Covalent bonds: formed from sharing of two electrons, usually one donated from each of the two bonding atoms. Sharing electrons is one way that atoms can satisfy the "Octet Rule" which as stated by Gilbert Lewis "atoms, by sharing electrons to form an electron-pair bond, can acquire a stable, noble-gas structure".
There are 3 kinds of covalent bonds based on number of pairs of electrons shared between the two atoms:
Bond energy: single < double < triple
- single covalent bond - 1 shared pair
- double covalent bond - 2 shared pairs
- triple covalent bond - 3 shared pairs
Bond length: single < double < triple
Lewis Structures: The simplest covalent bond is between two hydrogen atoms. Electron-dot (or Lewis) symbol for hydrogen is H.. The two electrons from two hydrogen atoms pair up to make H:H. The pair of dots can be replaced by a dash to represent a bond, as in H-H. Hydrogen as the element exists as these "diatomic molecules" and its molecular formula is shown as H2.
The halogens also exist as diatomic molecules. Fluorine, chlorine, bromine and iodine all have the same type of Lewis structure, with seven outer electrons (2 on the right side, 2 on the top, 2 on the left side, and 1 on the bottom, so two F atoms share their unpaired electrons with each other to form the single covalent bond, F:F or F-F, to get F2. The rest of the pairs of electrons around each halogen are called "unshared pairs". The number of paired and unpaired electrons around an atom play an important part in determining the shape of a molecule, which in turn contribute to chemical and physical properties of the molecule.
Oxygen atoms, with six outer level electrons each ( 2 on the right side, 2 on the top, and 1 each on left side and bottom), share two pairs of electrons and thus have a double bond between them. This is shown by O::O or O=O for O2.
Nitrogen atoms, with only five outer electrons has 2 on the right side, and 1 each on the top, left side and bottom. Two nitrogen atoms share three electrons each, to form three pairs between them. This bonding is written as :N:::N:, N=N for N2
Things get a bit more complicated when atoms of different kinds of elements bond. One of the simplest bonds is between hydrogen and chlorine. The single unpaired electron on H pairs up with the unpaired electron on the chlorine and the result is H:Cl or H-Cl, or HCl.
Writing/Drawing Lewis Structures for simple molecules.
- Count the number of valence electrons in each of the atoms in the molecule.
- Count the total number of electrons that are needed to give each atom in the molecule an octet in the valence level, or two electrons for hydrogen.
- Find the difference between the number of electrons available for sharing (#1) and the total number needed to complete the octet for each (#2). This gives the number of electrons that will be shared among the bonding atoms.
- Divide the number of shared electrons by 2 to give the number of electron pairs, thus the number of bonds between all of the atoms in the molecules. A double or triple bond count as two or three bonds because of the number of electron pairs they represent.
There are several rules to remember when writing Lewis structures:
- H is always terminal, only one single bond per hydrogen
- Oxygen almost always has two shared pairs and two unshared pairs of electrons. The two shared pairs may be in two separate single bonds or in one double bond.
- Carbon atoms will always have 4 bonds, either as 4 separate single bonds, or a combination of single, double and/or triple bonds so that there are 4 pairs of electrons around each carbon atom in the molecule. Carbon atoms may single, double, or triple bond to each other, single or double bond to oxygen atoms, or single, double, or triple bond to nitrogen atoms. Halogen atoms form single bonds only, and usually form just one bond. An example of an exception is the chlorate, ClO31-, ion.
Send questions, comments or suggestions to
Gwen Sibert, at the
Roanoke Valley Governor's School
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