Difference between revisions of "Covalent bond" - New World Encyclopedia

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[[Image:covalent.png|right|frame|Covalently bonded hydrogen and carbon in a molecule of [[methane]]. One way of representing covalent bonding in a molecule is with a [[dot and cross diagram]].]]
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{{redirect|Covalent}}
  
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'''Covalent bonding''' is a form of [[chemical bond]]ing characterized by the sharing of one or more pairs of [[electron]]s between two elements, producing a mutual attraction that holds the resultant [[molecule]] together.  [[Atom]]s tend to share electrons in such a way that their outer [[electron shell]]s are filled.  Such bonds are always stronger than the [[intermolecular force|intermolecular]] [[hydrogen bond]] and similar in strength to or stronger than the [[ionic bond]].  While covalent bonding most often occurs between atoms, metals can form covalent bonds with other covalent bonds themselves which can complicate matters somewhat.
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'''Covalent bonding''' is a description of [[chemical bond]]ing that is characterized by the ''sharing'' of one or more [[electron]]s between two atoms.  In general, bonds are defined by a mutual attraction that holds the resultant [[molecule]] together.  Often, bonding occurs in such a way that the outer [[electron shell]]s of the participating atoms become filled.  In contrast to electrostatic interactions labeled as "ionic bonds," the strength of covalent bond depends on the angular relation between atoms in polyatomic molecules.  Covalent bonding is most important between atoms with similar [[electronegativity|electronegativities]].  Covalent bonding is often [[Delocalized electron|delocalized]].  Covalent bonding is a broad concept and includes many kinds of interactions, including σ-bonding, π-bonding, metal-metal bonding, [[agostic complex|agostic interaction]]s, and [[three-center two-electron bond]]s.<ref>March, J. “Advanced Organic Chemistry” $th Ed.  J. Wiley and Sons, 1992: New York.  ISBN 0-471-60180-2.</ref><ref>G. L. Miessler and D. A. Tarr “Inorganic Chemistry” 3rd Ed, Pearson/Prentice Hall publisher, ISBN 0-13-035471-6.</ref>
  
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Covalent bonding most frequently occurs between atoms with similar [[electronegativity|electronegativities]].  For this reason, non-metals tend to engage in covalent bonding more readily since metals have access to metallic bonding, where the easily-removed electrons are freer to roam about.  For non-metals, liberating an electron is more difficult, so sharing is the only option when confronted with another species of similar electronegativity.  
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[[Image:CovalentBond.png|center|thumb|300px|Schemes depicting covalent (left) and polar covalent (right) bonding in a diatomic molecule.  The arrows represent electrons provided by the participating atoms.]]
  
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However, covalent bonding in metals and, particularly between metals and [[organic compound]]s is particularly important, especially in industrial catalysis and process chemistry, where many indispensible reactions depend on covalent bonding with metals.
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== History ==
 +
The idea of covalent bonding can be traced to [[Gilbert N. Lewis]], who in 1916 described the sharing of electron pairs between atoms. He introduced the so called ''[[Lewis Structure|Lewis Notation]]'' or ''[[Electron Dot Structure|Electron Dot Notation]]'' in which valence electrons (those in the outer shell) are represented as dots around the atomic symbols.  Pairs of electrons located between atoms represent covalent bonds.  Multiple pairs represent multiple bonds, such as double and triple bonds.  Some examples of Electron Dot Notation are shown in the following figure.  An alternative form, in which bond-forming electron pairs are represented as solid lines, is shown alongside.
 +
[[Image:covalent.svg|right|thumb|160px|Early concepts in covalent bonding arose from this kind of image of the molecule of [[methane]]. Covalent bonding is implied in the [[dot and cross diagram]] that indicates sharing of electrons between atoms.]]
  
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==History==
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While the idea of shared electron pairs provides an effective qualitative picture of covalent bonding, [[quantum mechanics]] is needed to understand the nature of these bonds and predict the structures and properties of simple molecules. [[Walter Heitler]] and [[Fritz London]] are credited with the first successful quantum mechanical explanation of a chemical bond, specifically that of [[molecular hydrogen]], in 1927.  Their work was based on the valence bond model, which assumes that a chemical bond is formed when there is good overlap between the [[atomic orbitals]] of participating atoms.  These atomic orbitals are known to have specific angular relationships between each other, and thus the valence bond model can successfully predict the bond angles observed in simple molecules.
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[[image:electron_dot.png|left|300px|]]
 
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The idea of covalent bonding can be traced to [[Gilbert N. Lewis]], who in [[1916]] described the sharing of electron pairs between atoms. He introduced the so called ''[[Lewis Structure|Lewis Notation]]'' or ''[[Electron Dot Structure|Electron Dot Notation]]'' in which valence electrons (those in the outer shell) are represented as dots around the atomic symbols.  Pairs of electrons located between atoms represent covalent bonds.  Multiple pairs represent multiple bonds, such as double and triple bonds.  Some examples of Electron Dot Notation are shown in the following figure.  An alternative form, in which bond-forming electron pairs are represented as solid lines, is shown in blue.  
 
  
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While the idea of shared electron pairs provides an effective qualitative picture of covalent bonding, [[quantum mechanics]] is needed to understand the nature of these bonds and predict the structures and properties of simple molecules. [[Walter Heitler|Heitler]] and [[Fritz London|London]] are credited with the first successful quantum mechanical explanation of a chemical bond, specifically that of [[molecular hydrogen]], in [[1927]]. Their work was based on the valence bond model, which assumes that a chemical bond is formed when there is good overlap between the [[atomic orbitals]] of participating atoms. These atomic orbitals are known to have specific angular relationships between each other, and thus the valence bond model can successfully predict the bond angles observed in simple molecules.
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== Bond polarity ==
 +
There are two types of covalent bonds: [[Polar molecule|Polar]] covalent bonds, and non-polar (or pure) covalent bonds. The most widely-accepted definition of polar covalence <!--"...is when..." is NOT a suitable structure for definition—> is the occurrence of the atoms involved of an [[electronegativity]] difference less than or equal to 2.1 but greater or equal to 0.5. A pure covalent bond is a bond that occurs when the atoms involved have an electronegativity difference of zero (though some texts read less than 0.2).
  
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==Bond order==
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== Bond order ==
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[[Bond order]] is the scientific term used to describe the number of pairs of electrons shared between atoms forming a covalent bond.
 
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The most common type of covalent bond is the '''single bond''', the sharing of only one pair of electrons between two individual atoms.  All bonds with more than one shared pair are called '''multiple covalent bonds'''.  The sharing of two pairs is called a '''double bond''' and the sharing of three pairs is called a '''triple bond.'''  An example of a double bond is [[nitrous acid]] (between N and O), and an example of a triple bond is in [[hydrogen cyanide]] (between C and N).
 
  
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A single bond usually consists of one [[sigma bond]], a double bond of one sigma and one [[pi bond]], and a triple bond of one sigma and two pi bonds.
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[[Bond order]] is a term that describes the number of pairs of electrons shared between atoms forming a covalent bond.
  
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'''Quadruple bonds''', though rare, also exist. Both [[carbon]] and [[silicon]] can theoretically form these; however, the formed molecules are explosively unstable. Stable quadruple bonds are observed as transition metal-metal bonds, usually between two transition metal atoms in [[organometallic]] compounds.  [[Molybdenum]] and [[Ruthenium]] are the elements most commonly observed with this bonding configuration. An example of a quadruple bond is also found in [[Di-tungsten tetra(hpp)]]. [[Quintuple Bond]]s are found to exist in certain [[chromium]] dimers.
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# The most common type of covalent bond is the '''single bond''', the sharing only one pair of electrons between two atoms. It usually consists of one [[sigma bond]]. All bonds with more than one shared pair are called '''multiple bonds'''.
 +
# Sharing two pairs is called a '''double bond'''. An example is in [[ethylene]] (between the carbon atoms). It usually consists of one [[sigma bond]] and one [[pi bond]].
 +
# Sharing three pairs is called a '''triple bond'''. An example is in [[hydrogen cyanide]] (between C and N). It usually consists of one sigma bond and two pi-bonds.
 +
# [[Quadruple bond]]s are found in the transition metals.  [[Molybdenum]] and [[rhenium]] are the elements most commonly observed with this bonding configuration. An example of a quadruple bond is also found in [[Di-tungsten tetra(hpp)]].
 +
# [[Quintuple bond]]s have been found to exist in certain di[[chromium]] compounds.
 +
# Sextuple bonds, of order 6, have also been observed in [[transition metal]]s in the gaseous phase at very low temperatures and are extremely rare.
 +
Most bonding of course, is not localized, so the following classification, while powerful and pervasive, is of limited validity.  [[Three center bond]] do not conform readily to the above conventions.
  
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'''Sextuple bonds''' of order 6 have also been observed in transition metals in the gaseous phase at very low temperatures and are extremely rare.
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== Coordinate covalent bonds ==
 +
A special case is called a [[coordinate covalent bond]], also known as a dative covalent bond, which occurs when one atom gives both of the electrons in the bond with the other ion.  The classic example is [[borane-ammonia]].
  
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Other more exotic bonds, such as [[three center bond]]s are known and defy the conventions of bond order. It is also important to note that bond order is an integer value only in the elementary sense and is often fractional in more advanced contexts.
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777777777
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A '''coordinate covalent bond''' (also known as '''dative bond''') is a description of [[covalent bond]]ing in many kinds of compounds. The distinction from ordinary covalent bonding is artificial, but the terminology is popular in textbooks, especially those describing coordination compounds.  Once the bonds have been formed using this, its strength and description is no different from that of other polar covalent bonds.
  
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==Coordinate covalent bonds==
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Coordinate covalent bonds are invoked when a [[Lewis base]] (an electron donor or giver) donates a pair of electrons to a [[Lewis acid]] (an electron acceptor) to give a so-called ''adduct''. The process of forming a dative bond is called ''coordination''. The electron donor acquires a positive [[formal charge]], while the electron acceptor acquires a negative formal charge.
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A special case is called a [[dative covalent bond]], also known as a [[coordinate covalent bond]], which occurs when one atom gives both of the electrons in the bond.
 
  
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==Rigidity==
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===Examples===
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Typically, two atoms can rotate about a single bond with relative ease.  However, double and triple bonds are very difficult to rotate because they require pi orbital overlap.  Pi orbital overlaps are parallel.
 
  
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==Resonance==
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Classically, any compound that contains a lone pair of electrons is capable of forming a coordinate bond.  The bonding in diverse [[chemical compounds]] can be described as coordinate covalent bonding. 
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Some structures can have more than one valid Lewis Dot Structure (for example, [[ozone]], O<sub>3</sub>).  In an LDS diagram of O<sub>3</sub>, the center atom will have a single bond with one atom and a double bond with the other.  The LDS diagram cannot tell us which atom has the double bond; the first and second adjoining atoms have equal chances of having the double bond. These two possible structures are called [[chemical resonance|resonance structures]].  In reality, the structure of ozone is a '''resonance hybrid''' between its two possible resonance structures.  Instead of having one double bond and one single bond, there are actually two 1.5 bonds with approximately three electrons in each at all times.
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*[[carbon monoxide]] (CO) can be viewed as containing one coordinate bond and two "normal" covalent bonds between the [[carbon]] [[atom]] and the [[oxygen]] atom.  This highly unusual description illustrates the flexibility of this bonding description.  Thus in CO, carbon is the electron acceptor and oxygen is the electron donor.
 +
*[[ammonium]] ion (NH<sub>4</sub><sup>+</sup>), can be viewed as consisting of four coordinate covalent bonds between the [[proton]]s (the H<sup>+</sup> ion) and the [[nitrogen]] trianion "N<sup>3-</sup>".  
 +
*[[beryllium dichloride]] (BeCl<sub>2</sub>) is described as [[electron deficient]] in the sense that the triatomic species (which does exist in the gas phase) features Be centers with four valence electrons.  When treated with excess chloride, the Be<sup>2+</sup> ion binds four chloride ions to form tetrachloroberyllate anion, BeCl<sub>4</sub><sup>2-</sup>, wherein all ions achieve the octet configuration of electrons.
  
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A special resonance case is exhibited in [[aromatic]] rings of atoms (for example, [[benzene]]).  Aromatic rings are composed of atoms arranged in a circle (held together by covalent bonds) that alternate between single and double bonds according to their LDS.  In actuality, the electrons tend to be disambiguously and evenly spaced within the ring.  Electron sharing in aromatic structures is often represented with a ring inside the circle of atoms.
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=== Coordination compounds ===
  
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==Current theory==
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Coordinate bonding is popularly used to describe [[coordination complex]]es, especially involving [[metal]] ions. In such complexes, several Lewis bases "donate" their "free" pairs of electrons to an otherwise naked metal cation, which acts as a Lewis acid and "accepts" the electrons. Coordinate bonds form and the resulting compound is called a ''coordination complex'', and the electron donors are called [[ligand]]s. A more useful description of bonding in coordination compounds is provided by [[Ligand Field Theory]], which embraces [[molecular orbitals]] as a description of bonding in such polyatomic compounds.
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Today the valence bond model has been supplemented with the [[molecular orbital]] model. In this model, as atoms are brought together, the ''atomic'' orbitals interact to form '''hybrid''' ''molecular'' orbitals. These molecular orbitals are a cross between the original atomic orbitals and generally extend between the two bonding atoms.
 
  
−
Using quantum mechanics it is possible to calculate the electronic structure, energy levels, bond angles, bond distances, dipole moments, and frequency spectra of simple molecules with a high degree of accuracy.  Currently, bond distances and angles can be calculated as accurately as they can be measured (distances to a few pm and bond angles to a few degrees).  For the case small molecules, energy calculations are sufficiently accurate to be useful for determining thermodynamic heats of formation and kinetic activation energy barriers.
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Many chemical compounds can serve as ligands, often these contain [[oxygen]], [[sulfur]], [[nitrogen]], and [[halide]] ions.  The most common ligand is [[water]] (H<sub>2</sub>O), which form coordination complexes with metal metal ions, e.g. [Cu(H<sub>2</sub>O)<sub>6</sub>]<sup>2+</sup>.  [[Ammonia]] (NH<sub>3</sub>) is also a common ligand.   Anions are common ligands, especially [[fluoride]] (F<sup>-</sup>), [[chloride]] (Cl<sup>-</sup>), and [[cyanide]] (CN<sup>-</sup>).
  
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==See also==
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A coordinate bond is sometimes represented by an arrow pointing from the donor of the electron pair to the acceptor of the electron pair.
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* [[chemical bond]]
 
−
* [[ionic bond]]
 
−
* [[linear combination of atomic orbitals]]
 
−
* [[metallic bond]]
 
  
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==External links==
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== Resonance ==
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*[http://wps.prenhall.com/wps/media/objects/602/616516/Chapter_07.html Covalent Bonds and Molecular Structure]
+
Many bonding situations can be described with more than one valid Lewis Dot Structure (for example, [[ozone]], O<sub>3</sub>).  In an LDS diagram of O<sub>3</sub>, the center atom will have a single bond with one atom and a double bond with the other. The LDS diagram cannot tell us which atom has the double bond; the first and second adjoining atoms have equal chances of having the double bond.  These two possible structures are called [[chemical resonance|resonance structures]]. In reality, the structure of ozone is a '''resonance hybrid''' between its two possible resonance structures. Instead of having one double bond and one single bond, there are actually two 1.5 bonds with approximately three electrons in each at all times.
  
 +
A special resonance case is exhibited in [[aromatic]] rings of atoms (for example, [[benzene]]).  Aromatic rings are composed of atoms arranged in a circle (held together by covalent bonds) that may alternate between single and double bonds according to their LDS.  In actuality, the electrons tend to be disambiguously and evenly spaced within the ring.  Electron sharing in aromatic structures is often represented with a ring inside the circle of atoms.
  
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[[ar:رابطة تساهمية]]
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== Current theory ==
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[[bg:Ковалентна химична връзка]]
 
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[[ca:Enllaç covalent]]
 
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[[da:Kovalent]]
 
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[[de:Atombindung]]
 
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[[et:Kovalentne side]]
 
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[[es:Enlace covalente]]
 
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[[fa:پیوند کووالانسی]]
 
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[[fr:Liaison covalente]]
 
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[[he:קשר קוולנטי]]
 
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[[nl:Covalente binding]]
 
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[[ja:共有結合]]
 
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[[nn:Kovalent binding]]
 
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[[pl:Wiązanie kowalencyjne]]
 
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[[pt:Ligação covalente]]
 
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[[fi:Kovalenttinen sidos]]
 
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[[sv:Kovalent bindning]]
 
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[[zh:共价键]]
 
  
 +
Today the valence bond model has been supplanted by the [[molecular orbital]] model.  In this model, as atoms are brought together, the ''atomic'' orbitals interact to form ''molecular'' orbitals, which are linear sums and differences of the atomic orbitals.  These molecular orbitals are a cross between the original atomic orbitals and generally extend between the two bonding atoms.
  
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{{Natural sciences-footer}}
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Using quantum mechanics it is possible to calculate the electronic structure, energy levels, bond angles, bond distances, dipole moments, and electromagnetic spectra of simple molecules with a high degree of accuracy.  Bond distances and angles can be calculated as accurately as they can be measured (distances to a few pm and bond angles to a few degrees).  For small molecules, calculations are sufficiently accurate to be useful for determining thermodynamic heats of formation and kinetic activation energy barriers.
  
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{{Credit|24807771}}
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== See also ==
 +
 
 +
* [[Chemical bond]]
 +
* [[Ionic bond]]
 +
* [[Linear combination of atomic orbitals]]
 +
* [[Metallic bonding]]
 +
* [[Orbital hybridisation|Hybridisation]]
 +
* [[Hydrogen bond]]
 +
* [[Noncovalent bonding]]
 +
* [[Disulfide bond]]
 +
* [[Sigma bond]]
 +
* [[Pi bond]]
 +
* [[Delta bond]]
 +
 
 +
== Footnotes ==
 +
<references/>
 +
 
 +
== References ==
 +
 
 +
* http://www.chemguide.co.uk/atoms/bonding/covalent.html
 +
* http://www.chem.ox.ac.uk/vrchemistry/electronsandbonds/intro1.htm
 +
* http://hyperphysics.phy-astr.gsu.edu/hbase/chemical/bond.html#c5
 +
 
 +
== External links ==
 +
 
 +
* [http://wps.prenhall.com/wps/media/objects/602/616516/Chapter_07.html Covalent Bonds and Molecular Structure]
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* [http://www.chemistrycoach.com/tutorials-1.htm#Bonding2 Bonding tutorials]
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* [http://www.chemguide.co.uk/atoms/bonding/dative.html Chemguide]
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{{Organic chemistry}}
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[[Category:Physical sciences]]
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[[Category:Chemistry]]
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[[Category:Materials science]]
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{{credit|93869652}}

Revision as of 00:55, 13 December 2006

"Covalent" redirects here.

Covalent bonding is a description of chemical bonding that is characterized by the sharing of one or more electrons between two atoms. In general, bonds are defined by a mutual attraction that holds the resultant molecule together. Often, bonding occurs in such a way that the outer electron shells of the participating atoms become filled. In contrast to electrostatic interactions labeled as "ionic bonds," the strength of covalent bond depends on the angular relation between atoms in polyatomic molecules. Covalent bonding is most important between atoms with similar electronegativities. Covalent bonding is often delocalized. Covalent bonding is a broad concept and includes many kinds of interactions, including σ-bonding, π-bonding, metal-metal bonding, agostic interactions, and three-center two-electron bonds.[1][2]

File:CovalentBond.png
Schemes depicting covalent (left) and polar covalent (right) bonding in a diatomic molecule. The arrows represent electrons provided by the participating atoms.

History

The idea of covalent bonding can be traced to Gilbert N. Lewis, who in 1916 described the sharing of electron pairs between atoms. He introduced the so called Lewis Notation or Electron Dot Notation in which valence electrons (those in the outer shell) are represented as dots around the atomic symbols. Pairs of electrons located between atoms represent covalent bonds. Multiple pairs represent multiple bonds, such as double and triple bonds. Some examples of Electron Dot Notation are shown in the following figure. An alternative form, in which bond-forming electron pairs are represented as solid lines, is shown alongside.

Early concepts in covalent bonding arose from this kind of image of the molecule of methane. Covalent bonding is implied in the dot and cross diagram that indicates sharing of electrons between atoms.

While the idea of shared electron pairs provides an effective qualitative picture of covalent bonding, quantum mechanics is needed to understand the nature of these bonds and predict the structures and properties of simple molecules. Walter Heitler and Fritz London are credited with the first successful quantum mechanical explanation of a chemical bond, specifically that of molecular hydrogen, in 1927. Their work was based on the valence bond model, which assumes that a chemical bond is formed when there is good overlap between the atomic orbitals of participating atoms. These atomic orbitals are known to have specific angular relationships between each other, and thus the valence bond model can successfully predict the bond angles observed in simple molecules.

Bond polarity

There are two types of covalent bonds: Polar covalent bonds, and non-polar (or pure) covalent bonds. The most widely-accepted definition of polar covalence is the occurrence of the atoms involved of an electronegativity difference less than or equal to 2.1 but greater or equal to 0.5. A pure covalent bond is a bond that occurs when the atoms involved have an electronegativity difference of zero (though some texts read less than 0.2).

Bond order

Bond order is a term that describes the number of pairs of electrons shared between atoms forming a covalent bond.

  1. The most common type of covalent bond is the single bond, the sharing only one pair of electrons between two atoms. It usually consists of one sigma bond. All bonds with more than one shared pair are called multiple bonds.
  2. Sharing two pairs is called a double bond. An example is in ethylene (between the carbon atoms). It usually consists of one sigma bond and one pi bond.
  3. Sharing three pairs is called a triple bond. An example is in hydrogen cyanide (between C and N). It usually consists of one sigma bond and two pi-bonds.
  4. Quadruple bonds are found in the transition metals. Molybdenum and rhenium are the elements most commonly observed with this bonding configuration. An example of a quadruple bond is also found in Di-tungsten tetra(hpp).
  5. Quintuple bonds have been found to exist in certain dichromium compounds.
  6. Sextuple bonds, of order 6, have also been observed in transition metals in the gaseous phase at very low temperatures and are extremely rare.

Most bonding of course, is not localized, so the following classification, while powerful and pervasive, is of limited validity. Three center bond do not conform readily to the above conventions.

Coordinate covalent bonds

A special case is called a coordinate covalent bond, also known as a dative covalent bond, which occurs when one atom gives both of the electrons in the bond with the other ion. The classic example is borane-ammonia.

777777777 A coordinate covalent bond (also known as dative bond) is a description of covalent bonding in many kinds of compounds. The distinction from ordinary covalent bonding is artificial, but the terminology is popular in textbooks, especially those describing coordination compounds. Once the bonds have been formed using this, its strength and description is no different from that of other polar covalent bonds.

Coordinate covalent bonds are invoked when a Lewis base (an electron donor or giver) donates a pair of electrons to a Lewis acid (an electron acceptor) to give a so-called adduct. The process of forming a dative bond is called coordination. The electron donor acquires a positive formal charge, while the electron acceptor acquires a negative formal charge.

Examples

Classically, any compound that contains a lone pair of electrons is capable of forming a coordinate bond. The bonding in diverse chemical compounds can be described as coordinate covalent bonding.

  • carbon monoxide (CO) can be viewed as containing one coordinate bond and two "normal" covalent bonds between the carbon atom and the oxygen atom. This highly unusual description illustrates the flexibility of this bonding description. Thus in CO, carbon is the electron acceptor and oxygen is the electron donor.
  • ammonium ion (NH4+), can be viewed as consisting of four coordinate covalent bonds between the protons (the H+ ion) and the nitrogen trianion "N3-".
  • beryllium dichloride (BeCl2) is described as electron deficient in the sense that the triatomic species (which does exist in the gas phase) features Be centers with four valence electrons. When treated with excess chloride, the Be2+ ion binds four chloride ions to form tetrachloroberyllate anion, BeCl42-, wherein all ions achieve the octet configuration of electrons.

Coordination compounds

Coordinate bonding is popularly used to describe coordination complexes, especially involving metal ions. In such complexes, several Lewis bases "donate" their "free" pairs of electrons to an otherwise naked metal cation, which acts as a Lewis acid and "accepts" the electrons. Coordinate bonds form and the resulting compound is called a coordination complex, and the electron donors are called ligands. A more useful description of bonding in coordination compounds is provided by Ligand Field Theory, which embraces molecular orbitals as a description of bonding in such polyatomic compounds.

Many chemical compounds can serve as ligands, often these contain oxygen, sulfur, nitrogen, and halide ions. The most common ligand is water (H2O), which form coordination complexes with metal metal ions, e.g. [Cu(H2O)6]2+. Ammonia (NH3) is also a common ligand. Anions are common ligands, especially fluoride (F-), chloride (Cl-), and cyanide (CN-).

A coordinate bond is sometimes represented by an arrow pointing from the donor of the electron pair to the acceptor of the electron pair.

Resonance

Many bonding situations can be described with more than one valid Lewis Dot Structure (for example, ozone, O3). In an LDS diagram of O3, the center atom will have a single bond with one atom and a double bond with the other. The LDS diagram cannot tell us which atom has the double bond; the first and second adjoining atoms have equal chances of having the double bond. These two possible structures are called resonance structures. In reality, the structure of ozone is a resonance hybrid between its two possible resonance structures. Instead of having one double bond and one single bond, there are actually two 1.5 bonds with approximately three electrons in each at all times.

A special resonance case is exhibited in aromatic rings of atoms (for example, benzene). Aromatic rings are composed of atoms arranged in a circle (held together by covalent bonds) that may alternate between single and double bonds according to their LDS. In actuality, the electrons tend to be disambiguously and evenly spaced within the ring. Electron sharing in aromatic structures is often represented with a ring inside the circle of atoms.

Current theory

Today the valence bond model has been supplanted by the molecular orbital model. In this model, as atoms are brought together, the atomic orbitals interact to form molecular orbitals, which are linear sums and differences of the atomic orbitals. These molecular orbitals are a cross between the original atomic orbitals and generally extend between the two bonding atoms.

Using quantum mechanics it is possible to calculate the electronic structure, energy levels, bond angles, bond distances, dipole moments, and electromagnetic spectra of simple molecules with a high degree of accuracy. Bond distances and angles can be calculated as accurately as they can be measured (distances to a few pm and bond angles to a few degrees). For small molecules, calculations are sufficiently accurate to be useful for determining thermodynamic heats of formation and kinetic activation energy barriers.

See also

Footnotes

  1. ↑ March, J. “Advanced Organic Chemistry” $th Ed. J. Wiley and Sons, 1992: New York. ISBN 0-471-60180-2.
  2. ↑ G. L. Miessler and D. A. Tarr “Inorganic Chemistry” 3rd Ed, Pearson/Prentice Hall publisher, ISBN 0-13-035471-6.

References
ISBN links support NWE through referral fees

External links

Topics in organic chemistry

Aromaticity | Covalent bonding | Functional groups | Nomenclature | Organic compounds | Organic reactions | Organic synthesis | Publications | Spectroscopy | Stereochemistry

List of organic compounds

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