CHEMICAL BONDING.pptx

(1)

CHEMICAL

BONDING

Diadem L. Cruz

III- St. Matthew

(2)

(3)

Introduction



_{A c}_{A c}_{hemical bo}_{hemical bo}_{nd is an attraction be}_{nd is an attraction be}_{tween atoms tha}_{tween atoms tha}_{t allows the formati}_{t allows the formati}_{on of chemical s}_{on of chemical s}_{ubstances that}_{ubstances that}

contain two or more atoms. The bond is caused by the electrostatic force of

contain two or more atoms. The bond is caused by the electrostatic force of attraction between opposattraction between oppositeite chares! either between electrons and nucle

chares! either between electrons and nuclei! or as i! or as the result of a the result of a dipole attraction. The strenth ofdipole attraction. The strenth of chemical bonds "aries considerably# there are $stron bonds$ such as co"alent or ionic bonds and chemical bonds "aries considerably# there are $stron bonds$ such as co"alent or ionic bonds and $wea% bonds$ such as

$wea% bonds$ such as dipole&ddipole&dipole interactions! the London dispersion force ipole interactions! the London dispersion force and hydroen bondin.and hydroen bondin.



_{Since opposite chares attract "ia a}_{Since opposite chares attract "ia a}_{simple electroma}_{simple electroma}_{netic force! the}_{netic force! the}_neati"el_neati"el_{y chared electrons that}_{y chared electrons that}

are orbitin the nucleus and the positi"ely chared protons in the nucleus attract each other. An 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 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

toward electrons in this position. This attraction constitutes the chemical bond. Due to the chemical bond. Due to the matter wa"ematter wa"e nature of electrons and their smaller mass! they must occupy a much larer amount of "olume

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 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 far apart! as compared with the size of the nuclei themsel"es. This phenomenon limits the distance between nuclei and atoms in a bond.

between nuclei and atoms in a bond.



_{In eneral! stron chemical bondin is associated wi}_{In eneral! stron chemical bondin is associated wi}_{th the sharin or tr}_{th the sharin or tr}_{ansfer of electrons between the}_{ansfer of electrons between the}

participatin atoms. The atoms in molecules! crystals! metals and diatomic ases' indeed most of participatin atoms. The atoms in molecules! crystals! metals and diatomic ases' indeed most of thethe physical en"ironment around us' are held toether by chemical bonds! which dictate the structure and physical en"ironment around us' are held toether by chemical bonds! which dictate the structure and the bul% properties of matter.

(4)

(5)

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.

_{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.

(6)



_{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.

(7)

_{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.

_{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.

(8)

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 $;uery <5$ of his 0ptic%s! whereby atoms attach to each other by some $force$. Specifically! after ac%nowledin the "arious popular theories in "oue at the time! of how atoms were reasoned to attach to each other! i.e. $hoo%ed atoms$! $lued toether by rest$! or $stuc% toether by conspirin motions$! :ewton states that he would rather infer from their cohesion! that $particles attract one another by some force! which in immediate contact is eceedinly stron! at small distances performs the chemical operations! and reaches not far from the particles with any sensible effect.$

_{In 5=5>! 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! $An electron may form a part of the shell of two different atoms and cannot be said to belon to either one eclusi"ely.$

(9)

_{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+.

_{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.

(10)



_{In 5><<! ,. ,. ?ames and A. S. Coolide carried out a calculation}

on the dihydroen molecule that! unli%e all pre"ious calculation

which used functions only of the distance of the electron from the

atomic nucleus! used functions which also eplicitly added the

distance between the two electrons.H< /ith up to 5< adustable

parameters they obtained a result "ery close to the eperimental

result for the dissociation enery. Later etensions ha"e used up

to J9 parameters and i"e ecellent areement with eperiment.

This calculation con"inced the scientific community that (uantum

theory could i"e areement with eperiment. ,owe"er this

approach has none of the physical pictures of the "alence bond

and molecular orbital theories and is difficult to etend to larer

molecules.

(11)

&aence !ond theory



_{In 5>67! "alence bond theory was formulated and it arues that a chemical}

bond forms when two "alence electrons! in their respecti"e atomic orbitals!

wor% or function to hold two nuclei toether! by "irtue of effects of lowerin

system eneries. uildin on this theory! the chemist Linus 3aulin published

in 5><5 what some consider one of the most important papers in the history of

chemistryK $0n the :ature of the Chemical ond$. In this paper! elaboratin on

the wor%s of Lewis! and the "alence bond theory *1+ of ,eitler and London!

and his own earlier wor%s! 3aulin presented si rules for the shared electron

bond! the first three of which were already enerally %nownK



_{5. The electron-pair bond forms throuh the interaction of an unpaired electron}

on each of two atoms.



_{6. The spins of the electrons ha"e to be opposed.}

(12)



_{,is last three rules were newK}



_{9. The electron-echane terms for the bond in"ol"es only one wa"e function from each}

atom.



_{J. The a"ailable electrons in the lowest enery le"el form the stronest bonds.}



_{E. 0f two orbitals in an atom! the one that can o"erlap the most with an orbital from another}

atom will form the stronest bond! and this bond will tend to lie in the direction of the

concentrated orbital.



_{uildin on this article! 3aulin)s 5><> tetboo%K 0n the :ature of the Chemical ond}

would become what some ha"e called the $ible$ of modern chemistry. This boo% helped

eperimental chemists to understand the impact of (uantum theory on chemistry. ,owe"er!

the later edition 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.

(13)

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.

(14)



_{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.

(15)



_{In the 5><8s the two methods stronly competed until it was realised}

that they are both approimations to a better theory. If we ta%e the

simple "alence bond structure and mi in all possible co"alent and ionic

structures arisin from a particular set of atomic orbitals! we reach what

is called the full confiuration interaction wa"e function. If we ta%e the

simple molecular orbital description of the round state and combine

that function with the functions describin all possible ecited states

usin unoccupied orbitals arisin from the same set of atomic orbitals!

we also reach the full confiuration interaction wa"efunction. It can be

then seen that the simple molecular orbital approach i"es too much

weiht to the ionic structures! while the simple "alence bond approach

i"es too little. This can also be described as sayin that the molecular

orbital approach is too delocalised! while the "alence bond approach is

too localised.

(16)



_{The two approaches are now rearded as}

complementary! each pro"idin its own insihts into

the problem of chemical bondin. Modern

calculations in (uantum chemistry usually start from

*but ultimately o far beyond+ a molecular orbital

rather than a "alence bond approach! not because

of any intrinsic superiority in the former but rather

because the M0 approach is more readily adapted

to numerical computations. ,owe"er better "alence

bond prorams are now a"ailable.

(17)

Bonds in chemica

formuas



_{The fact that atoms and molecules are three-dimensional ma%es it difficult to}

use a sinle techni(ue for indicatin orbitals and bonds. In molecular formulas

the chemical bonds *bindin orbitals+ between atoms are indicated by "arious

methods accordin to the type of discussion. Sometimes! they are completely

nelected. Bor eample! in oranic chemistry chemists are sometimes

concerned only with the functional roups of the molecule. Thus! the molecular

formula of ethanol may be written in a paper in conformational!

dimensional! full two-dimensional *indicatin e"ery bond with no

three-dimensional directions+! compressed two-three-dimensional *C,<&C,6&0,+!

separatin the functional roup from another part of the molecule *C6,J0,+!

or by its atomic constituents *C6,E0+! accordin to what is discussed.

Sometimes! e"en the non-bondin "alence shell electrons *with the

two-dimensional approimate directions+ are mar%ed! i.e. for elemental carbon .)C).

Some chemists may also mar% the respecti"e orbitals! i.e. the hypothetical

ethene9 anion *NCOCN 9+ indicatin the possibility of bond formation.

(18)

"tron' chemica !onds



_{Stron chemical bonds are the intramolecular forces which hold atoms}

toether in molecules. A stron chemical bond is formed from the transfer or

sharin of electrons between atomic centers and relies on the electrostatic

attraction between the protons in nuclei and the electrons in the orbitals.

Althouh these bonds typically in"ol"e the transfer of inteer numbers of

electrons *this is the bond order! which represents one transferred electron

or two shared electrons+! some systems can ha"e intermediate numbers of

bonds. An eample of this is the oranic molecule benzene! where the bond

order is 5.J for each carbon atom! meanin that it has 5.J bonds *shares

three electrons+ with each one of its two neihbors.



_{The types of stron bond differ due to the difference in electroneati"ity of}

the constituent elements. A lare difference in electroneati"ity leads to

more polar *ionic+ character in the bond.

(19)

Ionic !ond



_{Ionic bondin is a type of electrostatic interaction between atoms which ha"e a}

lare electroneati"ity difference. There is no precise "alue that distinuishes

ionic from co"alent bondin! but a difference of electroneati"ity of o"er 5.7 is

li%ely to be ionic! and a difference of less than 5.7 is li%ely to be co"alent.HJ Ionic

bondin leads to separate positi"e and neati"e ions. Ionic chares are

commonly between <e to G<e. Ionic bondin commonly occurs in metal salts

such as sodium chloride *table salt+. A typical feature of ionic bonds is that the

species form into ionic crystals! in which no ion is specifically paired with any

sinle other ion! in a specific directional bond. 4ather! each species of ion is

surrounded by ions of the opposite chare! and the spacin between it and each

of the oppositely chared ions near it! is the same for all surroundin atoms of

the same type. It is thus no loner possible to associate an ion with any specific

other sinle ionized atom near it. This is a situation unli%e that in co"alent

crystals! where co"alent bonds between specific atoms are still discernible from

the shorter distances between them! as measured "ia such techni(ues as P-ray

diffraction.

(20)



_{Ionic crystals may contain a miture of co"alent and ionic}

species! as for eample salts of comple acids! such as

sodium cyanide! :aC:. Many minerals are of this type.

P-ray diffraction shows that in :aC:! for eample! the bonds

between sodium cations *:aG+ and the cyanide anions

*C:-+ are ionic! with no sodium ion associated with any

particular cyanide. ,owe"er! the bonds between C and :

atoms in cyanide are of the co"alent type! ma%in each of

the carbon and nitroen associated with ust one of its

opposite type! to which it is physically much closer than it is

to other carbons or nitroens in a sodium cyanide crystal.

(21)



_{/hen such crystals are melted into li(uids! the ionic bonds are}

bro%en first because they are non-directional and allow the

chared species to mo"e freely. Similarly! when such salts dissol"e

into water! the ionic bonds are typically bro%en by the interaction

with water! but the co"alent bonds continue to hold. Bor eample!

in solution! the cyanide ions! still bound toether as sinle

C:-ions! mo"e independently throuh the solution! as do sodium C:-ions!

as :aG. In water! chared ions mo"e apart because each of them

are more stronly attracted to a number of water molecules! than

to each other. The attraction between ions and water molecules in

such solutions is due to a type of wea% dipole-dipole type chemical

bond. In melted ionic compounds! the ions continue to be attracted

to each other! but not in any ordered or crystalline way.

(22)

#ypica !ond en'ths in pm

and !ond ener'ies in ()*mo+

_{ond Lenth} _{*pm+ 2n er y} _*%?Nmol+ _{, ' ,ydroen} _,&, _{79 9<E} _{,&0 >E <EE} _,&B _{>6 JE=} _{,&Cl 567} _9<6 _{C ' Carbon} _C&, _58> _95< _C&C _5J9 _<9= _{C&CO 5J5} _OC&CQ ₅₉₇ _OC&CO ₅₉₌ _{COC 5<9} _E59 _{CQC 568} _=<> _C&: ₅₉₇ _<8= _{C&0 59<} _<E8 _C&B _5<9 ₉₌₌ _{C&Cl 577} _<<8 

_{: ' :itroen}



_:&,

₅₈₅

_<>5



_:&:

_59J

₅₇₈



_:Q:

₅₅₈

_>9J



_{0 ' 0yen}



_{0&0 59=}

_59J



_{0O0 565}

_9>=



_{B! Cl! r! I ' ,aloens}



_B&B

₅₉₆

_5J=



_{Cl&Cl 5>>}

_69<



_{r&, 595}

_<EE



_{r&r 66=}

_5><



_I&,

_5E5

_6>=



_{I&I 6E7}

_5J5

(23)

Covaent !ond

_{Co"alent bondin is a common type of bondin! in which the electroneati"ity difference}

between the bonded atoms is small or noneistent. onds within most oranic compounds are described as co"alent. See sima bonds and pi bonds for LCA0-description of such bondin.

_{A polar co"alent bond is a co"alent bond with a sinificant ionic character. This means that the}

electrons are closer to one o f the atoms than the other! creatin an imbalance of chare. They occur as a bond between two atoms with moderately different electroneati"ities! and i"e rise to dipole-dipole interactions. The electroneati"ity of these bonds is 8.< to 5.7 .

_{A coordinate co"alent bond is one where both bondin electrons are from one of the atoms}

in"ol"ed in the bond. These bonds i"e rise to Lewis acids and bases. The electrons are shared rouhly e(ually between the atoms in contrast to ionic bondin. Such bond in occurs in

molecules such as the ammonium ion *:,9G+ and are shown by an arrow pointin to the Lewis acid. Also %nown as non-polar co"alent bond! the electroneati"ity of these bonds rane from 8 to 8.<.

_{Molecules which are formed primarily from non-polar co"alent bonds are often immiscible in}

(24)

(25)



_{0ne-electron bondin in the dihydroen cation.}



_{onds with one or three electrons can be found in radical}

species! which ha"e an odd number of electrons. The simplest

eample of a 5-electron bond is found in the dihydroen cation!

,6G. 0ne-electron bonds often ha"e about half the bond

enery of a 6-electron bond! and are therefore called $half

bonds$. ,owe"er! there are eceptionsK in the case of dilithium!

the bond is actually stroner for the 5-electron Li6G than for the

6-electron Li6. This eception can be eplained in terms of

hybridization and inner-shell effects.HE



_{Comparison of the electronic structure of the three-electron}

(26)

(27)



_{The simplest eample of three-electron bondin can be found in the helium}

dimer cation! ,e6G. It is considered a $half bond$ because it consists of only

one shared electron *rather than two+ in addition to one lone electron on

each atom# in molecular orbital terms! the third electron is in an anti-bondin

orbital which cancels out half of the bond formed by the other two electrons.

Another eample of a molecule containin a <-electron bond! in addition to

two 6-electron bonds! is nitric oide! :0. The oyen molecule! 06 can also

be rearded as ha"in two <-electron bonds and one 6-electron bond! which

accounts for its paramanetism and its formal bond order of 6.H7 Chlorine

dioide and its hea"ier analoues bromine dioide and iodine dioide also

contain three-electron bonds.



_{Molecules with odd-electron bonds are usually hihly reacti"e. These types}

(28)

Bent Bonds



_{ent bonds! also %nown as banana bonds!}

are bonds in strained or otherwise sterically

hindered molecules whose bindin orbitals

are forced into a banana-li%e form. ent

bonds are often more susceptible to reactions

than ordinary bonds.

(29)

$esonant !ondin'

(30)

Hypervaent !ondin'



_{In hyper"alent molecules! there eists bonds}

which ha"e sinificant non-bondin ionic

(uality to them. This manifests as

non-bondin orbital le"els in molecular orbital

theory! while in "alence bond theory it is

analyzed as a form of resonant bondin.

(31)

Eectron,de-cient !ondin'



_{In three-center two-electron bonds *$<c&6e$+ three atoms share two}

electrons in bondin. This type of bondin occurs in electron deficient

compounds li%e diborane. 2ach such bond *6 per molecule in diborane+

contains a pair of electrons which connect the boron atoms to each other

in a banana shape! with a proton *nucleus of a hydroen atom+ in the

middle of the bond! sharin electrons with both boron atoms. In certain

cluster compounds! so-called four-center two-electron bonds also ha"e

been postulated.



_{In certain conuated R *pi+ systems! such as benzene and other}

aromatic compounds *see below+! and in conuated networ% solids such

as raphite! the electrons in the conuated system of R-bonds are

spread o"er as many nuclear centers as eist in the molecule! or the

networ%.

(32)

Aromatic !ondin'



_{In oranic chemistry! certain confiurations of electrons and orbitals infer etra stability}

to a molecule. This occurs when R orbitals o"erlap and combine with others on different

atomic centres! formin a lon rane bond. Bor a molecule to be aromatic! it must obey

,c%el)s rule! where the number of R electrons fit the formula 9n G 6! where n is an

inteer. The bonds in"ol"ed in the aromaticity are all planar.



_{In benzene! the prototypical aromatic compound! 5= *n O 9+ bondin electrons bind E}

carbon atoms toether to form a planar rin structure. The bond $order$ *a"erae

number of bonds+ between the different carbon atoms may be said to be *5=NE+N6O5.J!

but in this case the bonds are all identical from the chemical point of "iew. They may

sometimes be written as sinle bonds alternatin with double bonds! but the "iew of all

rin bonds as bein e(ui"alently about 5.J bonds in strenth! is much closer to truth.



_{In the case of heterocyclic aromatics and substituted benzenes! the electroneati"ity}

differences between different parts of the rin may dominate the chemical beha"iour of

aromatic rin bonds! which otherwise are e(ui"alent.

(33)

Metaic !ond



_{In a metallic bond! bondin electrons are}

delocalized o"er a lattice of atoms. y

contrast! in ionic compounds! the locations of

the bindin electrons and their chares are

static. The freely-mo"in or delocalization of

bondin electrons leads to classical metallic

properties such as luster *surface liht

reflecti"ity+! electrical and thermal conducti"ity!

ductility! and hih tensile strenth.

(34)

Intermoecuar !ondin'

There are four basic types of bonds that can be formed between two or more *otherwise

non-associated+ molecules! ions or atoms. Intermolecular forces cause molecules to be

attracted or repulsed by each other. 0ften! these define some of the physical

characteristics *such as the meltin point+ of a substance.



_{A lare difference in electroneati"ity between two bonded atoms will cause a}

permanent chare separation! or dipole! in a molecule or ion. Two or more molecules or

ions with permanent dipoles can interact within dipole-dipole interactions. The bondin

electrons in a molecule or ion will! on a"erae! be closer to the more electroneati"e

atom more fre(uently than the less electroneati"e one! i"in rise to partial chares on

each atom! and causin electrostatic forces between molecules or ions.



_{A hydroen bond is effecti"ely a stron eample of an interaction between two}

permanent dipoles. The lare difference in electroneati"ities between hydroen and

any of fluorine! nitroen and oyen! coupled with their lone pairs of electrons cause

stron electrostatic forces between molecules. ,ydroen bonds are responsible for the

hih boilin points of water and ammonia with respect to their hea"ier analoues.

(35)



_{The London dispersion force arises due to}

instantaneous dipoles in neihbourin atoms. As the

neati"e chare of the electron is not uniform around

the whole atom! there is always a chare imbalance.

This small chare will induce a correspondin dipole

in a nearby molecule# causin an attraction between

the two. The electron then mo"es to another part of

the electron cloud and the attraction is bro%en.



_{A cation&pi interaction occurs between a pi bond and}

(36)

"ummary. eectrons in

chemica !onds

_{In the *unrealistic+ limit of $pure$ ionic bondin! electrons a re perfectly localized on one of the two}

atoms in the bond. Such bonds can be understood by classical physics. The forces between the atoms are characterized by isotropic continuum electrostatic potentials. Their manitude is in simple proportion to the chare difference.

_{Co"alent bonds are better understood by "alence bond theory or molecular orbital theory. The}

properties of the atoms in"ol"ed can be understood usin concepts such as oidation number. The electron density within a bond is not assined to indi"idual atoms! but is instead delocalized between atoms. In "alence bond theory! the two electrons on the two atoms a re coupled toether with the bond strenth dependin on the o"erlap between them. In molecular orbital theory! the linear combination of atomic orbitals *LCA0+ helps describe the delocalized molecular o rbital structures and eneries based on the atomic orbitals of the atoms they came from. nli%e pure ionic bonds! co"alent bonds may ha"e directed anisotropic properties. These may ha"e their own names! such as sima bond and pi bond.

_{In the eneral case! atoms form bon ds that are intermediates between ionic and co"alent!}

dependin on the relati"e electroneati"ity of the atoms in"ol"ed. This type of bond is sometimes called polar co"alent.