Initially, the term\ncovalence concerning bonding was used in 1919, by Irving Langmuir in an\nAmerican journal known as the Journal of the American Chemical Society under\nthe title of “The Arrangement of Electrons in Atoms and Molecules.”\nLangmuir wrote in his journal that \u201cwe shall denote by the term covalence the\nnumber of pairs of electrons that a given atom shares with its neighbors.\u201d<\/p>\n\n\n\n
The idea of covalent\nbonding is very old. Before 1919, several years back, it was known to Gilbert\nN. Lewis, who in 1916 explains the sharing of electron pairs between atoms.\nLewis introduced the lewis dot structure, or Lewis notation or electron dot\nnotation, in which the valence electrons i.e. those present in an outer shell\nwas represented by dots around the atomic symbol and the pairs of electrons\nthat were located between the atoms were represented by the covalent bonds. The\nmultiple pairs represent multiple bonds for example double bonds and triple\nbonds.<\/p>\n\n\n\n
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<\/figure><\/div>\n\n\n\n<\/a><\/p>\n\n\n\n The old concept of a covalent bond has been\nrepresented by using Methane molecule. A covalent bond is represented in the\nLewis structure that shows electron sharing between atoms. <\/p>\n\n\n\n A covalent bond which is additionally known as a\nmolecular bond is a chemical bond that involves the sharing of electron pairs\nbetween atoms. <\/em><\/p>\n\n\n\n Covalent bonds are mostly\nformed between two non-metals. For example, water, where hydrogen (H) and\noxygen (O) forms a bond to make H2<\/sub>O. Here, the oxygen atom shares\none electron with each hydrogen atom. It shows that the oxygen has a bit\npositive charge and that of hydrogen has a small negative charge which as a\nresult are attracted to each other due to the presence of electromagnetic\nforce.<\/p>\n\n\n\n Atomic orbitals have\nspecific directional properties that lead towards different types of a covalent\nbond. Sigma (\u03c3<\/strong>,) bonds are the strongest type\nof covalent bonds and this is due to head-on overlapping of orbitals on two\ndifferent atoms. A single bond is normally called a sigma bond. The pi (\u03c0) bonds are relatively weaker\nand this is due to overlapping between p (or d) orbitals. A double bond between\nthe given two atoms consists of one sigma and one pi bond while a triple bond\nconsists of one sigma and two pi bonds.<\/p>\n\n\n\n A double covalent bond is the type of chemical bond\nin which two electron pairs are shared between the two atoms.<\/em><\/p>\n\n\n\n This type of covalent\nbond includes four bonding electrons between atoms rather than the usual two\nbonding electrons that are involved in a single bond. Double covalent bonds are\nconsidered reactive bonds because of the presence of a large number of\nelectrons. The double bonds are shorter and much stronger than single bonds.<\/p>\n\n\n\n Double covalent bonds can\nbe represented as two parallel lines. A sign of equal is used to show a double\nbond in a formula. In the mid-19th<\/sup> century. A Russian chemist\nAlexander Butlerov tells us the structural formula of a double covalent bond.<\/p>\n\n\n\n <\/a><\/strong><\/p>\n\n\n\n \nGeometrical representation of ethylene<\/p>\n\n\n\n Double covalent bonds are\ncreated by the sharing four electrons of the two atoms in order to fulfill the\noctet rule. The octet rule tells that the atoms will share, lose or gain their\nelectrons to achieve eight valence electrons. Double covalent bonds are usually\nmade from nonmetals like carbon, nitrogen, and oxygen. <\/p>\n\n\n\n Electrons move freely and\nvery fast around an atom. Depending on the energy level and closeness to the\nnucleus, the electrons will stay within a certain shape. The specific shape\ngained by the electrons is called electron orbitals. There are four shapes of\nelectron orbitals namely s, p, d, and f. in double covalent bonds s and p\norbitals are involved.<\/p>\n\n\n\n Some atoms donate more\nelectrons while some donate much less to a double covalent bond. They shuffle\nthemselves around the present electrons to share four electrons.<\/p>\n\n\n\n Examples with simple molecules<\/strong><\/p>\n\n\n\n Oxygen,\nO2<\/sub>:<\/strong><\/p>\n\n\n\n The two atoms of oxygen can gain a stable structure by sharing two pairs of electrons as shown in a figure below.<\/p>\n\n\n\n Carbon\ndioxide, CO2<\/sub>:<\/strong><\/p>\n\n\n\n Ethene,\nC2<\/sub>H4<\/sub>:<\/strong><\/p>\n\n\n\n Ethene forms a double covalent bond between two carbon atoms. As shown in the above figure.<\/p>\n\n\n\n Ethene is made from\nhydrogen atom (1s1<\/sup>) and carbon atoms (1s2<\/sup> 2s2<\/sup>\n2px<\/sub> 1<\/sup>2py<\/sub>1).<\/p>\n\n\n\n Promotion\nof an electron:<\/strong><\/p>\n\n\n\n The carbon atom doesn\u2019t\nhave satisfied unpaired electrons to form the required number of bonds. So it\nrequires to donate one of the 2s2<\/sup> pair to the empty orbit 2pz<\/sub>.\nThe same thing happens with the formation of a carbon bond.<\/p>\n\n\n\n Now, this carbon atom is in an excited state.<\/em><\/p>\n\n\n\n The newly formed orbitals\nare known as sp2 hybrids because they are formed by a combination of one s and\ntwo p orbitals. These sp2 hybrid orbitals arrange themselves as far as they can\ni.e. at 120\u1d52 to each other in a plane. The remaining p orbitals align\nthemselves at right angles to them.<\/p>\n\n\n\n Ethene,\nC2<\/sub>H4<\/sub><\/strong><\/p>\n\n\n\n Ethene is comprised of\nfour hydrogen atoms and two carbon atoms and before joining of these molecules,\nthe arrangement would look like the form given below:<\/p>\n\n\n\n The different atomic\norbitals that are pointing towards each other will now merge to give molecular\norbitals. Each orbit will contain a bonding pair of electrons. The molecular\norbitals that are formed by end-to-end overlapping of atomic orbitals are known\nas sigma bonds.<\/p>\n\n\n\n In the above diagram, the\nblack dots are representing the nuclei of the atoms.<\/p>\n\n\n\n It is also noticed that\nthe p orbitals are very close to each other and they are overlapping sideways.<\/p>\n\n\n\n This sideway overlapping\ncreates a molecular orbital but it is of a different kind. Here the electrons\naren’t held on a line between the two nuclei, above and below the molecule. A\nformation of bond in this way is called a pi bond.<\/p>\n\n\n\n Note:<\/strong><\/p>\n\n\n\n The sigma bond is\nrepresented by using the lines \u2013 and each line sows one pair of shared\nelectrons. The other lines represent the direction of the bond in it. In a pi\nregion, two electrons make up a bond and these two electrons can move freely\nwithin that space.<\/p>\n\n\n\n https:\/\/www.chemguide.co.uk\/atoms\/bonding\/doublebonds.html<\/a><\/p>\n\n\n\n https:\/\/en.m.wikipedia.org\/wiki\/Carbon%E2%80%93carbon_bond<\/a><\/p>\n\n\n\n https:\/\/study.com\/academy\/lesson\/double-bond-definition-formation-example.html<\/a><\/p>\n\n\n\n https:\/\/en.wikipedia.org\/wiki\/Covalent_bond<\/a><\/p>\n\n\n\n https:\/\/www.thoughtco.com\/definition-of-double-bond-605044<\/a><\/p>\n\n\n\n https:\/\/en.wikipedia.org\/wiki\/Double_bond<\/a><\/p>\n","protected":false},"excerpt":{"rendered":" History Initially, the term covalence concerning bonding was used in 1919, by Irving Langmuir in an American journal known as the Journal of the American Chemical Society under the title of “The Arrangement of Electrons in Atoms and Molecules.” Langmuir wrote in his journal that \u201cwe shall denote by the term covalence the number of … Read more<\/a><\/p>\n","protected":false},"author":6,"featured_media":0,"comment_status":"closed","ping_status":"open","sticky":false,"template":"","format":"standard","meta":{"rating_form_position":"","rating_results_position":"","mr_structured_data_type":""},"categories":[6],"tags":[],"multi-rating":{"mr_rating_results":[{"adjusted_star_result":4.5,"star_result":4.5,"total_max_option_value":5,"adjusted_score_result":4.5,"score_result":4.5,"percentage_result":90,"adjusted_percentage_result":90,"count":4,"post_id":850}]},"_links":{"self":[{"href":"https:\/\/chemdictionary.org\/wp-json\/wp\/v2\/posts\/850"}],"collection":[{"href":"https:\/\/chemdictionary.org\/wp-json\/wp\/v2\/posts"}],"about":[{"href":"https:\/\/chemdictionary.org\/wp-json\/wp\/v2\/types\/post"}],"author":[{"embeddable":true,"href":"https:\/\/chemdictionary.org\/wp-json\/wp\/v2\/users\/6"}],"replies":[{"embeddable":true,"href":"https:\/\/chemdictionary.org\/wp-json\/wp\/v2\/comments?post=850"}],"version-history":[{"count":5,"href":"https:\/\/chemdictionary.org\/wp-json\/wp\/v2\/posts\/850\/revisions"}],"predecessor-version":[{"id":873,"href":"https:\/\/chemdictionary.org\/wp-json\/wp\/v2\/posts\/850\/revisions\/873"}],"wp:attachment":[{"href":"https:\/\/chemdictionary.org\/wp-json\/wp\/v2\/media?parent=850"}],"wp:term":[{"taxonomy":"category","embeddable":true,"href":"https:\/\/chemdictionary.org\/wp-json\/wp\/v2\/categories?post=850"},{"taxonomy":"post_tag","embeddable":true,"href":"https:\/\/chemdictionary.org\/wp-json\/wp\/v2\/tags?post=850"}],"curies":[{"name":"wp","href":"https:\/\/api.w.org\/{rel}","templated":true}]}}Covalent\nbond and its types<\/strong><\/h2>\n\n\n\n
Definition\nof a double covalent bond:<\/strong><\/h2>\n\n\n\n
Representation:<\/strong><\/h2>\n\n\n\n
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<\/figure><\/div>\n\n\n\nFormation\nof a double covalent bond:<\/strong><\/h2>\n\n\n\n
<\/figure>\n\n\n\n<\/figure><\/div>\n\n\n\n<\/figure><\/div>\n\n\n\n<\/figure><\/div>\n\n\n\nReferences:<\/strong><\/h4>\n\n\n\n