© 2006 brooks/cole - thomson writing lewis formulas: the octet rule the octet rule states that...

© 2006 Brooks/Cole - Thomson

Writing Lewis Formulas: The Octet Rule• The octet rule states that representative elements usually attain stable noble gas electron

configurations in most of their compounds.• Lewis dot formulas are based on the octet rule.• We need to distinguish between bonding (or shared) electrons and nonbonding (or

unshared or lone pairs) of electrons.

• N - A = S rule– Simple mathematical relationship to help us write Lewis dot formulas.

• N = number of electrons needed to achieve a noble gas configuration.– N usually has a value of 8 for representative elements. – N has a value of 2 for H atoms.

• A = number of electrons available in valence shells of the atoms.– A is equal to the periodic group number for each element. – A is equal to 8 for the noble gases.

• S = number of electrons shared in bonds.• A-S = number of electrons in unshared, lone, pairs.


Writing Lewis Formulas: The Octet Rule1. For ions we must adjust the number of electrons available, A.

a. Add one e- to A for each negative charge.

b. Subtract one e- from A for each positive charge.

2. The central atom in a molecule or polyatomic ion is determined by:

a. The atom that requires the largest number of electrons to complete its octet goes in the center.

b. For two atoms in the same periodic group, the less electronegative element goes in the center.

3. Select a reasonable skeletona. The least electronegative is the central atomb. Carbon makes 2,3, or 4 bondsc. Nitrogen makes 1(rarely), 2,3, or 4 bondsd. Oxygen makes 1, 2(usually), or 3 bondse. Oxygen bonds to itself only as O2 or O3, peroxides, or superoxidesf. Ternary acids (those containing 3 elements) hydrogen bonds to the oxygen, not

the central atom, except phosphatesg. For ions or molecules with more than one central atom the most symmetrical

skeleton is used4. Calculate N, S, and A


Writing Lewis Formulas:The Octet Rule

• Write Lewis dot and dash formulas for hydrogen cyanide, HCN.• N =• A =• S (shared electrons)• A-S (lone pair electrons)



• Write Lewis dot and dash formulas for the sulfite ion, SO32-.

N =

A =

S (shared electrons)

A-S (lone pairs)



• What kind of covalent bonds, single, double, or triple, must this ion have so that the six shared electrons are used to attach the three O atoms to the S atom?

O S O

O··

····

····

··

··

··

····

··

····

2-O S

O

O·· ·· ··

······ ··

······

2-or


Resonance

• Write Lewis dot and dash formulas for sulfur trioxide, SO3.

N =A =S A-S


Resonance

• There are three possible structures for SO3.– The double bond can be placed in one of three places.

O S

O

O·· ······ ··

······

OS

O

O·· ···· ·· ··

··

······

O S

O

O·· ····

·· ··

····

oWhen two or more Lewis formulas are necessary to show the bonding in a molecule, we must use equivalent resonance structures to show the molecule’s structure.

oDouble-headed arrows are used to indicate resonance formulas.


Writing Lewis Formulas:Limitations of the Octet Rule

• There are some molecules that violate the octet rule.– For these molecules the N - A = S rule does not apply:

1. The covalent compounds of Be.2. The covalent compounds of the IIIA Group.3. Species which contain an odd number of electrons.4. Species in which the central element must have a share of

more than 8 valence electrons to accommodate all of the substituents.

5. Compounds of the d- and f-transition metals.



• In those cases where the octet rule does not apply, the substituents attached to the central atom nearly always attain noble gas configurations.

• The central atom does not have a noble gas configuration but may have fewer than 8 (exceptions 1, 2, & 3) or more than 8 (exceptions 4 & 5).



• Write dot and dash formulas for BBr3.

– This is an example of exception #2; The covalent compounds of the IIIA Group.



• Write dot and dash formulas for AsF5.

– This is an example of rule 4; Species in which the central element must have a share of more than 8 valence electrons to accommodate all of the substituents.


Covalent Bonding• Covalent bonds are formed when atoms share electrons.• If the atoms share 2 electrons a single covalent bond is

formed.• If the atoms share 4 electrons a double covalent bond is

formed.• If the atoms share 6 electrons a triple covalent bond is

formed.– The attraction between the electrons is electrostatic

in nature• The atoms have a lower potential energy when

bound.


Formation of Covalent Bonds

• We can use Lewis dot formulas to show covalent bond formation.

1. H molecule formation representation.+H. H . H H.. or H2

H Cl H Cl+...

.... ..

..

..

... or HCl

2. HCl molecule formation

• Some examples of nonpolar covalent bonds.

• H2 H HorH H..

N N········ ·· N N·· ··or• N2


Polar and Nonpolar Covalent Bonds

• Covalent bonds in which the electrons are shared equally are designated as nonpolar covalent bonds.– Nonpolar covalent bonds have a symmetrical charge

distribution.• To be nonpolar the two atoms involved in the bond

must be the same element to share equally.



• Covalent bonds in which the electrons are not shared equally are designated as polar covalent bonds– Polar covalent bonds have an asymmetrical charge

distribution• To be a polar covalent bond the two atoms involved in the bond

must have different electronegativities.• Some examples of polar covalent bonds.• HF

bondpolar very 1.9 Difference

4.0 2.1 ativitiesElectroneg

F H

1.9

2.1 4.0



• Compare HF to HI.

bondpolar slightly 0.4 Difference

2.5 2.1 ativitiesElectroneg

I H

0.4

2.1 2.5



• Polar molecules can be attracted by magnetic and electric fields.


Dipole Moments• Molecules whose centers of positive and negative charge do not coincide, have an

asymmetric charge distribution, and are polar.

– These molecules have a dipole moment.

• The dipole moment has the symbol . is the product of the distance,d, separating charges of equal magnitude and opposite

sign, and the magnitude of the charge, q.

units Debye0.38 units Debye1.91

I- H F- H

--

• There are some nonpolar molecules that have polar bonds.• There are two conditions that must be true for a molecule to be polar.

1. There must be at least one polar bond present or one lone pair of electrons.2. The polar bonds, if there are more than one, and lone pairs must be arranged

so that their dipole moments do not cancel one another.


The Continuous Range of Bonding Types• Covalent and ionic bonding represent two extremes.

1. In pure covalent bonds electrons are equally shared by the atoms.

2. In pure ionic bonds electrons are completely lost or gained by one of the atoms.

• Most compounds fall somewhere between these two extremes.

• All bonds have some ionic and some covalent character.

– For example, HI is about 17% ionic• The greater the electronegativity differences the more polar the bond.


Formal Charges

― The hypothetical charge on an atom in a covalently bonded molecule or ion; bonding electrons are counted as if equally shared.

1. FC = (Group #) – [(number of bonds) + number of unshared e-)]

2. For Lewis dot formulas an atom that has the same number of bonds as its group number has an FC of zero

3. The sum of formal charges is equal to zero for molecules and equal to the charge of the ion for ions.

Ammomia – NH3 Ammonium Ion – NH4+

Lewis Dot

FC =


Formal Charges

― The most likely formula for a molecule or ion is usually the one in which the formal charge on each atom is zero or as near zero as possible

― Negative formal charges are more likely to occur on the more electronegative elements

― Lewis dot formulas in which adjacent atoms have formula charges of the same sign are usually not accurate

Cl=N-O Cl=N-O


Stereochemistry

• Stereochemistry is the study of the three dimensional shapes of molecules.

• Valence Shell Electron Pair Repulsion Theory

• Commonly designated as VSEPR

• Principal originator

– R. J. Gillespie in the 1950’s

• Valence Bond Theory

• Involves the use of hybridized atomic orbitals

• Principal originator

– L. Pauling in the 1930’s & 40’s


• The same basic approach will be used in every example of molecular structure prediction:


Polar Molecules: The Influence of Molecular Geometry

• Molecular geometry affects molecular polarity.– Due to the effect of the bond dipoles and how they either

cancel or reinforce each other.

A B A

linear molecule nonpolar

A B A

angular molecule

polar

• Polar Molecules must meet two requirements:1. One polar bond or one lone pair of electrons on central

atom.2. Neither bonds nor lone pairs can be symmetrically

arranged that their polarities cancel.


VSEPR Theory• Regions of high electron density around the central atom are arranged as far apart as

possible to minimize repulsions.

• There are five basic molecular shapes based on the number of regions of high electron density around the central atom.

• Lone pairs of electrons (unshared pairs) require more volume than shared pairs.

– Consequently, there is an ordering of repulsions of electrons around central atom.

• Criteria for the ordering of the repulsions:

1 Lone pair to lone pair is the strongest repulsion.

2 Lone pair to bonding pair is intermediate repulsion.

3 Bonding pair to bonding pair is weakest repulsion.

• Mnemonic for repulsion strengths

lp/lp > lp/bp > bp/bp

• Lone pair to lone pair repulsion is why bond angles in water are less than 109.5o.


VSEPR Theory

• Frequently, we will describe two geometries for each molecule. 1.1. Electronic geometryElectronic geometry is determined by the locations of regions of

high electron density around the central atom(s).2.2. Molecular geometryMolecular geometry determined by the

arrangement of atoms around the central atom(s).Electron pairs are not used in the molecular geometry determination just the positions of the atoms in the molecule are used.


VSEPR Theory• Two regions of high electron density around the central atom.

• Three regions of high electron density around the central atom.

• Four regions of high electron density around the central atom.


VSEPR Theory• Five regions of high electron density around the

central atom.

• Six regions of high electron density around the central atom.


VSEPR Theory• An example of a molecule that has different electronic and molecular

geometries is water - H2O.

• Electronic geometry is tetrahedral.• Molecular geometry is bent or angular.

H

C

HHH

H

C

HHH

• An example of a molecule that has the same electronic and molecular geometries is methane - CH4.

• Electronic and molecular geometries are tetrahedral.


Valence Bond (VB) Theory

Regions of High Electron Density

Electronic Geometry

Hybridization

2 Linear sp

3 Trigonal planar sp2

4 Tetrahedral sp3

5 Trigonal bipyramidal

sp3d

6 Octahedral sp3d2


Molecular Shapes and BondingMolecular Shapes and Bonding

• In the next sections we will use the following terminology:A = central atomB = bonding pairs around central atomU = lone pairs around central atom

• For example:AB3U designates that there are 3 bonding pairs and 1 lone pair

around the central atom.


Linear Electronic Geometry:AB2 Species (No Lone Pairs of

Electrons on A)

1s 2s 2p

Be

1s sp hybrid 2p


Trigonal Planar Electronic Geometry: AB3 Species (No Lone

Pairs of Electrons on A)

1s 2s 2p

B 1s sp2 hybrid


Tetrahedral Electronic Geometry: AB

4 Species (No Lone Pairs of

Electrons on A)

2s 2pC [He]


Tetrahedral Electronic Geometry: AB4 Species

Tetrahedral Electronic Geometry: AB3U Species

Valence Bond Theory (Hybridization)

2s 2p

N [He] four sp3 hybrids

2s 2p

C [He] four sp3 hybrids

2s 2pO [He]

four sp3 hybrids

Tetrahedral Electronic Geometry: AB2U2 Species


Tetrahedral Electronic Geometry: ABU3 Species (Three Lone Pairs of Electrons on

A)Valence Bond Theory (Hybridization)

2s 2pF [He]

four sp3 hybrids

tetrahedral

FH

··

····


Trigonal Bipyramidal Electronic Geometry: AB5, AB4U, AB3U2, and AB2U3

4s 4p 4dAs [Ar] 3d10

five sp3 d hybrids 4d


Compounds Containing Double Bonds

• Ethene or ethylene, C2H4, is the simplest organic compound containing a double bond.

Lewis dot formulaN = 2(8) + 4(2) = 24A = 2(4) + 4(1) = 12 S = 12

• Compound must have a double bond to obey octet rule.

CC

H

HH

H

C CH

H

H

H····

·· ·· ··

··or



• VSEPR Theory suggests that the C atoms are at center of trigonal planes.

C C

H

HH

H



Valence Bond Theory (Hybridization)

C atom has four electrons.Three electrons from each C atom are in sp2 hybrids.One electron in each C atom remains in an unhybridized p orbital

2s 2p three sp2 hybrids 2p

C • An sp2 hybridized C atom has this shape.

Remember there will be one electron in each of the three lobes.

Top view of an sp2 hybrid



• The single 2p orbital is perpendicular to the trigonal planar sp2 lobes.The fourth electron is in the p orbital.

Side view of sp2 hybrid with p orbital included.



• Two sp2 hybridized C atoms plus p orbitals in proper orientation to form C=C double bond.

• The portion of the double bond formed from the head-on overlap of the sp2 hybrids is designated as a bond.

• The other portion of the double bond, resulting from the side-on overlap of the p orbitals, is designated as a bond.


Compounds Containing Triple Bonds

• Ethyne or acetylene, C2H2, is the simplest triple bond containing organic compound.

Lewis Dot FormulaN = 2(8) + 2(2) = 20A = 2(4) + 2(1) =10S = 10

• Compound must have a triple bond to obey octet rule.



Lewis Dot Formula

C C HHCH HC·· ·· ···· ·· or

VSEPR Theory suggests regions of high electron density are 180o apart.

H C C H

H

C

HHH



Valence Bond Theory (Hybridization)Carbon has 4 electrons.Two of the electrons are in sp hybrids.Two electrons remain in unhybridized p orbitals.

2s 2p two sp hybrids 2p

C [He]


Compounds Containing Triple Bonds A bond results from the head-on overlap of two

sp hybrid orbitals.

Note that a triple bond consists of one and two p bonds.

The unhybridized p orbitals form two bonds.


Summary of Electronic & Molecular Geometries

© 2006 brooks/cole - thomson writing lewis formulas: the octet rule the octet rule states that...

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