Béda révisi "Beungkeut kovalén"

11 bita ditambahkeun ,  4 tahun yang lalu
m
Ngarapihkeun éjahan, replaced: rea → réa (4), ea → éa (8), eo → éo using AWB
(Removing Link FA template as it is now available in wikidata)
m (Ngarapihkeun éjahan, replaced: rea → réa (4), ea → éa (8), eo → éo using AWB)
'''Beungkeut kovalén''' hartina [[beungkeut kimia]] nu cirina ''babagi'' hiji atawa leuwih [[éléktron]] antara dua [[atom]]. Sacara umum, beungkeut téh ditangtukeun ku ayana pabetot-betot atawa daya tarik nu ngajadikeun kabentukna hiji [[molekul]]. Mindengna mah ieu beungkeut kabentuk ku dieusina [[kulit éléktron]] luar. Sabalikna ti interaksi éléktrostatik dina [[beungkeut ionik]], kakuatan beungkeut kovalén mah gumantung kana hubungan antara atom-atom dina molekul poliatomik. <!--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>
 
[[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.]]
 
== History ==
The ideaidéa 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.]]
 
While the ideaidéa 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 eachéach 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 molecule|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 greatergréater 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 readréad less than 0.2).
 
== Bond order ==
# [[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 gaseousgaséous 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 readilyréadily to the above conventions.
 
== Coordinate covalent bonds ==
 
== Resonance ==
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 realityréality, the structure of ozone is a '''resonance hybrid''' between its two possible resonance structures. InsteadInstéad of having one double bond and one single bond, there are actually two 1.5 bonds with approximately three electrons in eachéach 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 linearlinéar 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 measuredméasured (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 heatshéats of formation and kinetic activation energy barriers.
 
== References ==
 
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