A Nitrogen Atom Has Seven Electrons How Many Can Be Shared With Other Atoms
3.2: Covalent Bonding
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A second method past which atoms can achieve a filled valence shell is by sharing valence electrons with another cantlet. Thus fluorine, with one unpaired valence electron, can share that electron with an unshared electron on another fluorine to form the chemical compound, F ii in which the two shared electrons course a chemic bond holding the ii fluorine atoms together. When you do this, each fluorine now has the equivalent of eight electrons in its valence trounce; three unshared pairs and one pair that is shared between the 2 atoms. Notation that when you are counting electrons, the electrons that are shared in the covalent bond are counted for each atom, individually. A chemic bond formed by sharing electrons between atoms is called a covalent bond. When two or more atoms are bonded together utilizing covalent bonds, the compound is referred to as a molecule.
There is a simple method, given below, that nosotros can use to construct Lewis diagrams for diatomic and for polyatomic molecules:
- Begin by adding up all of the valence electrons in the molecule. For F 2 , each fluorine has 7, giving a total of 14 valence electrons.
- Adjacent, draw your cardinal atom. For a diatomic molecule like F two , both atoms are the aforementioned, but if several different atoms are present, the central atom will be to the left (or lower) in the periodic table.
- Next, draw the other atoms around the fundamental cantlet, placing 2 electrons between the atoms to class a covalent bond.
- Distribute the remaining valence electrons, as pairs, around each of the outer atoms, so that they all are surrounded by viii electrons.
- Place any remaining electrons on the central atom.
- If the cardinal atom is not surrounded past an octet of electrons, construct multiple bonds with the outer atoms until all atoms take a complete octet.
- If there are an odd number of valence electrons in the molecule, leave the remaining single electron on the key cantlet.
Permit's utilize these rules for the Lewis diagram for chlorine gas, Cl two . At that place are xiv valence electrons in the molecule. Both atoms are the same, so nosotros draw them adjacent to each other and place two electrons between them to form the covalent bond. Of the twelve remain electrons, we now identify six around one chlorine (to give an octet) and then identify the other half-dozen around the other chlorine (our central cantlet). Checking, we encounter that each cantlet is surrounded by an octet of valence electrons, so our construction is consummate.
\[:\underset{..}{\overset{..}{Cl}}\cdot \cdot \underset{..}{\overset{..}{Cl}}:\]
All of the Group 7A elements (the halogens), have valence shells with 7 electrons and all of the common halogens be in nature equally diatomic molecules; fluorine, F 2 ; chlorine, Cl ii ; bromine, Br ii and iodine I 2 (astatine, the halogen in the sixth menstruum, is a brusk-lived radioactive element and its chemical properties are poorly understood). Nitrogen and oxygen, Group 5A and 6A elements, respectively, also exists in nature as diatomic molecules (N 2 and O ii ). Let'due south consider oxygen; oxygen has six valence electrons (a Group 6A element). Following the logic that nosotros used for chlorine, nosotros draw the two atoms and place one pair of electrons betwixt them, leaving ten valence electrons. We place three pairs on one oxygen cantlet, and the remaining two pairs on the 2nd (our central atom). Because we only have 6 valence electrons surrounding the second oxygen cantlet, we must move i pair from the other oxygen and form a second covalent bond (a double bond) betwixt the two atoms. Doing this, each atom now has an octet of valence electrons.
\[:\underset{..}{O}::\overset{..}{O}:\]
Nitrogen has five valence electrons. Sharing 1 on each atom gives the showtime intermediate where each nitrogen is surrounded by six electrons (not enough!). Sharing another pair, each nitrogen is surrounded past seven electrons, and finally, sharing the third, we get a construction where each nitrogen is surrounded by eight electrons; a noble gas configuration (or the "octet rule"). Nitrogen is a very stable molecule and relatively unreactive, existence held together by a strong triple covalent bail.
\[:N\vdots \vdots N:\]
As nosotros have constructed Lewis diagrams, thus far, we accept strived to achieve an octet of electrons around every element. In nature, however, there are many exceptions to the "octet rule". Elements in the outset row of the periodic table (hydrogen and helium) tin simply suit two valence electrons. Elements below the second row in the periodic tabular array can conform, 10, 12 or even xiv valence electrons (we will encounter an example of this in the next section). Finally, in many cases molecules exist with unmarried unpaired electrons. A classic example of this is oxygen gas (O 2 ). Nosotros have previously drawn the Lewis diagram for oxygen with an oxygen-oxygen double bail. Physical measurements on oxygen, however, suggest that this film of bonding is non quite accurate. The magnetic properties of oxygen, O ii , are about consistent with a structure having ii unpaired electrons in the configuration shown below:
\[:\underset{.}{\overset{..}{O}}\cdot \cdot \underset{.}{\overset{..}{O}}:\]
In this Lewis diagram, each oxygen cantlet is surrounded by seven electrons (not eight). This electronic configuration may explain why oxygen is such a reactive molecule (reacting with iron, for instance, to form rust); the unpaired electrons on the oxygen molecule are readily available to interact with electrons on other elements to form new chemical compounds.
Some other notable exception to the "octet rule" is the molecule NO (nitrogen monoxide). Combining one nitrogen (Group 5A) with 1 oxygen (Group 6A) gives a molecule with eleven valence electrons. There is no way to arrange eleven electrons without leaving one electron unpaired. Nitric oxide is an extremely reactive molecule (past virtue of its unshared electron) and has been found to play a central function is biochemistry as a reactive, brusk-lived molecule involved in cellular communication.
\[\cdot \underset{.}{\overset{..}{N}}\cdot \cdot \underset{.}{\overset{..}{O}}:\]
As useful equally Lewis diagrams can be, chemists tire of drawing little dots and, for a shorthand representation of a covalent bond, a short line (a line-bond) is often drawn between the two elements. Whenever you lot see atoms continued by a line-bond, you are expected to understand that this represents 2 shared electrons in a covalent bond. Further, the unshared pairs of electrons on the bonded atoms are sometimes shown, and sometimes they are omitted. If unshared pairs ore omitted, the chemist reading the structure is assumed to understand that they are present.
\[\underset{..}{\overset{..}{F}}\cdot \cdot \underset{..}{\overset{..}{F}}:\; or\; :\underset{..}{\overset{..}{F}}-\underset{..}{\overset{..}{F}}:\; or\; F-F\]
\[:N\vdots \vdots N:\; or\; :N\equiv North:\; or\; N\equiv N\]
Contributors and Attributions
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Paul R. Immature, Professor of Chemical science, University of Illinois at Chicago, Wiki: AskTheNerd; PRY
askthenerd.com - pyounguic.edu; ChemistryOnline.com
Source: https://chem.libretexts.org/Bookshelves/Introductory_Chemistry/Book%3A_Introductory_Chemistry_Online_(Young)/03%3A_Chemical_Bonding_and_Nomenclature/3.02%3A_Covalent_Bonding
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