t, - Anderson Junior GollegeJC 1 Chemistry Lecture

mfm/bonding/2004

Lecture Outline:1 lonic (electrovalent) bond2 Metallic bond3 (l) Covalent bond

(lt) Co-ordinate (dative covalent) bond(lll) Bond energy, bond length and bond polarity

4 lonic bond with covalent character5 Simple molecules(l) Polar & non-Polar molecules

(ll) Intermolecular forces: van der Waals' forces & hydrogen bonding6 Shapes of simple molecules: VSEPR approach7 Atomic orbital approach to shapes of molecules

References:1 Michael Freemantle. Chemistry in Action-2 Brown, LeMay, Bursten. Chemistry, The Central Science.3 E. N. Ramsden. A-Level Chemistry.4 Chris Conoley & Phil Hills.Chernlsfry.5 JGR Briggs. Longrnan AJevel Course in Chemistry

Assessment Objectives:(a) describe ionic (electrovalent) bonding, as in sodium chloride and magnesium oxide,including the use of'dot-ad cross' diagrams(b) describe, including the use of'dotad cross' diagrams,(i) covalent bonding, as in hydrogen; oxygen; chlorine; hydrogen chloride; carbondioxide; methane; ethene(ii) co-ordinate (dativecovalent) bonding, as in BFe.NHa(c) explain the shapes of, and bond angles in, molecules by using the qualitative model ofelectron-pair repulsion (including lone pairs), using as simple examples: BF3 (trigonal)' CO2(linear); CHl (tetrahedral); NHa (pyramidal)i H2O (nonlinear); SF6 (oclahedral)(d) describe covalentbonding in terms of orbital overlap, giving and bonds(e) explain the shape of, and bond angles in, the ethane, ethane and benzene molecules interms of and bonds.(f) predict the shapes of, and bond angles in, molecules analogous to those specmed in (c)and (e)(g) describe hydrogen bonding, using ammonia and water as simple examples of moleculescontaining -NH and -OH groups(h) explain the terms bond energy, bond length and bond polarity and use them to comparethe reactivities of covalent bonds(i) describe intermolecular forces (van der Waals' forces) based on permanent and induceddipoles, as in CHCI3(|); 8rz(l) and liquid noble gasesO describe metallic bonding in terms of a lattice of positive ions sunounded by mobileelectrons(k)describe, interpret and / predict the effect of different types of bonding (ionic bonding;covalent bonding; hydrogen bonding; other intermolecular attractions; metallic bonding) onthe physical properties of substances(l) deduce the type of bonding present from given information(m) show understanding of chemical reactions in terms of energy transfers associated withbreaking and making of chemical bonds.

1 lntroduction

. A chemical bond is a force that holds together two or more atoms, ions,molecules or any combination of these.

. The type of chemical bond present in a substance largely determines theproperties of the substance.

. Generally, there are two main types of bonds based on the strength of theforce of attraction.

(a) Stronq BondinqThere are 3 main types of strong bonds:

(i) Metallic bonding(ii) lonic(electrovalent)bonding(iii) Covalent bonding (including co-ordinate/dative bonding)

. The type of bond in a substance is primarily determined by theouter-shell electronic structure of its atoms.

. In forming bonds, atoms either gain, lose, or share electrons toachieve the lowest energy and thus most stable electronicconfiguration.

. The stable configuration is the configuration of a noble gas closestto the element in the periodic table.

. This guideline is called the octet rule.

(b) Weak Bonds: Intermolecular Forces of Attraction(Bonding between molecules). Some atoms are bonded together by covalent bonds to form small,

discrete molecules.

. Weak forces of attraction exist between these molecules. Theseintermolecular forces include

(i) Permanent dipole - permanent dipole aftraction I Also called van der(ii) Instantaneous dipole - induced dipole attraction J Waals' forces(iii) Hydrogen bonding

2 Metallic Bond. This is the strong electrostatic attraction between the lattice of metal

cations and the 'sea' of delocalised electrons present in a metalelement.

. The electrons involved in metallic bonding are the outer-shell or valenceelectrons of the metal atoms.

. These electrons are relatively easily removed from the atoms resulting inthe formation of metal cations.

. The electrons are delocalised, not controlled by specific metal nuclei. Theelectrons move randomly throughout the lattice of regularly ananged metalcations.

. Each metal cation is attracted to the delocalised electrons and vice versa.

. The strength of the metallic bond is reflected in the high melting andboiling points of most metals-

Eg. metal m.p.fo b.o.fCK 63.4 757Ca 851 1482Cu 1083 2582

Question 1:

.a a"hin t +. Vt.* 4h ,tif,lr,llt, lo.)r

Whot do you think contributes to the increose in the melting points(ond boiling points) of the three metols?

-

!.calt :h iu-L,v t l

J ,o- | y rhf,lo6J,a ettc|.o'.s w

- e{4 4i}6 h r^ ('!A}n

,tr}/ ' \^. fr

el?. i 'd 'rJ ta^. i 'y irvt) SO2 tb (o to

-,a +r" ! l r^ (n,

*t .r1^lu. \ ." l -r f- ,4rf" de 'h.t . l l ; r L. l .^r,

u,;l.,r' A lL,' ^4ttg

Properties associated with metallic bonds (\v p_ r oe ra rr ,1 Generally high melting points and boiling points.

- ge,{"i9 i"..r * ? .

-

-.tta+ra1g ,{t,?L N+nltt' |.o'.dt

2 Metals are malleable and ductile +|'o'' o''r'.y {arau'-

8-r6r

. This bond is the strong electrostatic forces of attraction between ionsof opposite charges.

. The ions may be formed from atoms as a result of electron transfer fromone atom to the other.

. Typically, binary (2 elements) ionic compounds consist of metallic cationsfrom Group I and ll and non-metallic anions from Group Vl and Vll.

Example of formation of ions by electron transferNa+

1 s22s22p63s1-) Na* +

1 s22s22p63s23p5 1s22s22p6

@'-+ 9{'.o.5 ,t.t,rnr{, a+lyrr-

oafog qrl ei,

1s22s22p63s23p0

7 c+,t tnDot-and

-cross diagramf , . l t f

^ . ' l -1;r . ta! l l l r r : it . . J L {( J

Other examples of ionic compounds' l MgO r- r2rr

I . I I I ^ .xl :un, l l : ol , "J. l l '^ ' "

\ r . J \xl

f .:", J Fi;J'.f :;;JCaFzr - ' t t a

-?-I I | { r Il ' l r I l .^^ ll . - , I I o v . II r t . .J

LizO f',, J'NOTE:1 There is no discrete molecules in ionic compounds. The formula of an

ionic compound is an empirical formula for the infinite number of cationsand anions in the compound. Thus, it is incorrect to write "molecule ofNaCl".

2 There are other examples of ionic compounds which are not formed bydirect electron transfer.Example : NaNOg - ionic bonding between Na* and NO3-(a polyatomic ion)

. The strength of ionic bond depends mainly on the magnitude of chargedensity on the cation and on the anion.,1 4I Charge density on ion = amount of charoe | ,:

lonic radius J

. Strength of electrostatic attraction between ions of high charge densitiesdf tw t-alt, .ha-t lrl[ ,t4a clo.y &a!\ ,t Na'

"

ls,,*q ao^ ,3t'

. The laftice energy of an ionic compound is a measure of the strength ofthe ionic bond in the compound. 9v , crcy &'+;ti "r t' ' 1"5 4r'^' Nd '

Y,.;};j'j";*,,:1.' ;' r+@3'v

) Xoat Lo.r u r : . t . t io,q, +l .n lv\ ^ ' ' i r(Definition: Lattice energy is the energy released when one mole of an ionic solid isformed from its gaseous ions. - will be studied in greater detail in later lectures)

(For your info only) Lattice eneroies of some ionic solidsCompound Lattice energy(kJ/mol)

Compound Lattice energy(kJ/mol)LiF 1030 MqClz 2326Licl 834 SrCl, 2127Lil 730NaF 9'10 Mso 3795NaCl 788 CaO 3414Nal 682 SrO 3217

Phvsical Properties of lonic Compounds

. High melting point and high boiling point.

. ,.'1,r Soluble in water and other polar solvents.

. Good electrical conductors when in molten state or when dissolved asions in aqueous solutions.

. lonic solids are brittle. it fractures under stress.(Refer to Solid Structure lecture for explanation of thse properties)

is stronger.

EXample Lit I 'r ' , 'h^rtrv ; 'r ' ' 'abvt t1"cr r ' o-

+ cW, &r!r 5 "1 ri*

. I ar ,f v .ll"o' t'16'

The relative strength of ionic bond in compounds is also reflected in therelative melting point or boiling point of the *tpounOt. t J'O-;

';. '""^d + ,l

'' 1 t b p

Question 2: *, ..' At' *'The oxides NozO ond AlzOg hove the melting points 1275 'C ond 2O7? "Crespectively.Whot do you think qccount for the difference in the melting points?-lf t 't ,nu,r, \ '1hv cl'vr4r &xl3 Mar' 1u+-

Ekclrrt ' i ( " t+vr.. t ' o.. 4a

Alv ", t o";t ,r4h r+-eqtu '1lzr

t /w Nat arl c ' '- G.?eer ar*.f ol e\Ntt\l 4qttlfll 'fo h.l+ A?) 0a tlen da, d

+ 1t1 Covalent Bond. Atoms can also achieve stable outer electronic configurations (noble gas

confi guration) by sharing electrons.. A covalent bond consists of a pair of electrons shared between 2

adjacent atoms. Each atom contributes one electron.. The strong electrostatic forces of attraction between the two

positively charged nuclei and the negatively charged, shared pair ofelectrons between them constitutes a covalent bond which holds the twoatoms together in a molecule or in a polyatomic ion.

. A covalent bond tends to form between atoms that do not easily lose orgain electrons from the other; i.e. between atoms of similar electron-attracting ability (similar electronegativity).

Dot-and-cross diagrams and Lewis diagrams/structures in representingcovalent molecules

r h l1.,ts 4.,/.t\d , uo^ ),'g Pa'rlelkl a,, r!

'{ P,r-r..tr} bl ,

-

t

6 ."1 bw F o,y ! .4

, . . ' / ' rx ' ,

Molecule Dot-and-cross diagram Lewis structure

H2 t-\ i l-r g *H

Oz o o =O

HCI Hl c l ; H - CI I

Coz! t tc: lo 0 = c=0

4 (ll) Solid state structure of covalent substancesCovalent substances in solid state are molecular. There are two types ofsolid state structures:

o Simple molecular structure1 eg: iodine crystals, lz(s); solid carbon dioxide, CO2(s); ice HzO(s)&-*.---

4 (lll) Goordinate (dative) bond. This is a type of covalent bond in which both the shared electron pair

comes from one atom.x!Y

. lt has the same strength as a covalent bond.

. lt usually occurs when atoms in molecules have pairs of electrons in theirouter shells that are not involved in covalent bonds. These non-bondingelectrons are called lone-pairs. Sometimes these lone-pairs are used toform a rcrehrnt bond to an atom that can accommodate two moreelectrons. a'ttt'

'Dono/ atom : donates the shared pair of electrons,: has at least 1 lone pair of electrons available in thevalence shell

Acceptor' : accepts and shares the shared pair of electrons, * "".t: has at least 1 vacant orbital in its valence shell,

usually a metal ion, transition metal or an atom in u ,*j,- ).,ru, r

molecule. Dative bond is represented by an anow, * , from the'donor'atom to

the 'acceptor' atom.

Examples1 In polyatomic ions

HtO' (hvdroxonium ion): formed when a lone pair of electrons on Oatom of H2O forms a dative bond with a proton, H'.

.) a'q'

+ H't f*H

L,ar. t ! i . f r i ,r l . .+

t t t l " "H

2 ln complex ions

--+ [";'"i 1' " [,vi^)'NHg* (ammonium ion): formed when a lone pair of electrons on N atomof NHg forms a dative bond with a proton, H'.

r { r {| . . I--) f ..-'ry'-. I(" ; "J

Coordinate bonds are responsible for the formation of complex ionssuch as [Cu(HzO)e]2'and [Fe(CN)6]3- in which each H2O or CN-usesone of its lone pair to form a coordinate bond to the central positivemetal ion'

l cN cN'Es. Fe3* + 6 !cN- -+ Fe(cN).3- I ,,f_!; ^-i

cN'

I '---tar.ttNC - - l _\I ICN.I

r CN-n

4 (lV) Covalent bondinq in polvatomic ions. In polyatomic ions, two or more atoms are bound together predominantly

by covalent bonds.

. These atoms form a stable grouping that canies a charge, either positiveor negative.

. lonic bond binds these ions to other oppositely charged ions'

Eg. CO:2- NOr-

006t

Eg. The types of bonds present in solid NHa-Cl- are:

covalent (normal and coordinate) bonds in NHat ion andionic bond between NHa' ions and Cl- ions.

(l) The Octet RuleMany atoms undergoing reactions often gain, lose or share electrons toachieve eight valence electrons (noble gas configuration, except He).

An octet of electrons consists of full s and p subshells on an atom (i.e.,four pairs of valence electrons). s'?bMost representative (s- and p- block) elements tend to achieve this stableconfiguration in most of their compounds.

This rule is simple and useful in understanding the basic concepts ofbonding. However, it has its limitation in dealing with ionic compounds ofthe transition metals (as seen in the iron in Fe(CN)6" mentioned above). ltalso fails in many situations involving covalent bonding, as seen on thenext page.

r , il r - l ll + |lNll ' r \ |(f' rr ut7

dHr+

( i )( i i )

(i)(il)Molecules with an odd number (or unpaired) electrons.In most molecules, the number of electrons is even, and completepairing of electrons occurs.

However, in a few molecules such as ClO2, NO, NOz, the number ofelectrons is odd. Complete pairing of these electrons is impossible, andan octet around each atom cannot be achieved.

Particles with unpaired or odd electrons are called ftee radicals. Freeradicals are very reactive substances.

Examoles

Nitrogen monoxide, NO

' * (o"r

. '1, " J: (t ') )

.+

. , , r (9a)

r 'N !.1 o ;. .

xx 1r e).N.. Or

Nitrogen dioxide, NOz) t ' " , ! . ( l ' \ . ' . , "

-( ' t t \'..

tg - ,tni ,o;

+

| /-dr ttnr6rr ri,, 5., ,o.tI artiv. &, ' '

["* *''* ' Iru^*owf L n-" )

fr+L ao.fc

8c' t " . g,0

-tt ctt,+aag4;,1j

I l^l,q (t)

)I dr.t (L)

0z

* 0c ' N- o,8trt ,ntg a?- 4, A)

&- LxJ

lv7

[ .aa,g o al

*J", ! t klbr,'

( i i ) Molecules in which the central atom has less than an octet.1 A I so cal bd EGEiro-nTeRGilmo lec u les )Often encountered in covalent compounds containing beryllium(Be),boron (B) and aluminium (Al) as central atoms. (o$ rn 6,7I )

,lv g L \1.BeCl, (S), BFg and AlClg (g) are typical examples.Be has 4 electrons in its valence shell in BeClz

B and Al atoms are each sunounded by 6 electrons in BF3 andAlCl3 respectively.

f t Y(

iClr tBe,xCl ixt F,

(r)

Yl- {axR

xY x ix

FI F (t \ ' a

, " ;

AInt at ' * r ta\t ,_rlft ..-'.'

( 6c)

These molecules tend to react with other molecules or ions so thatthe central atom can achieve an octet of electrons. Comooundsformed in this way are called addition compounds.

Examples(a) Reaction between ammonia and boron trifluoride

toc)

iNH3 + BF: -+ HgNiBF3 or H3N---+BF3u1'^

-

*'-tt v l6at BYqt '

lltla

u, ' f r .d t tV.+T t

Question 4Suggest the structurol formuloe of the products formed inthe reoction belween

Beryllium dichloride, BeClz ond ommonioi: r,l- \

' ^'r{t'nrs -) le ' ' t fsr a. er l ! S*rr ,

atBoron trifluoride, BFg ond fluoride ion, F-

sF, . [ +i] -+ eF.1.. E' l

l!"'

(b) Dimerisation () xorcd'r 6tvrt u? 3 fu* w* *Ydt)r-.t----/f | . . r ' i rntv

Electron deficient molecules such as aluminium chloride tend todimerise (two molecules bonded to form 1) so that each central Al atomcan achieve an octet of electrons.

h'.n Y

2 AlCls + AlzCle

Structure of AleClo dimerJ. 2

',1.L 1.\.r L'"&J 1"4 f Jrtw

lrdt { '$o" a lin tv '

AI. / \

Molecules in which the central atom has more than an octet.

. Central atoms in these molecules are from Period 3 or greater.

. Typical examples include PFs, SFo, and XeFr.

cl

cl

\AI

cl

cl

cl

cl

(i i i)

FF^. l .^F

F*. . ,F

toc -

P

FF .r .

-

used in bonding.eight electrons in

Hence. the atom can accommodate more thanits valence shell.

tr3,J

t l t141"P c)

a Ce.r ark-qd 'a16'Yav' e/ ll e I

r*^revt I "r"al! '-q r" 'r :a.

(a)

Size of the atom partly determines whether the atom canaccommodate more than eight electrons.

The larger the central atom, the larger the number of atoms thatCan SUffound il-

- ,o, o.co*oloQ iot gtlr'c:t 4lar ; 't'alb

alsvr '

The size of the sunounding atoms is also important. Expandedvalence shells occur most often when the central atom is bonded tothe smallest and most electronegative atoms, such as F' Cl, and O.

Bond Enerqv (Bond EnthalPv)Strength of a covalent bond is measured by the bond energy.

Bond energy is the amount of energy required to break one moleof covalent bond in one mole of a gaseous covalent molecule (orthe amount of energy evolved in forming one mole of covalentbond).Example:

Bond energy for Cl-Cl is 242 kJ mol-1

ch(s)+2cl(g)2 cl(g) + Cl2 (s)

( rUlo,U r )Bond (dissociation) energy = + 242 kJ mol-'Bond (formation) energy ='242kJ mol-1

tc! l"tld)

. The stronger the covalent bond, the greater is the bond energy.

N2 molecule has a very large N -N bond energy of 994 kJ mol-l ' Pequi''- u(r cth,lt

'

44+ tJ ,cl ?"r'al +.

^1, {l)

-> J N (5) r,.F I kr { .HdJh

Covalent molecules with higher bond energies generally have lower fl;i'.' ,:.T:,'tendency to undergo decomposition . t/.,\^ t b,.a.lHte

l3

HCI (g) does not decomposeHl (g) decomposes easily atred heat.

Average bond enerqy

For bonds that occur only in polyatomic molecules such as the C-Hbond, an average bond enthalpy is used.

Example: Average C-H bond energy of CHa is 410 kJ mol-l

(Bond energy values for common covalent bonds are listed in the DataBooklet)Quantitative use of bond energy will be dealt with in 'Reaction Energetics/Thermochemistry' lecture

Covalent Bond Lenqth

The covalent bond length is defined as the distance between the nuclei ofthe two atoms in the bond. lt is also the sum of the covalent radius of thetvvo atoms.

Hydrord chbnd. molccutc HCt

The table below shows the bond length and bond energies of somehalogens.

(b)

Bond dinancc = 0lal .n\

Bond Bond lenqth, nm Bond enerqy, kJ mol-'ct- cl 0.199 242Br- Br o.228 ,t o(I I 0.266 151

Question 5Suggest on explonotion for the decre.rise in bond energy from Clz to fzosshown obove.

-

gla, * a+66 \.tl,,rr',r ,** cl ,o Bt 4 A

-

'fu, +t. 4'o"rrna fd*,r. ltv $rtihe Far. .+ ct.l*a"r ".] to''rf i 'q..le;

p'('o'r.r f,-"'"ct, L z.-

-

gsd Ur.tll }9.c{.ra aL ine|9dl" 4 ala .a'/ Aa/t j62'n L4{r" 'la^^ ClL +. t, ,thd

, so"i .w't{ &c46tt.

l4

bond bond eneroy/kJ mol-' bond length/nmc-c 350 0.154

610 0.134c:c 840 0.'120

Number of bonding electrons and bond enerqv

. The more bonding electrons(or more bonds) between two atoms, thehigher is the bond energy, and thus the shorter and stronger is the bond.

The bonding electrons in a covalent bond are pulled between the twonuclei of the atoms sharing the electrons.

In homonuclear diatomic molecules (H2, Cl2, Oz, Nz etc), the bondingelectrcns are shared equally. The bond is a non'polar covalent bond.

onding nuclei

When two different atoms are joined by covalent bond, the bondingelectrons are shared unequally. The bond is a polar covalent bond, orsimply, a polar bond.

Polarcovalent bond

(Id^rr r -

.&,f,J$.

In a polar covalent bond, the atom that has a stronger attraction for thebonding electrons acquires a partial negative charge; the atom that has aweaker attraction acquires a partial positive charge.

Examole: t '*clt-

t r - ! f o =o N: N

\ t /.,lo^

- FlaY

covah^+ 1,.'x

6",1 ,rlthls,'r, 3- &.*,

l5

The ability of an atom to attract electrons to itself in a covalent bondis called the atom's electronegativity.

-

Electronegativity value is used to estimate whether a given bond will benon-polar covalent, polar covalent, or ionic.

ir^.r{d",

. The difference in electronegativity between tvvo atoms can be used togauge the polarity of the bonding between them.

$eater difference in electronegativity => more polar covalent bond.

Examples

Bond Difference inelectroneqativitV

Type of bond

H-H 0.0 rtw. -

Pole v cwel rrt

H- Cl no ?dav r"ral r"lH_F 1.9 (w0) p"lar coqplo.lNa-Cl- 2.1

Generally, if difference in electronegativity is more than 2, bond is likely tobe ionic.

a

a03e5

t ci'e-ve

ir ruv

.; 80

.9q- 60

e?. 4n

Paulinq's electroneqativitv values of some elements

I tl ltl H2.1

IV VI vtlLi1.0

Be1.5

B2.0

c2.5

N3.0

o 6,Na0.9

Mg1.2

AI1.5

Si1.9

P2.1

S2.5 3.0

K0.8

Ca1.0

Gez.v

As2.0

Se2.4

Br2.8

Rb0.8

Sr1.0,

Sn1.7

Sb1.9

Te2.1

I

Cs0.7

Ba0.9

l6

(t)

Bond polaritv and bond enerqv

. The increase in polarity causes increase in ionic character of the covalentbond, thus lncreasing the strength of the covalent bond'

bond electronegativitydifference betvveen atoms

bond energy/kJ mol-'

H-H 0.0 +3BB

o-H 1.4 +468

F-H1.9 +562

SIMPLE MOLECULES

Tvpes of simple molecules : Non-polar and Polar Molecules

The unequal sharing of the bonding electrons in a polar covalent bondcreates a dipole.

. A dipole is represented by -*-+ pointing towards the moreelectronegative atom.

b+ ' \ - - th-. i < ' H-c1

Size of a dipole is measured by its dipole moment.

more polar bond => stronger dipole moment.

Dipole moment of a molecule = vector sum of all the bond dipole moments(refer to examples on pages 18 & 19)

The dipole moment determines whether a molecule is polar or non-polar.

A non-polar molecule has zero dipole moment aYl\-_\_/A polar molecule is one in which one end of the molecule has a partialpositive charge and the other end has a partial negative charge.

6- ,,--\b (_-_p u-Whether a molecule is polar or non-polar depends on:

(i) the bond moments (i.e. the polarities of bonds in the molecule)(ii) geometry or shape of the molecule (witt be discussed in detait in section 9)

I

I

t

a

I

t1

Non-polar molecules

There are two types of non-polar molecules:

(i) molecules with non-polar bondsExamples: Hz , Clz l-l H | 6 g1 { l'e!. qt"tlu 'u

Polar molecules

A molecule is polar if

. its bonds are polar and,r its shape is not slmmetrical

As a result, the chaqe separation along the bonds does not cancel eachother.

'+ overall dipole moment of molecule is not zero

Examples of oolar molecules

Molecule Shape

NHs.1-

"'ui-i;- -r.- tIT5-

Trigonal pyamidalt+

Hzo bentF{ . . t -H- 0: V/*, 4

H- .

cHct3 Tetrahedrall-lI o '

'7,

I (ll) Bonding between simple molecules: INTERMOLECULAR FORCES

Intermolecular forces are weak aftractive forces between discretemolecules in a simple molecular substance.

These forces are much weaker (less than 15%) than the intramolecularforces (covalent bonds ) that bond the atoms together within themolecules.

Intermolecularattraction (weak)

The strength of the intermolecular forces in a simple molecular substancedetermines its physical properties such as its melting point, boiling pointand solubility.

stronger intermolecular force => higher melting and boiling points

Three types of intermolecular attractive forces exist between neutralmolecules:

. Permanent dipole - permanent dipole attraction \ ,,. &t Mate, &vot

. Instantaneous dipole - induced dipole attraction J

. Hydrogen bonding

Another type of attractive forces, the ion-dipole force, is important inaqueous solution chemistry. This type of force exists between an ion andthe partial charge on the end of a polar molecule.

' ^ ' /H

@' U lrr - Mr$ i!^f'+a't rti '1'zo4 'tv''\

All these four types of forces are electrostatic in nature, involvingattractions between positive and negative species.

20

(a) Permanent dipole-permanent dipole attraction\ah &' h'aalrr l*tt 1ld.

^ Occurs belyeen polar molecules.

Examples of polar.molecules: H11 , tl'o , tolt''

. . Polar molecules aftract each other by weak electrostatic attraction whenthe positive end of one molecule is near the negative end of another-

1@@Such dipole-dipole forces are effective only when polar molecules are veryclose together.

For molecules of similar mass and size, the strcngth of permanent dipole-permanent dipole forces increase with increasing polarity.

molecule M. Dioole moment Boiling point (K)Dimethvl ether, CHTOCH: 46 1.3 244Ethanal, CH3CHO 44 2.7 294Ethanonitrile, CH3CN 41 3.9 355

51 ' - - . '6

- 6r^5-. ) - t\_\,i

^.*.__-/po. *'i'

. This results in a weak and temporary force of electrostatic attractionbetween the dipoles. (ucuatlo)

. Instantaneous dipole-induced dipole attractions are significant only whenthe molecules are very close together.

. This type of attraction is usually weaker than permanent dipole-permanentdipole aftractions.

Strenqth of id-id aftraction

Strength of id-id attraction is mainly determined by two factors:

(i) Number of electrons in molecule ( or size of molecule)Molecule with larger number of electrons (or Larger molecule)=> stronger id-id attraction between the molecules

Reason: Molecules with more electrons are more easily polarised togive a momentary dipole.

Question 6The boiling points of the hologens ore given below.Hologen

FzClzBrzIz

Boilino Loint(K)85.1

238.6332.O457.6

Suggest on explonotion for the voriotion in the boiling point of the hologens.-

Ih haryhr I'ac s1,p ,?holor/ fiLirr "*ale,1 { daret r.olrr,rrr h rl !5 19l

Van' Jrv v0crb' .bvr-

lls l*ry- nar. LAa d ron -Dol av , {& vtrtl Yar' &. toaals' ,

(ii) Shape of the molecule (especially organic molecule)Example

pentane 2.2-dimethvlpropane

cH3cH2cH2cH2cH3 ^. i1:. k"^g

tu,ra4 a.4ta *",21 "tt!;ft "t'fJi

Bp: 36'C Bp: 10 "CM,:.72 M,:72

NOTE:. Instantaneous dipole-induced dipole attractive forces operate between all

molecules, whether polar or non-polar.

. ln fact, id-id forces between large polar molecules may contribute more lointermolecular aftractions than do permanent dipole-permanent dipoleforces.

Question 7Pentone arrd Z,Z-dimethylpropone cre isomers.Boiling point of pentone is 26C higher thon 2,2-dimethylpropone-

n!.!4^t l :(\ '

-

luu' n' l-qtu/al ' )--lG

- !dr' r,tlta

- Efqr'frt *wk,dz t .ot e4E ddal *"nlt tJz a *< e ar,t1 *,rrr :n .r}1 el*V&L

-

tt|!1 i*t.^42 -la,t

er l'4+ro\ |{,Ia..^ clrry]dalf, vc .aild &"r.1 {,tn [r{'r:rn spplrar

-

na , Po'&ac |lrr t.'\r'v f 'e .

sy'l lcal 'ndttula ,

*+'t t'on*taa4

,'oV uiA-

grtr *V/y,tt ir cl^"64Lr> $*.*l llry 61a slol

Z\{'.''l-

4n, t - ,4

. . ra,clc./ ! a

^elr i ,$ ntlrkla l.D a.tlt{-e- . tda 4 6b4. t*z ata4, $** Ia

l' rrulPi 5,a

(c ) Hvdroqen Bondinq.**"What is it?. A special and ptticularly strong type of dipole-dipole attraction between

some molecules.

. lt is an electrostatic attraction between the Ftial positirrely chargedhydrcgen atom in a polar bond (in H-F, H-O or H-N bond only) anda firnepair of electrons on a nearby electronegative atom ( F, O, or N ) on\,another molecule or ion' ik

'v3at'u

,@u"' ,@,@,",',6)

How does it occur?. As F, N, and O are very electronegative atoms, a covalent bond between

hydrogen and any of these three elements is very polar.6-

' , 'o. ' ' '1 -

F"4--H l l r l r r r lF-

L, t-.o o 'ro,'t ..,. ta,r a Ve'X evo"6 prtr cu'4p.

. Thus, the hydrogen atom in the molecule is almost a bare proton of thehydrogen nucleus. This strong positive charge on the hydrogen is stronglyattracted to the negative charge of a lone pair of electrons in F or O or Natom in a nearby molecule.

lntermolecular Hvdrogen bondinq in Water. Hvdroqen Fluoride and Ammonia

"'@:;l@---.**24 "ol " o,- '"?;

evara)a 6

@ry'*7"*' Bra" u-l

Substance H20(t) HF(I) NH3(t)Mr 18 20 17Boilinq poinV'C 100 19

-33Intermolecular Hbonding

.. K -

.J!4.

f,',o1- ,., "tfi".';s 2' A'.s -

".1, ". ".

o. /" , '" H 11 \5.

Two intermolecularH bonds per moleculeil4rt b\ I qq6 3 !ttt.'./

-

r.o trfrlf* tuloY' 6"*1

-

?lt;a ,.a+ erF Q! ? *taa &..V uU' A,st

6-

. Molecular substances held together by hydrogen bonds havehigher melting points and boiling points than molec.ules heldtogether by van der Waals' forces.

Example:Melting points and boiling points of HF, HzO and NH:are higher than other hydrides in the correspondinggroups as illustrated in the following figures.

1'* 'boiling poinl'C

O.

relative molecular mass --...*

- o.? 4 f l9.r ' ' '& rt 2tts ' l 'd 2

ihttrr #a tt 'o 'G H'f. ar

r r+r rgl-' r! v D,$ *t421 i1n14a1*,|t. ruraec:-X no

"l .c- .,ra f

tro )\g ;L d.lr, ,";,r^ D $o'y,l,

2 Solubilitv of some substances

. Some molecular substances dissolve in water by forming hydrogenbonds with water molecules.

Examole: Ethanol in water

A{6- r.r

t. ,/r i l l t \ :0 ,

\brH

Anomalous densitv of ice (to be studied in detail in Solid Structure). The structure of ice reveals the maximum number (4) of hydrogen

bonding interactions between water molecules. These interactionsallow the H2O molecules in ice to be spaced further apart than inliquid water, resulting in a lower density of ice compared to liquid

3

water.

. This accounts for the fact that ice floats on water.A

--(ta,ca- \arYiiAou--{}-

I

Anomalous relative molecular mass. M.

Example 1. Ethanoic acid, CHgCOOH, dissolved in non-aqueous solvents such

as benzene, has an apparent M, of 12O instead of the expectedvalue of 60.

. This is due to presence of two hydrogen bonds which join twoCH:COOH molecules together to form a dimer. The process offorming these dimers is called dimerisation.

. r ro l l l l l l l l l l l l l rH

- 6 a, tc -

cta 3o _ t\ nunt7l|o ,/

Example 2. Liquid (and solid) hydrogen fluoride, HF, has larger M, than

expected because the molecules are associated in groups of twos(HF)2 or threes (HF)3 by very strong hydrogen bonding.

o' -6_

d.r l -

f -7i trttn' H - ':

/ .

- - -0 idtr{ . \\ l t l l t l

26

5 Effect of intramolecular vercus intermolecular hvdroqen bondinq

.t ,o" has a lower boiling point than

X.x/q9

OH2-hydroxyphenol 4-htsoxyphend

u a& y &- 2-hvdoxvohenolO-l r'T,'ria,. ,t Able to form intrarnolecular hydrogen bonding.| 6-n q'

.-t-^-/" " => Hence less sites for intermolecular hydrogenl( ') | bonding to occur. I rntlr 6rt.A.a" U.e norl,t:V

= lowerboiling point

I

^,,H ltlutt 0\ 4_hvdoxvohenolO- \q ---t--,

ts No intramolecular hydrogen bonding can occur.-\ More sites for intermolecular hydrogen bonding.

\-/, ) vpt i**sb& a{rtr'5 '\-

u Y =higherboilingPoint.\ ,oOIITIH-I

27

Tvoe of force Acts between Strength depends onInstantaneous dipole-induced dipole

Between all simplediscrete molecules.Only type of force betweennon-polar molecules.

Permanent dipole-permanent dipole

Between polar molecules(molecules withpermanent dipole)

Hydrogen bonding Between lone pair ofelectrons on F or O or Natom and H atomcovalently bonded to F, O,orNinaneighbour ingmolecule

Summarv of lntermolecular Forces

Relative Strenqth of Ghemical Bonds:

covalent, ionic bond or metallic >>> H-bond > pd-pd aftractions > id-id aftractions

van der Waals forces

Question 8Identify the moin type of intermoleculor force in each of these substonces.

Substcnce moin type of intermoleculor force

Neon, Ne vlrJ {vas l{r ? ;J-i8 arinc{.vl

Corbon dioxide. COz vrd frta & t" tr-b dt&toh

Hydrogen chloride, HCI(g) utN 1{,' &a a- p}- Fd a}t,,.i.o, 1-,;r-t)r} \!. iv, b'r'l''.t

Hydrogen f luoride, HF(l) h{a'+n Bplrrnr| (.{C-?r + D-it) ,,. ^a

plt a.||t &r r01 t

Woter, HeO(l) Hvl'Eh ftv$h,a arpr-PJ + i/-rJ) -ai''reFl

nmmonio, NH:(l) tht{r sna\ (.r0}-C{;r-lr)

28

VSEPR approach-

v\Jr-r r \ qyPr \ - rq\

. lonic and metallic bonds are non-directional, i.e. they extend indefinitely inall directions.

. Covalent bonds, however, have a prefened direction in space. Covalentmolecules have definite geometrical shapes characterised by their bondlengths and bond angles.

. Experimental methods are available for determining the bond lengths andbond angles and thus the shapes of molecules and ions, e.g. microwavespectroscopy, X-ray diffraction, neutron diffraction and electron diffractiontechniques.

. Prediction of the shapes and angles can also be done using the Valence-Shell Electron-Pair Repulsion (VSEPR) model.

Valence-Shell Electron-Pair Reoulsion (VSEPR) model - to predictshapes

.. This model assumes that the spatial arrangement of molecules and ions isdetermined by the repulsion between electron pairs around the centralatom.

' The electron pairs in the valence shell of a central atom in a molecule areof two types:- bond pair (ie, the two electrons in a covalent bond),- lone pair (ie, pair: of electrons not involved in forming a covalent bond)

Between each electron pair and another electron pair, there is a force ofelectrostatic repulsion, which forces the electron pairs as far apart aspossible.

The anangement of electron pairs around the central atom in a moleculedepends on the number of electron pairs.

The anangement of a given number of electron pairs, which minimizesthe repulsion among them, dictates the shape of the molecule or ion.

The shapes of molecules and ions are thus determined by thearrangement of the electron pairs around the central atom, rather thanby the atoms.

29

Relative strenqth of repulsion of lone pai6 and bond pairs

. Lone pairs of electrons in the central atom also repel the bonding pairs ofelectrons.

. Since these ffi6'6ldron pairs are c-loser to the nucleus of the centralatom than bonding pairs, they exercise a greater force of repulsion.

. Consequently, repulsion between electon pairs decreases in the order:

Lone pairJone pair ) Lone pair-bonding pair ) Bonding pair- bonding pairrepulsion p repulsion @ repulsion @

. The unequal repulsion of the different types of electrcn pairs around acentral atom affects the bond angle in a molecule. (refer to examples on page 32-35)

Electron-pairs distribution for minimum repulsionToble 1

\CI. . ' / / - - - )

t-9r-'t ,d) t

w.l z

h'.orLi&

. The spatial anangement of electron pairs about the central atom givenabove is called its electron-pair geometry.

30

Guidelines to predict the shApe of molecules and ions

1 Write the dot-and-cross structure of the molecule or ion.

NHaCHc d'op" rrOo,r0

H

2 Consider the central atom:(i) Countthe IllETtrof *cfron pairs (both bonding and non-bonding type) around the central atom.

(ii) Note that a double or triple bond is counted as one bonding pairwhen predicting geometry. An unpaired electron is counted as alone pair.

(iii) Anange the electron pairs in the way thatGlfrs theilrebetweenthem.

& - ?d'rc e.,''rrl tr "

tt,.f lf l oroQttt l,A

lF,^-V

Thet#ar$t6rfir8'are t:lftlFd from the way the electronbond pairs are oriented in space.

.&4 Determine the shape of the molecule by considedng the positions of

the atoms (i.e. the molecular geometry).Note: the position of lone pairs of elechons cennot b6 located (or 'seen') in theexpedmental determination of shapes of molecules, lhus the shape of themolec1lle is dictated by the arangement of aloms, as prodictd by the electronpair geometry of the molecule.

CHA i6

---

, t " l 'x

rlL'['' h A6, . s\qF,

NHsCH+

3

&..,

NHg A" rT-"'

A;or qya.'Jal

Bond Angle

. Table 'l on page 29 gives the bond angles for molecules with regular orsymmetrical shapes. These shapes and bond angles are obtained only ifall the electron pairs around the central atom are bond pairs only, and thecentral atom is bonded to other identical atoms.

. For some molecules or polyatomic ions, the bond angle may be larger orsmaller than expected of a particular electron-pair geometry.

This may be due to:

(a) Presence of both bond pairs and lone pairs of electrons around thecentral atom

. A bond pair of electrons is attracted by two nuclei of the bondedatoms.

. Lone pair is attracted only by one nucleus (of the central atom)and thus held closer to the central atom and occupy morespace.

. Thus, lone pairs exert greater repulsion forces on adjacentelectron pairs and thus tend to compress the bond anglesbetween the bond pairs.

Thus repulsion between:

Lone pair-lone pair ) Lone pair-bond pair ) Bond pair- bond pair

EgT

'4t-- 'H HI

.1.Ni

'H

,1.;frt^t

Bond angle:

Explanation: As the number of lone pairs of electrons increases,electron repulsion from the lone pair+increases. This compresses thebond pairs of electrons together and thus decreases the bond angle.

(b) Presence of multiole bondsMultiple bonds contain a higher electron charge density than dosingle bonds. Thus, electrons in multiple bonds exert a greaterrepulsive force on adjacent electron pairs than do single bonds.

" \^_,- ,Cl'-- '"-"

Eg.

32

Molecule/ ion

Lewisstrucfure

No. ofbondPair

No. oflonepair

ShapeBondangle

BeClz:Ct '8.-c l ! L 0 cL -b-cL

ltwo,rE0"

Coz02 C z 0 o 0=c=o

lvq,90"

HCNh- a =-Ni ') a H- c = d rgu'

BFs : i :I

IF F"0

Io

I.ttgnot gyral

t70

SnClu

@1a!:ii- si - .i: I

(-')V.. > ,ro"s")

/ - \,X 4r i 'u

u tl 'sls72

zt2o"

NOi l0 = N -r oi.,

o

- lo.tf b''k f i . rv 0

, , n, t* | a ' t to '' Nl

o /-\ o

> t20"

J'iq*,4\ ?len6r.

l?s'

6\drl\o*-:l:1i: r

i 0:r l * g:

Soz- ,^ '23 =; ' .- .o I

a,'.3)' ( t r-c '

x f, A3S. {/.ln,' 6 | ..a, 1L.i7 , 7'

-ra tatt , Aa aa a' 33

Molecule/ ion

Lewisstructure

No. ofbondpair

No. oflonepair

Shape Bondangle

CHrAI

H- c -H}l

+ o

riri. nF'l$.tr{ rr,..t +.k&,I pi,. \, a

bp-bp.1-o,t F ) lnnr 4 ?ri'^

9 q|ol.or*.. laq4 ff;::

los

Soq2-

o g otg-- p, . + n

,.r - + I"* bond

-

' eFPFts

'* ( ) 'a':-/,ilo0 ),.c.

f{.lnz.L&ol

loq

btapt\

So12':.Q..

i0:

s = o? e-- j : '?i4p*t

' |g ' r" ' i la l

rO?

@rfi)

_de $11 , (At ! h,0 16,/*lCnru * shel * # p.o /r"trc.lrr -

34

NorE: lf more than one anangement of lone pairs and bond pairs is possible,choose one that minimises lone pair repulsion.

D In trigonal bipyramidal electron-pair anangement, repulsion is minimisedwhen every lone pair is in the plane of the triangle.

) In an octahedral arrangement with two lone pairs, repulsion is minimisedwhen these electrons are on opposite sides of central atom.

Refer

* ,

ltr {.11 a.l

!p oh

-

9rlOtr

' ( latgA

,.r -!r- r\49. , r.tels{n ldWa" Lp *

l2oo 4, n0 .lt^* I borl lhl t ore {v

I fr|ah}.rlrv' Flort l5

* *" qv att\ 4f ir p..!i |rtr ) ctlo OU owa! ) {*-

^r n.ht I ool 7ll'rr e3

\g "e

nwtntu)$a.a plo/t t

,.ttf t lrv p4tvyar iLk,

ko,c' +a I p39r aorap f"- ?-

ir t,ytL.:r ,.ra!

to exa )elow.Molecule

/ ionLewis

structureNo, ofbondPaar

No. oflonePair

Shape Bondangle

PCl5icLit . .

i (1 -P-cr l

"../ \ i.',!cr cl :

o.

TrigorrlC | "bl nrarlal

t?.. . . l r ,oo{,z- n

c1 /Tto"c7

40"A

t20"

SFa;F;t . .

' .F -

s - F!. . t . .

?t i

+

t -16r tax' l F "or'rl ut

.

06 ? ,4*t'1

F .. l. ?t'rr a' ' l'

Molecule/ ion

Lewisstructure

No. ofbond pair

No. oflone pair

Shape Bondangle

lFs"o i f !ar l F. '" \ - . / . .

t-

. . /, .7 i . '

5E I oo' Ea"

t ly ,

P! !^ ' c l

e/A\ s(Jaot

lc14'['i1.." , ;iil[ri''1 -sI:J

.t

t' ,-,| -,t l I

34,e-tl."x: I '**"[ "u )q0

More on deviations of bond angles

1 Same central atom but different terminal atoms.

Eg. NHg and NFge,,./,.ii--..n

ll o * i" ,F is more electronegative than H, hence the bond pair of electrons inN-F bond is drawn closer to F than its position in N-H bond. Thisresults in less 'crowding' arcund N, which leads to less repulsion andthus smaller bond angle in NFg .

2 Same terminal atoms but different central atom.

Eg. HzO and H2S

e,-"

'-iuH

,Fu.rt

$ is less electronegative than O, henc the bonding pair is further awayfrom S than its position in O-H. This results in less crowding around S,which leads to less repulsion and smaller bond angle in H2S.

JO

Atomic Orbital Aporoach to Shapes of Molecules(to explain/rationalize observed molecular geometry)

10

The VSEPR model provides a simple means for predicting the shapesof molecules. However, it does not explain why bonds exist betweenatoms. In developing theories of covalent bonding, chemists haveapproached the problem from another direction, namely usingquantum mechanics.

(a) Formation of Covalent Bond bv of Overlapplnq of AtomicOrbitals

. When two atoms approach sufficiently close, the valence atomicorbital of one atom merges with that of another atom.

. The atomic orbitals then share a region of space. These orbitalsare said to overlap to form a molecular orbital.

' The overlap allows two electrons of opposite spins to share acommon space between the bonding nuclei, thus forming acovalent bond.

X holr.r\^/ ,fbllr l: o'Y olt* ttad '

Tvpes of covalent bond

o (sigma) bond- can be formed by :

(i) ovedap of two s atomic orbitals,eg. cpvalent bond in a H2 molecule is a o bond

Atoms approach each other

"-*@ @=--n YnY@\overlap

region

(b)

1s atomicorbital ofH atom

1s atomicorbital ofH atom

overlap oftvro 1s orbitalsin H2 molecule

o bond of H,

77

(i i) Jrule headon overlap of two p atomic orbitals,eg. Covalent bond in Cl2 molecule

overlap of an s and a p atomic orbital

eg covalent bond in HCI

d lp,) molecula.orbital

d- bo'r l . l Xal

single region of overlap is

(i i i)

ls 3P

Note that in a sigma bond formation, aproduced.

-R"tF$ 'l q/.'raP i! ab\ {fr firtr 1"r"t'5 {lt ct'tr'r 'l{h 'h^c Lo-' lYt

^u'lt i

2 n (pi) bond- is formed by the sude-wa y overlap of two p orbitals, resulting in 2regions of overlap.

Note that the 2 regions of overlap form 1 covalent n bond.

++ $f*'**"';1- e

q. -trir 16 E.l u|' r^r'kp A"- *

rr (prl molecularo.biral

) wehr {^c.n o- try'l-

n bond formation only occurs where the two atoms are already bondedby a sigma bond in multiple-bonded molecules. For example, thedouble bond of the Oz molecule consists of one a bond and one nbond. et 0] o"o (ro.rn) l {11. ,u

" ."=r- ' ( ,6-+rtr)

Nz a =p ( ,c. : , r )J , r / . h.

The regions of overlap(where the bonding electrons reside) in the nbond are located away from the line of centres of the nuclei. Thus thebonding electrons in the I bond are less firmly held than the electronsin o bond, and the n bond is more easily broken than the o bondduring chemical reactions.

38

SECTION FOR SELF-SfUDY(only required for the understanding of structures of organicmolecules

- methane, ethane and benzene _ in later lectures)

The orbital ovedap model cannot satisfactorily explain the bond formation andobserved geometries in some polyatomic molecuies. The concept ofhydridisation is developed to describe the atomic orbitals used by the centralatom in the bonding of such molecules.

Hybridisation : mixing of two or more atomic orbitals on an atom to createhybrid atomic orbitals. The number of hybrid orbitals created is the sum of thenumber of atomic orbitals used. Each hybrid orbital is identical to the other,but points in a different direction.

Example

sp2 hybrid orbitalsstlown together(larte lobes only)/o-E--=tr

Twop orbitals

Threesp2lrybrid ori/ ita:j

and two p orbitals (anhyb.idize to fo.m threeequivatent d hybrid o.birak.Ine |arge lobes of the hvbrido.bitals point toward th;co.ners of an equilateralt.ian9le.

W eHvbridize F.=--

Tu'o sp hybrid orbitals sp hybrid orbitals shoi^,n together(iarge lobes only)

Olle s orbital

_.19

Bondinq and shape of methane (GHe) molecule

C (ground state) '. 1s22s22p,12pr12p.oCarbon atom in the ground state has two unpaired electrons and henceexpected to form only two covalent bonds. However, it shows tetravalency(forms 4 bonds) in most compounds, for instance, CH4 , CO2

A sharing of four electrons can only be achieved by having the carbon atomundergo excitation and promoting one of the 2s electrons into the 2p orbital.

C'(excited state) : 1s22s12p^12pr12p.1With 4 unpaired electrons, an excited C' should form 4 bonds with 4 H atoms.

-one bond formed byoverlapof C" 2s orbital with lsorbital of one H- three other bonds formed by overlap of each C" 2p orbital with 1s ofeach of the other three H atoms

The bond formed using C* 2s orbital would be expected to be different fromthe other three using the 2p orbitals, and the bond angles would be 90o(sincethe p orbitals are at 90o to each other).

However, experiments have shown that all the C-H bonds in CHr are identical,having the same bond length and strength. All bond angles are 109.5o.This implies that in the CH4 molecule, carbon atom must have used 4equivalent atomic orbitals (instead of using one 2s and three 2p orbitals) forbondingwith4Hatoms.

Hvbridisation model to account for the shape of methane

The C' atom undergoes hybridisation to form four hybrid orbitals of equalenergy, having the same size and shape, before overlapping with the sorbitals of the H atoms.Since one s orbital and three p orbitals are involved in the hybridisationprocess, the new orbitals formed are known as sp' hybrid orbitals.(refer to ne)d page for diagrams)Note:sp3 hybrid orbital is more concentrated in direction than a p orbital, hence it isable to overlap more extensively and form stronger bond than a p orbital.The energy released by the bond formation more than offsets the energyexpended to excite and promote the electrons.

40

sp3 hvbridisation of excited carbon atom

Diagram showsformation of four sp3hybrid orbitals from aset of one s orbital andthree o orbitals.

@to".rdhyb.idrsn \,/-

Overlap of oibitals in methaneThe spa]ti"l

"tra"gernent of th" four sp3 hybrid orbitals is in the form of a

tetrahedron.1s orbital of each H atom overlaps with each sp3 hybrid orbital of the excitedcarbon atom to form a o C-H covalent bond.Thus the shape of the methane molecule is tetrahedral and the bond angle is109.50

.5"

\

4 l

and shape of Ethene lGzll4ilnolecule

C=CH/ 'H

Each C. in ethene undergoes sp2 hybridisation to form three sp' hybridorbitals. One 2p orbital remain unhybridised.

c. 1s 2s 2pl4- tt_l IT-TI-TT-II rv I I r I

\----\'----------l L--unhvbridised2porbital

Orbitals used for sp'^hybridisationTo form 3 hybrid sP'orbitals

Each C atom uses - one sp2 hybrid orbital to form o bond with the othercarbon atom, and

- each of the other two sp' hybrid orbitals to form obond with hydrogen atom.

The three o bonds (C-C , two C-H) formed are in a trigonal planararrangement.

The unhybridised 2p orbital is directed perpendicular to the plane thatcontains the three sp' hybrid orbitals.The unhybridised 2p orbitals on the two carbon atoms overlaplaterally to produce a rr bond.

H

H:Two lobesone ?r burii

Hence, the ethene molecule is a planar (flat shape) molecule with all theatoms on the same plane and the pi-charge cloud above and below the plane.

Bondinq and shape of Benzene molecule

Benzene, CoHo

structural formula of benzene

HHt l.c-c.

H-1/ '-c-H- -

- \ /C:C/ \

HH

arrangement of atoms(this is not the actual structure of benzene)

Each carbon in benzene is sp2 hybridized (i.e. it has three sp2 hybridatomic orbitals and one unhybridised p orbital)

Each carbon uses - two of the sp2 hybrid orbitals to form sigma, o , bondswith two neighbouring carbon atoms, and

- one sp'orbital to form a o bond with the 1s orbital ofhydrogen atom

of the carbon and hydrogen atoms.

U.rtrkn.i,(5 ( n.rrtt r(krckt..d.kj i&{.. Each carbon still has one unhybridised p atomic orbital which is at rightangles to the plane of the o bonds.

. These six p orbitals overlap to form a continuous n molecular orbital aboveand below the plane of the molecule.

. The continuous n molecular orbital is also called delocalised n electroncloud, as the electrons are not confined between any two atoms, but arefree to move over the entire carbon skeleton.

(Note : The ring inside the structural formula of benzene indicates thepresence of this detocalised n electron cloud.). The delocalisation of these n electrons makes the benzene molecule

energetically very stable.

' - \- ,- f-"b4 -\

41