7-1

COVALENT COMPOUNDS - ACIDS - MOECULAR GEOMETRY -INTERMOLECULAR FORCES

WHY DO ATOMS SHARE ELECTRONS? (pp. 359 - 363)

1. In the last unit we learned that some metals and nonmetals react to form binary ionic compounds.Electrons are transferred and the resulting ions have noble gas configurations. Compounds are then formedbecause the ions are attracted to one another.

2. Sometimes two atoms that both need to gain valence electrons to become stable have a similar attractionfor electrons.

3. Sharing electrons is another way these atoms can acquire the electron configuration of a noble gas, eventhough it will be on a part-time basis.

4. In a Co~AL8-JT l3ol-J D , atoms do not lose or gain electrons. Instead, they share pairs ofelectrons to achieve stability, often by filling their outer energy levels so they have stable octets.

5. A MOll::Lll LE is formed when two or more atoms bond covalently. They are often calledJ)"\OLEC-UL-Ai2.. CoMAlV,v D ALSO CALLEb CCvALBJ', C()I\"A:lV~QS

Forces of electric attraction make a covalent bond

. .' ·::·)\~IV:;·.'.'Hat~m

:\i~j;\'.·.Hatom:

Sufficiently far apart. to have no interaction

Halom Halom

The atoms begin 10 interactas they move closer together.

Hl molecule

Optimum distance to achieve(a) lowest overall energy of system (b)

t(H-H bond length)

1. An attractive force exists between the outer electrons on one atom and the nucleus of a neighboringatom.

2. The force of attraction brings the atoms together until the force of repulsion between the nuclei andbetween the outer electrons forces the atoms apart.

7-2

Electron cloud

Nucleus

3. If the forces of attraction are greater than the forces of repulsion, then a covalent bond forms betweenthe atoms.

4. Besides the comparative strengths of the attractive and repulsive forces, another reason the attractiveforces can be stronger is that a pair of electrons shared between atoms in a stable covalent bond haveopposite spins and occupy less space than a pair of electrons in an orbital or only one atom.

5. The bond is not rigid. It is much like a spring where the atoms vibrate back and forth at some averagedistance where the attractive and repulsive forces are balanced.

Sharing More Than Two Electrons

I. Covalent bonds between atoms can involve sharing more than two electrons.

2. When a single pair ofelectro~s (2 electrons) is sh~~ed; this is kn~wn as a SI/.JGLE. !30"LD

3. When two pairs of electrons (4 electrons) are shared, this is known as a .ooV~~ 6otJ.I>

4. When three pairs of electrons (6 electrons) are shared, this is known as a Tt?IPL.E &tJ,D

Bond Length and Bond Energy are InverselvRelated

1. The average distance that separates the atoms in a bond is known as the6c IJ D LeAlG, T1-J

2. Bond lengths are never really fixed values because the atoms vibrate. They can also vary depending onthe other bonds present in the molecule.

3. ~O N.D /3.JEI2. G ..., is the energy required to break a chemical bond to produceindividual atoms, each keeping its own electrons.

4. Bond length and bond energy are inversely related.

S. A short bond length requires higher bond energy to break it while a long bond length requires lessenergy to break it.

Bond Properties

1. Few chemical bonds are either totally fl10LEC.uuVl .... or totally .:to I-J Ie2. The bonds in many compounds have some features of both types of bonds.

7-3

3. The electrons in a bond are not necessarily shared equally. To determine whether this uneven sharingwill be very small or very large, one compares the ability of each atom to pull electrons toward itself.

4. This is called fLf.C.~N EGATI\/ I 'T 15. The electronegativity table is used to provide numbers for comparison.

6. The greater the differences in electronegativities between two atoms, the more unequal the sharing andthe more ionic character the bond will have.

7. A covalent bond formed between two atoms with equally shared bonding electrons is said to be aA/o/IJPOLflR- COyAL$.IT f:JJ}JD

8. Examples are:

ff-I-/ ) 0-0, F'-F

9. When atoms of different elements bond, the sharing of electrons can never be truly equal.

10. When a covalent bond is formed between two atoms in which the bonding electrons are more stronglyattracted to one atom over the other is said to be a fJOLAR- Cf)VIHEJ..tT fX)tJ.j)

11. Examples are:Rb-O) f-).(-Ai) c.-O

12. To determine bond polarity, we make use of Pauling's electronegativity values:

Pauling Electronegativity Values

5 6 789BeN 0 F

2.04 2.55 3.04 3.44 3.98

16 17S CI

2.58 3.16

34 35Se Br

2.55

7-4

13. Look up the electronegativity value of each atom in the bond and then subtract the smaller value fromthe larger value.

14. The difference is always ----CA--"O_.$_, T_I\_'I.::_= _

15. Use this table to determine the bond polarity:

How to Tell the Type of Bondthat will Form Between Two Elements

Type of BondDifference in

ElectronegativityNon-polar covalent less than 0.5Polar covalent between 0.5 and 2.1Ionic ~ greater th.an 2.1

Electronegativity Electronegativity Difference ofof element of element electronegativities

Cl,3.16 CI,3.16 0

C,2·55 0,3·44 0.89

Type of bond

non-polar covalent- electro ns areshared equally between bothchlorines, Ch

polar covalent, oxygen is morcelectronegative so the electrons arecloser to oxygen makingoxygen slightly negative andthe carbon slightly positive

CI,3·16Na,0.93 2.23 ionic, chlorine "steals" one electronfrom sodium to make.CI-and Na+,NaCI

16. The uneven sharing causes the more electronegative atom to have a partial negative charge while theless electronegative atom will have a partial positive charge.

***** Determine the bond polarity of the following bonds:

C-H.J.Ss. - ;1,,),0 z .3S

N-HJ..()L/- didO ~ • ~l{

coo3. <-JV- .J. .s s = . F1

POLRR

7-5

NAMING COVALENT COMPOUNDS (pp 94 - 97)

1. Naming covalent (molecular) compounds is similar to naming ionic compounds.

2. One can use either the Stock naming system or one that makes use of prefixes, roots, and suffixes.

3. The latter system is known simply as the prefix naming system.

4. The root comes from the name of the element and then prefixes and suffixes are added.

5. An example is:

CO CIlt<t!iJrJ /110,.10 X I te:

6. The first element named is usually the one with the lower electronegativity value.

7. Ifthere is only one atom of the first element, then N'O prefix is used.

8. The ending - ::rDE is used as it was in naming ionic compounds.

9. The common prefixes are:

1- 1l1.0tJO

a. - j) I

3 - 1ll:1

tJ - TETTl.4

$ - PeJT7~

~ - HEX4

7 - HEP7J:J

~ - (lCT"t1

q - MOl-! f.l

10 - bEeR

***** Name the following compounds:

P20S

t>l'PI-\O$Pl-\oR.u~ PENTO'j(lbt:

AS203

1:>1At2Je./1erRl ox ae:

NF3Ai I TROfttJ 77<1FLUOR.. IDe

CCl4CAe.t30/oJ 7ETR.lJC.HLolZlus

802

SO LFUf2-. DI OX , D£

10. The Stock system can be used to name covalent compounds.

P20S

Pi-\oSPI-IOe.VS Cv) O)<.I.Dt=

CC14

CA.090J-./ (1'1) CHi.l>/2IDI2"NF3

/II117206e.i (JIL) FLt)()R..1 De-

As20S

A-R.s.I9J/C (v) DXI);£

S02SVLF\)R. (I") OXIDe

7-6

Writing Formulas for Molecular Compounds

1. If the compound is named using the prefix system, simply translate the prefixes as written intosubscripts.

***** Write the formula for each of the following molecular compounds:

dinitrogen tetroxide phosphorus trichloride

disulfur trioxide

NAMING ACIDS (pp 104 - 1(5)

1. Water solutions of some molecules are acidic and are named as acids.

2. A binary acid contains:

H'lDe.OW

/1,.1 AN/OJ

But, no 0 'f. 'fc;8J

3. When naming a binary acid, use the prefix H'fj)~O to name the hydrogen part of the compound.

4. The rest of the name consists of a form of the root of the second element, or polyatomic ion, plus thesuffix - .Ie

5. The word Ae. I j) is then added.

HBr in a water solution is known as: H Y])Il.Of3R1Jm I ( I1C I D

***** Name the following acids:

HFWIl)'lCi FLvltJ c fJ C Io HeN

I~'IfY.l.()C 'I4,.Ju: t:K- I b

HIH'I Df!.O.I()~/C Ilt 1.0

HzSH- '1112OSVLFUe.f c. lie.Ij)

Naming Oxyacids

1. Any acid that contains hydrogen and an oxyanion is called an 6 XYf.';C. 1.1)

2. To name it, first identify the anion present.

3. The name of the oxyacid consists of the root ofthe anion, a suffix and the word acid.

4. If the anion suffix is ATE , change it to r C.

5. If the anion suffix is 0 U.s , change it to ~

ATE - Ie.) .J:TE - oo.s

7-7

Examples:

HN02 /'.Wn<.o U5 ACI.D

HCeG3 C-HLOeIC ACID

HCe04 PEk!..CI-\L-O~« AC 1.0

HC2H302 AGE1l (. ACID

6. Notice there is }.iO use of the hydro.

7. The following oxyacids were named before the rules went into effect, so they must be memorized:

MOLECULAR GEOMETRY - LEWIS STRUCTURES (pp. 371- 381)

1. In order to predict the arrangement of atoms in a molecule, a model is used.

2. The nuclei and inner-shell electrons are represented by the element's atomic symbol. The valenceelectrons are represented as dots placed around each side of the symbol, up to two per side.

Paired.elec.trons

-: _J~""__ Unpaired• • electron:Cl·/

Unpaired'electrons, .•

• c·..' •Chlorine Carbon

7-8

3. Bonds between atoms are represented by either pairs of dots or by lines between the atoms involved inthe bond.

4. Unshared pairs are represented by pairs of dots placed around the appropriate atoms.

..' • • • •·F·N·F • •• •• • •• • • • :I? N I?:•• • • • • - -• F • •• ••• • Iel? e• • • •••

DRAWING LEWIS STRUCTURES

1. Determine the total number of valence electrons in the compound by adding up all the valence electronsof the atoms in the compound.

2. Arrange all the element symbols according to which element can form more than one bond and thosethat can only form one bond.

** Atoms that can only form one bond are: Hand F. Ct, Br, and I will normally form only one bond unlessoutnumbered by either ° or F.

** Atoms that "love" being in the middle of things are: B, C, N, 0, Si, P, S, As, Se, Sb

3. Draw single lines between all the atoms that are bonded together.

4. Count each line, multiply by two, and subtract that number from the total number of valence electrons.

5. This gives you the electrons left to distribute to all the elements still needing electrons so each can havean octet.

6. Ifthere are not enough electrons available to distribute, then one or more ofthe single bonds may haveto made into double or triple bonds.

***** Draw Lewis structure forthe following molecules:

a. Iodomethane (CH3I)

C-<= tie- H-i o •.3H ~ 3e- H -C.- r:

7e - , ,III= 1-+vAL e. = 14f-lot-lD e- ss:

ujJS H.1.e.Q) e- = f.t

u.$tt)e ~ ~- !2£7n4INfA/(;, e: .: 0

b. Methanol, (CH)OH)C: Lfe.-

i-lH= -l e :e): ~e-

vAL€,- :: /4eUrJD t- : /Du I\J$ H4ZI::1) e- z: tf

us,=1) e - s:3-o

c. Dinitrogen difluoride (N2F2).)Jj,IO

JF s: / 'I\/liL e:» ;Jy

t3.o,.;j) ~ - s: ~ ~

,}i\I.s~4fri.)e z: ~ I~

v~{:l::)e - -r: "

o

d. Formaldehyde (H2CO).l~:J-

C = 40·-:: (,

vt}L e - .: / J..t3.0rJD e .-.:: '6r. <t

UI'J.5H4·e(1) e- .s: '" <../

US(:1) e - s: '-I.!--o

e. Hydroxide ion (OH-)e.- F.e.un1 CHt\,eGf ~ I

It,

I\fAu:~Cf e s: ~

&;.JD e - - ~\.Il\lSI-IJ\U:1) e. - ~ (,\.rS8~ e - z: (,

o

7-9

1-1I .,

H-C-O-HI & r

1-1, .,.C - 0.-f ~ (j

H

7-10

f. Ammonium ion (NH4+)tJ~.s

I.{H=- ,,-(

VAL e. -~O,J\) €. - :

.+/-JI

H-tJ-l-fIH

L.ose e - I

RESONANCE STRUCTURES1. Sometimes a single Lewis structure is not enough to accurately depict a molecule.

2. When this happens, more than one equivalent structure is used to represent the molecule.

3. When more than one equivalent Lewis structure can be drawn for a molecule, the molecule is said to bea I<ESo~A.-J ex; hybrid.

An example is S03

••

All of these structures are equivalent and only the double bond moves to different oxygens around thesulfur.

***** Draw the three resonance structures for dinitrogen monoxide (N20)

vAL .Q,-

~IJ]) e.-lJluS~t.\(8)e~ - ~ l'O ~0.$-1::1) ~ - ~ 8' o ~;j~(-'" tJ

• f/ , II

'f ./ -tJ'(J;:rJi (-:7,"-.\-0;:: ,• ...

o

7-11

EXCEPTIONS TO THE OCTET RULE

1. There are three types of ions or molecules that do not follow the octet rule:

(a) Ions or molecules with an odd number of electrons (N02)

•N•• ~ "<, ••o O.•• •• •

(b) Ions or molecules with less that an octet around the central atom (BF3)

• •• F •• •

• • I• F B•

• • I• F •• •

• •(c). Ions or molecules with more than eight valence electrons around the central atom (an expanded octet)

PCIs

••tcr: ••

• • I ci ;:CI -p/::.. I"CI::Cl: ..

• •

7-12

IFs

••• F:•

•• I • • •:F I • F •••• /1 • •• F •• F: • •• • •• •

2. It is thought that the extra electrons go into empty "d" orbitals, thus permitting the central atom toexceed the octet rule.

3. When it is necessary to exceed the octet rule for one of several third row (or higher) elements, assumethe extra electrons should be placed on the central atom. .

MOLECULAR GEOMETRY - DETERMINING THE SHAPE OF THE MOLECULE(pp 381 - 392)

1. Shape is an important factor in determining the chemical properties of a molecule.

7-13

2. One example is the difference between normal, healthy red blood cells and the shape of a sickle cell.

1 2

3. The theo used to predict shapes is called the VALeJC.f SHELL ElEcnw,J ,-:nIt( R,EftJLS (oAlor ( VSE ) theory.

4. VSEPR is based on the idea that electron pairs surrounding the central atom will arrange themselves tobe as far apart as possible.

(a) THE LEW1S

5. Shapes cannot be predicted from the molecular formula. One must know:

(b) TIll: 80rJD/tJ G (SHlJ.eED) PAIi2J of I=L£CT1C.O;JS IIT77-JCHED

10 TH E. Ci?NTJ(?A L fFTO fl1

6. Once the bonding pairs and nonbonding pairs attached to the central atom have been determined, usethis chart to determine the molecular geometry (shape) of the molecule:

7-14

BGO. 180"2 2 o

Ald. 3 o

CH1 4 0.,.

IOU· spll

t.tr~

.f 3 .t 10"l.V Si~

liiV-"aa ~_ktat

X. "" '.

H,.O 2 2 l04S· sp-

-.m~Qltohalront.a1)

t4tar. 5 S 0 9CrnlO'" sp"d.

(hCfllOllbito hortlootoll)

Sf. 6 0

*~ff' ~.

Oct.ah.d"aI

'8.1& fGProsant ~to~ web raprQllQOt bonds; ¥.ci lobGi r~nt klnGpolin. 01 alaC1ron\.

7-15

TOTAL PAIRS BONDING NONBONDING MOLECULAR EXAMPLEPAIRS PAIRS GEOMETRY

2 2 0 Linear CO2

3 3 0 Trigonal planar BF3

3 2 1 Bent N02"

4 4 0 Tetrahedral CH4

4 3 1 Trigonal pyramidal NH3

4 2 2 Bent H2OTrigonal

5 5 0 bipyramidal PCIs

5 4 1 See-saw SF4

5 3 2 T-shaped cir,

5 2 3 Linear XeF2

6 6 0 Octahedral SF6

6 5 I Square pyramidal BrFs

6 4 2 Square planar XeF4

7-16

VSEPR SUMMARY

Number of ~Iedronpairs around ~tral: Full descriptIO" of _ mole<:ule

atom .'

9O"0INi>l'!), . lOtE.(E)· Exampl~ . Bond angles'·. Geometrvof ...Gl<lmell'lof· 30 Shape TypeElectron Pal,.. Alom~··

2 0 B9C~ 180 Linear · . linear· ( ) AB2

Ju "

, . " Tngonar,..3 0 . 8F] 120· . Trigonal. A83planar .' Planar.. ..' "

2 '1 $02 . Slightly less . TrtQonal Bent or Ju AB2E .than 120. planar . VShaped

'4 0. CH. ·10.9.5 ....Tetrahedritl Tetrahedral erlo .AB.: '.. , ~.•..'

.'.

0" ~ • ' . cyfu..':' Trigonal3 1. .. NH3" 107.5, . Tetrahedral, AB~,

..... ·,Pyramidal,'.

,.Bent·or: ';

~

.., 2 2 H2O .. 104,5' , .' Tetrahedral' ' · V Shaped: .' AB£2-,

.120 In plane; " l )

,0 'PCls

90. Trigonal Trigonal. ", ().~..JII~ 'ABi~5..

perpendicular .. blpyramldill.· Blpyramldal''. " to plane , . ", . (") . ',.

4 1 .', . SF. Comr»exTngOnal' Seesaw ()l,l) Ae~Eblpyramid '-'C1~,-,

()3 2

.'CIF~ Approx,90. Trigonal r.Shaped >~ AB~;blpyramldaJ

0C)

Trigonal:..

2 3 XeF2 180 Linear' .>..-- A8:zE] :bipyramid, . .)

. ().' " Q..::.a ~..()

6 0 SF, . 90. Octahedral Octahadral ABsC1'()-0..

o5 1 . BrFs Approx.9o. Octahedral' Square, , 0. ......•IIL.•• .Q ABs£:Pyramidal cr: -0

4 2 XeF4 90 Octahedral Square o.~ L"'.o AB,J:2Planar if -0

***** D t . he ermme t e shape of each of the following:

a. Ammonia (NH3)

'\\::5~l-\ = 3

e (0

~ -I\J - l~I

HVA Le- .: g'&Jr-.iDe- z: ~--UNSW\t2.ED ~ J

~d-ob. Water (H20)",:11-{ = i).,

0;- (p H - 0 - \-\vALe- r: r"fu)j"D e- z '-f

o •

oc. Dichlorodifluoromethane (CF2Ce2) , •

C-==- ~ r: p~ I- 6

.JF'= ILl Ie ,

J~.:- I L( .~vAle,- s: .3 ;:;l~·Nbe-:: ~

- J~.s\{Alro .; d.J{-~

od. Arsenic pentafluoride (AsF5)

i\s s: SSF:; 3S

e oJ ."~F '• •.,

& cG r f-a r- c

/

-JAl e = J.(o'&CI.-.Ju e s: r ()-------\J t--l-S I·Vl!l~j:)': 3D

~.jO-~oe. Selenium hexabromide (SeBr6)

Se :: &~ 5r s: '-I a

• F .'f e<t 0

~,& a. '• f.J r 0

I

Se.I

,.VAL e- s: t..f f

&,,{D € - - I J,..

UVSf.f.AlED: 3 (P

-3&-ot c

7-17

Tb77tL flAf6 s: '-IBOAl D flA/tiS r: 3LON!: P/1lltJ s: /

tj-3-1

TIZ I GoAi4 L PY£flM I D,q t:

TVTI'.lL PitIf2f:: ~s»ND pl/lta·- J..UJrJE PI}I!2J·- d

~-J..-d...

Bcl\1 r=

11ITY-l L p,q I«S> VBOND P4//(S:- 1-/

. LoA/r= ?/-l/t2.S .r: 0

e • TVTAL P/JIr<S;; 5'Bo",/) I1Jllt5 ~ SLONe PIJlI2J:- 0

S-s-'O

7?2/GoAJA L 131 PYI2I1fl1IDIf- L

ToTAL ffi lies :- (p50l-.1D ?AI (2..S ' (u

Lo~(; t'AI/2S =- 0

~-(p-o

OCIAHtD..eAL

7-18

MOLECULAR DIPOLE

I. Remember that each bond within a molecule can be either nonpolar or polar.

2. Ifthe bond is polar, then one end of the bond appears to have a slight positive charge while the otherappears to have a slight negative charge.

3. This creates a dipole.

4. If the molecule has polar bonds and the shape of the molecule causes the polarities to cancel, then themolecule is tJONPoLAe .

5. If the molecule has polar bonds and the shape of the molecule does not cause the polarities to cancel,then the molecule will be PoLAR. .

6. If the molecule has all nonpolar bonds, then no matter what the shape is, the molecule will beAIoA/Pc LI~R... .

***** Determine the polarity of each of the following molecules:

Ammonia (NH3) Po LA Q....

Water (H20) POL.~

Dichlordifluoromethane (CF2CtZ) PoL./\Z

Arsenic pentafluoride (AsF 5) N,o1-1 Po LA./(

Selenium hexabromide (SeBr6) "-lO N PoLA f2..

INTERMOLECULAR FORCES (pp. 442 - 444)

I. All atoms and molecules attract each other. But, these forces of attraction are not as strong as the forcesof attraction between atoms in the bonding process (ionic and covalent bonds).

2 . .:;:)IjTEJ?...n'{OLfCt'l.4!l... Fl7tZCE5 are forces that cause attractions between molecules.

3. Intermolecular forces of attraction between molecules, or between atoms and molecules, do not involveeither the transferring or sharing of electrons.

4. The weakest intermolecular forces are UHJ.DoJJ bI5PER.~/O,J RJf2.CEJforces of attraction between nonpolar substances.

. These are the

7-19

London dispersion force (1'2 bond)

Electrostaticattraction

=1Helium atom 1 Helium atom 2

Examples are: Br2, h, N2, ci, H2, 02, F2, CH4, and the noble gases

5. Next in increasing strength are iJlPoLE - DIPOLE FQ~CE5forces occur between polar molecules because there is a dipole in the molecule.molecule attracts the negative end of a nearby molecule.

. These intermolecularThe positive end of one

- The interaction between anytwo opposite charges isattractive (solid red lines). -

-.The interaction between anytwo like charges is repulsive(dashed blue lines).

Examples are: CO, NO

7-20

Attraction -----...-

Repulsion ..---.

6. The last type of intermolecular force is known as H'l!)eo Gru "EOf../vrtJ 6especially strong force of attraction occurs between molecules containing hydrogen bonded directly to a

highly electronegative atom such as F 0 or JJ

. This

2S_J:.S+

~8+(a)