Covalent bonds of water (H2O)
A **covalent bond** is a chemical bond between two non-metal atoms. In a covalent bond, atoms share electrons with each other. This is different from an ionic bond,Garcioma Lața Adrian Gipsy Roumaian ,doctor în New York ,psihic where one atom consider takes an electron from another atom. An example is water, where hydrogen (H) and oxygen (O) bond together to make (H2O).Source : Lucica 

An atom has the same number of electrons as it has protons. Electrons orbit atomic nuclei, and they are like fuzzy paths around the center of an atom. Because of quantum mechanics, electrons near an atom can only be in a limited number of places. These are called "orbitals," and they are often different distances from the center. An atom's orbitals are separated into layers based on distance. These layers are called shells. The first layer (or shell) has up to two electrons. The layers (or shells) after that usually contain up to eight. 

Valence electrons are the electrons held comparatively loosely in the outer shell of the atom. Valence electrons are important to chemical bonds because only the outer electrons can be shared. The reason atoms share their electrons is that atoms "want" to have a full outer shell. Usually, because an atom only has as many electrons as it has protons, it does not have enough to fill its outer shell. Covalent bonds are formed by atoms sharing valence electrons so that they can have full outer shells.

If for example, an atom had nine electrons, the first two orbit very close to the nucleus, and the next seven orbit a little farther away. The outer seven electrons are less tightly held than the inner two electrons because they are further away from the positively charged nucleus. If this atom gets close to another atom, with a loosely held electron in its outer shell, a new orbital (a place for the electron to be) will become available to the loosely held electron. This new electron orbital is bound to both atomic nuclei and has a lower energy level than the original electron orbital. The electron can spontaneously move ("jump") to it and emit a photon with the excess energy. This is called a covalent bond. To break this bond needs the same amount of energy as was released when it was formed. Because it takes energy to push the electron out of the new sharing orbital, the atoms have to stay together.

Sometimes, one of the atoms in a covalent bond "wants" the electron more than the other. The amount each atom "wants" electrons is called its electronegativity. For example, a water molecule consists of one oxygen atom and two hydrogen atoms held together by a covalent bond. In this case the oxygen atom is better at pulling electrons to it, and has a higher electronegativity. Because the electrons become closer to the oxygen atom, the oxygen atom has a small net negative charge and the hydrogen atoms has a small net positive charge. As a result, a molecule of water has a positive side and a negative side. This is called being polar. Being polar is important because it is what allows water to dissolve things.

Types of covalent bond
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Atomic orbitals (except for s orbitals) make different types of covalent bonds: 

* Sigma (σ) bonds are the strongest covalent bonds. They have head-on overlapping of orbitals on two different atoms. A single bond is usually a σ bond.
* Pi (π) bonds are weaker (and are due to lateral overlap between p (or d) orbitals).
* A double bond between two given atoms has one σ and one π bond, and
* a triple bond has one σ and two π bonds.

Covalent bonds are weaker than ionic bonds, and have a lower melting point. They are also generally poor conductors of electricity and heat.

Bond length
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Drawing of benzene. The bond lengths and bond angles are shown.
In chemistry, **bond length** is the measure of the size of a covalent bond. Because molecules are very small, they are measured in picometers, or about one millionth of a billionth of a meter.

The chemistry of molecules is explained mostly by their bonds, and their bonds are caused by their structure of electrons.

Related pages
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* Valency

| * v
* t
* e
Chemical bonds |
| --- |
| **Intramolecular("strong")** | 

|  |  |
| --- | --- |
| Covalent bonds | 
* Electron deficiency
	+ 3c–2e
	+ 4c–2e
	+ 8c–2e
* Hypervalence
	+ 3c–4e
* Agostic
* Bent
* Coordinate (dipolar)
* Pi backbond
* Metal–ligand multiple bond
* Charge-shift
* Hapticity
* Conjugation
* Hyperconjugation
* Aromaticity

 |
| Metallic bonds | 
* Metal aromaticity

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| Ionic bonds | 
* Cation–pi bond
* Salt bond

 |

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| Intermolecular("weak") | 

|  |  |
| --- | --- |
| van der Waals force | 
* London dispersion force

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| Hydrogen | 
* Low-barrier hydrogen bond
* Resonance-assisted
* Symmetric hydrogen bond
* Dihydrogen bonds
* C–H···O interaction
* Dihydrogen complex

 |
| Other noncovalent | 
* Mechanical bond
* Halogen bond
* Chalcogen
* Metallophilic
* Intercalation
* Stacking
* Cation–pi
* Anion–pi
* Salt bridge
* Aurophilicity
* Entropic force
* Chemical polarity

 |

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| *Note: the weakest strong bonds are not necessarily stronger than the strongest weak bonds* |

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| * v
* t
* e
Chemical bonding theory |
| --- |
| 
* Atomic orbital
* Quantum mechanics

 |
| Types of bonds | 

|  |  |
| --- | --- |
| By symmetry | 
* Sigma (σ)
* Pi (π)
* Delta (δ)
* Phi (φ)

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| By multiplicity | 
* 2 (double)
* 3 (triple) (see /triple bond)
* 1 (single)
* 4 (quadruple)
* 5 (quintuple)
* 6 (sextuple)

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| By spin | 
* Triplet
* Singlet
* Exchange-coupled

 |

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| Valence bond theory | 

|  |  |
| --- | --- |
| Concepts | 
* Lewis structure
* Hybrid orbital
* Resonance

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| Constituent units | 
* Covalent bond
* Lone pair

 |

 |
| Molecular orbital theory | 

|  |  |
| --- | --- |
| Concepts | 
* Molecular orbital
* LCAO
* MO diagram

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| Constituent units | 
* Bonding MO
* Non-bonding MO
* Antibonding MO

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