CHEMICAL BONDING.pptx

CHEMICAL BONDING Diadem L. Cruz III- St. Matthew

Introduction  A chemical chemical bond bond is an attraction between between atoms that that allows the formation formation of chemical substances substances that contain two or more atoms. The bond is caused by the electrostatic force of attraction between opposite opposite chares! either between electrons and nuclei! nuclei! or as the result of a dipole attraction. The strenth of chemical bonds "aries considerably# there are $stron bonds$ such as co"alent or ionic bonds and $wea% bonds$ such as dipole&d dipole&dipole ipole interactions! the London dispersion force and hydroen bondin.

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Since opposite chares attract "ia a simple electromanetic electromanetic force! the neati"el neati"ely y chared electrons that are orbitin the nucleus and the positi"ely chared protons in the nucleus attract each other. An electron positioned between two nuclei will be attracted to both of them! and the nuclei will be attracted toward electrons in this position. This attraction constitutes the chemical bond. Due to the matter wa"e nature of electrons and their smaller mass! they must occupy a much larer amount of "olume compared with the nuclei! and this "olume occupied by the electrons %eeps the atomic nuclei relati"ely far apart! as compared with the size of the nuclei themsel"es. This phenomenon limits the distance between nuclei and atoms in a bond.

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In eneral! stron chemical bondin is associated with with the sharin or tr ansfer of electrons between the participatin atoms. The atoms in molecules! crystals! metals and diatomic ases' indeed most of the physical en"ironment around us' are held toether by chemical bonds! which dictate the structure and the bul% properties of matter.

Overview of main types of chemica !onds  A chemical bond is an attraction between atoms. This attraction may be seen as the result of different beha"iors of the outermost electrons of atoms. Althouh all of these beha"iors mere into each other seamlessly in "arious bondin situations so that there is no clear line to be drawn between them! the beha"iors of atoms become so (ualitati"ely different as the character of the bond chanes (uantitati"ely! that it remains useful and customary to differentiate between the bonds that cause these different properties of condensed matter.

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In the simplest "iew of a so-called )co"alent) bond! one or more electrons *often a pair of electrons+ are drawn into the space between the two atomic nuclei. ,ere the neati"ely chared electrons are attracted to the positi"e chares of both nuclei! instead of ust their own. This o"ercomes the repulsion between the two positi"ely chared nuclei of the two atoms! and so this o"erwhelmin attraction holds the two nuclei in a fied confiuration of e(uilibrium! e"en thouh they will still "ibrate at e(uilibrium position. Thus! co"alent bondin in"ol"es sharin of electrons in which the positi"ely chared nuclei of two or more atoms simultaneously attract the neati"ely chared electrons that are bein shared between them. These bonds eist between two particular identifiable atoms! and ha"e a direction in space! allowin them to be shown as sinle connectin lines between atoms in drawins! or modeled as stic%s between spheres in models. In a polar co"alent bond! one or more electrons are une(ually shared between two nuclei. Co"alent bonds often result in the formation of small collections of better-connected atoms called molecules! which in solids and li(uids are bound to other molecules by forces that are often much wea%er than the co"alent bonds that hold the molecules internally toether. Such wea% intermolecular bonds i"e oranic molecular substances! such as waes and oils! their soft bul% character! and their low meltin points *in li(uids! molecules must cease most structured or oriented contact with each other+. / hen co"alent bonds lin% lon chains of atoms in lare molecules! howe"er *as in polymers such as nylon+! or when co"alent bonds etend in networ%s throuh solids that are not composed of discrete molecules *such as diamond or (uartz or the silicate minerals in many types of roc%+ then the structures that result may be both stron and touh! at least in the direction oriented correctly with networ%s of co"alent bonds. Also! the meltin points of such co"alent polymers and networ%s increase reatly.

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In a simplified "iew of an ionic bond! the bondin electron is not shared at all! but transferred. In this type of bond! the outer atomic orbital of one atom has a "acancy which allows addition of one or more electrons. These newly added electrons potentially occupy a lower enery-state *effecti"ely closer to m ore nuclear chare+ than they eperience in a different atom. Thus! one nucleus offers a more tihtly bound position to an electron than does another nucleus! with the result that one atom may transfer an electron to the other. This transfer causes one atom to assume a net positi"e chare! and the other to assume a net neati"e chare. The bond then results from electrostatic attraction between atoms! and the atoms become positi"e or neati"ely chared ions. Ionic bonds may be seen as etreme eamples of polarization in co"alent bonds. 0ften! such bonds ha"e no particular orientation in space! since they result from e(ual electrostatic attraction of each ion to all ions around them. Ionic bonds are stron *and thus ionic substances re(uire hih temperatures to melt+ but also brittle! since the forces between ions are short-rane! and do not easily bride crac%s and fractures. This type of bond i"es rise to the physical characteristics of crystals of classic mineral salts! such as table salt.

 A less often mentioned type of bondin is the metallic bond. In this type of bondin! each atom in a metal donates one or more electrons to a $sea$ of electrons that reside between many metal atoms. In this sea! each electron is free *by "irtue of its wa"e nature+ to be associated with a reat many atoms at once. The bond results because the metal atoms become somewhat positi"ely chared due to loss of their electrons! while the electrons remain attracted to many atoms! without bein part of any i"en atom. Metallic bondin may be seen as an etreme eample of delocalization of electrons o"er a lare system of co"alent bonds! in which e"ery atom participates. This type of bondin is often "ery stron *resultin in the tensile strenth of metals+. ,owe"er! metallic bonds are more collecti"e in nature than other types! and so they allow metal crystals to more easily deform! because they are composed of atoms attracted to each other! but not in any particularly-oriented ways. This results in the malleability of metals. The sea of electrons in metallic bonds causes the characteristically ood electrical and thermal conducti"ity of metals! and also their $shiny$ reflection of most fre(uencies of white liht.

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 All bonds can be eplained by (uantum theory! but! in practice! simplification rules allow chemists to predict the strenth! directionality! and polarity of bonds. The octet rule and 1S234 theory are two eamples. More sophisticated theories are "alence bond theory which includes orbital hybridization and resonance! and the linear combination of atomic orbitals molecular orbital method which includes liand field theory. 2lectrostatics are used to describe bond polarities and the effects they ha"e on chemical substances.

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HI"#O$% 

2arly speculations into the nature of the chemical bond! from as early as the 56th century! supposed that certain types of chemical species were oined by a type of chemical affinity. In 5789! Isaac :ewton famously outlined his atomic bondin theory! in $;uery ! on the heels of the in"ention of the "oltaic pile! ?@ns ?a%ob erzelius de"eloped a theory of chemical combination stressin the electroneati"e and electropositi"e character of the combinin atoms. y the mid 5>th century! 2dward Bran%land! B.A. e%ul! A.S. Couper! Aleander utlero"! and ,ermann olbe! buildin on the theory of radicals! de"eloped the theory of "alency! oriinally called $combinin power$! in which compounds were oined owin to an attraction of positi"e and neati"e poles. In 5>5E! chemist Filbert :. Lewis de"eloped the concept of the electron-pair bond! in which two atoms may share one to si electrons! thus formin the sinle electron bond! a sinle bond! a double bond! or a triple bond# in Lewis)s own words! $An electron may form a part of the shell of two different atoms and cannot be said to belon to either one eclusi"ely.$

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That same year! /alther ossel put forward a theory similar to Lewis) only his model assumed complete transfers of electrons between atoms! and was thus a model of ionic bonds. oth Lewis and ossel structured their bondin models on that of Abe)s rule *5>89+.

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In 5>67! the first mathematically complete (uantum description of a simple chemical bond! i.e. that produced by one electron in the hydroen molecular ion! ,6G! was deri"ed by the Danish physicist 0y"ind urrau.H6 This wor% showed that the (uantum approach to chemical bonds could be fundamentally and (uantitati"ely correct! but the mathematical methods used could not be etended to molecules containin more than one electron. A more practical! albeit less (uantitati"e! approach was put forward in the same year by /alter ,eitler and Britz London. The ,eitler-London method forms the basis of what is now called "alence bond theory. In 5>6>! the linear combination of atomic orbitals molecular orbital method *LCA0+ approimation was introduced by Sir ?ohn Lennard-?ones! who also suested methods to deri"e electronic structures of molecules of B6 *fluorine+ and 06 *oyen+ molecules! from basic (uantum principles. This molecular orbital theory represented a co"alent bond as an orbital formed by combinin the (uantum mechanical Schr@diner atomic orbitals which had been hypothesized for electrons in sinle atoms. The e(uations for bondin electrons in multi-electron atoms could not be s ol"ed to mathematical perfection *i.e.! analytically+! but approimations for them still a"e many ood (ualitati"e predictions and results. Most (uantitati"e calculations in modern (uantum chemistry use either "alence bond or molecular orbital theory as a startin point! althouh a third approach! Density Bunctional Theory! has become increasinly popular in recent years.

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In 5>J> failed to ade(uately address the problems that appeared to be better understood by molecular orbital theory. The impact of "alence theory declined durin the 5>E8s and 5>78s as molecular orbital theory rew in usefulness as it was implemented in lare diital computer prorams. Since the 5>=8s! the more difficult problems of implementin "alence bond theory into computer prorams ha"e been sol"ed larely! and "alence bond theory has seen a resurence. 

Comparison of vaence !ond and moecuar or!ita theory 

In some respects "alence bond theory is superior to molecular orbital theory. /hen applied to the simplest two-electron molecule! ,6! "alence bond theory! e"en at the simplest ,eitler-London approach! i"es a much closer approimation to the bond enery! and it pro"ides a much more accurate representation of the beha"ior of the electrons as chemical bonds are formed and bro%en. In contrast simple molecular orbital theory predicts that the hydroen molecule dissociates into a linear superposition of hydroen atoms and positi"e and neati"e hydroen ions! a completely unphysical result. This eplains in part why the cur"e of total enery aainst interatomic distance for the "alence bond method lies below the cur"e for the molecular orbital method at all distances and most particularly so for lare distances. This situation arises for all homonuclear diatomic molecules and is particularly a problem for B6! where the minimum enery of the cur"e with molecular orbital theory is still hiher in enery than the enery of two B atoms.

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The concepts of hybridization are so "ersatile! and the "ariability in bondin in most oranic compounds is so modest! that "alence bond theory remains an interal part of the "ocabulary of oranic chemistry. ,owe"er! the wor% of Briedrich ,und! 4obert Mulli%en! and Ferhard ,erzber showed that molecular orbital theory pro"ided a more appropriate description of the spectroscopic! ionization and manetic properties of molecules. The deficiencies of "alence bond theory became apparent when hyper"alent molecules *e.. 3BJ+ were eplained without the use of d orbitals that were crucial to the bondin hybridisation scheme proposed for such molecules by 3aulin. Metal complees and electron deficient compounds *e.. diborane+ also appeared to be well described by molecular orbital theory! althouh "alence bond descriptions ha"e been made.

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In the 5> 95<  C&C 5J9