieve a stable electron configurations?
Apr 20, Related questions Are polyatomic ions molecular compounds or ionic compounds? What type of atoms tend to form covalent bonds? What is an example of a covalent bonds practice problem? What are some examples of stable electronic configurations?
Sodium chloride is an ionic compound, and the crystalline solid has the structure shown on the right. Transfer of the lone 3s electron of a sodium atom to the half-filled 3p orbital of a chlorine atom generates a sodium cation neon valence shell and a chloride anion argon valence shell.
Electrostatic attraction results in these oppositely charged ions packing together in a lattice. The attractive forces holding the ions in place can be referred to as ionic bonds. By clicking on the NaCl diagram , a model of this crystal will be displayed and may be manipulated. The other three reactions shown above give products that are very different from sodium chloride. Water is a liquid at room temperature; carbon dioxide and carbon tetrafluoride are gases. None of these compounds is composed of ions.
A different attractive interaction between atoms, called covalent bonding, is involved here. Covalent bonding occurs by a sharing of valence electrons, rather than an outright electron transfer. Examples of covalent bonding shown below include hydrogen, fluorine, carbon dioxide and carbon tetrafluoride. These illustrations use a simple Bohr notation, with valence electrons designated by colored dots. Note that in the first case both hydrogen atoms achieve a helium-like pair of 1s-electrons by sharing.
In the other examples carbon, oxygen and fluorine achieve neon-like valence octets by a similar sharing of electron pairs. Carbon dioxide is notable because it is a case in which two pairs of electrons four in all are shared by the same two atoms. This is an example of a double covalent bond.
Non-bonding valence electrons are shown as dots. These formulas are derived from the graphic notations suggested by A. Couper and A. Some examples of such structural formulas are given in the following table. Multiple bonding , the sharing of two or more electron pairs, is illustrated by ethylene and formaldehyde each has a double bond , and acetylene and hydrogen cyanide each with a triple bond.
Boron compounds such as BH 3 and BF 3 are exceptional in that conventional covalent bonding does not expand the valence shell occupancy of boron to an octet. Consequently, these compounds have an affinity for electrons, and they exhibit exceptional reactivity when compared with the compounds shown above.
The number of valence shell electrons an atom must gain or lose to achieve a valence octet is called valence. In covalent compounds the number of bonds which are characteristically formed by a given atom is equal to that atom's valence.
From the formulas written above, we arrive at the following general valence assignments:. The valences noted here represent the most common form these elements assume in organic compounds.
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