Lewis Theory of Covalent Bondingprofkatz.com/courses/wp-content/uploads/2014/03/CH1710-CH9-Chemical-Bonding-II.pdfLewis theory implies that some combinations should be stable. (“octets”)

Lewis Theory of Covalent Bonding

Why Do Atoms Bond?A chemical bond forms when the potential energy of the bonded atoms is less than the potential energy of

the separate atoms.

To calculate this potential energy, you need to consider the following interactions:

nucleus–to–nucleus repulsionselectron–to–electron repulsionsnucleus–to–electron attractions

Covalent Bonds

Nonmetal atoms have relatively high ionization energies.

It is difficult to remove electrons from them.

When nonmetals bond, it is better in terms of potential energy for the atoms to share valence electrons rather than transferring electrons from one atom to another.

Covalent Bonds

Potential energy is lowest when the electrons are between the nuclei.

Shared electrons are attracted to both atoms.

Lewis Theory of Covalent Bonding

Atoms can achieve an octet of valence electrons by sharing valence electrons.

Shared electrons count for each atom’s octet.

Sharing of valence electrons is called covalent bonding.

Bonding pairs - electrons that are shared by atoms

Lone pairs (nonbonding pairs) - electrons that are not shared by atoms but belong to a particular atom

Single Covalent BondsWhen two atoms share one pair of electrons it is called

a single covalent bond (2 electrons)

F ••

••

••

• F •• •• • ••

F ••

••

••

••

•• F •• ••

F F

H • H • O ••

•

•

••

H H O •• •• ••

••

One atom may use more than one single bond to fulfill its octet.

Double Covalent Bond

When two atoms share two pairs of electrons the result is called a double covalent bond (four electrons)

O ••

•

•

•• O ••

• •

••

O ••

•• O ••

•• •• ••

⤵

O O

Triple Covalent Bond

When two atoms share three pairs of electrons the result is called a triple covalent bond (six electrons)

N ••

•

•

• N ••

• •

•

N •• ••

•• •• ••

N

⤵

N N

Covalent BondingModel vs. Reality

Lewis theory implies that some combinations should be stable. (“octets”)

Lewis theory allows us to predict the formulas of molecules of covalently bonded substances.

Hydrogen and the halogens are all diatomic molecular elements.

Oxygen generally forms either two single bonds or a double bond in its molecular compounds

Predictions of Molecular Formulas

Hydrogen is more stable when it is singly bonded to another atom.

+ H2

HCl+


Attractions between atoms are directional.

The shared electrons are most stable between the bonding atoms.

Covalently bonded compounds are found as individual molecules.

Compounds of nonmetals are made of individual molecule units.

Lewis theory predicts the melting and boiling points of molecular compounds should be

relatively low.

1) Melting or boiling involves breaking the attractions between the molecules, but not the bonds between the atoms.

2) The covalent bonds between atoms are strong, but the attractions between the molecules are generally weak


Intermolecular Attractions vs. Bonding

Lewis theory predicts that neither molecular solids nor liquids should conduct electricity.

There are no charged particles around to allow the material to conduct.


Molecular compounds do not conduct electricity in the solid or liquid state.

Molecular acids conduct electricity when

dissolved in water, due to them being ionized by the water.

Lewis theory predicts that the more electrons two atoms share, the stronger the bond should be.

In general, triple bonds are stronger than double bonds, and double bonds are stronger than single bonds.


Lewis theory predicts that the more electrons two atoms share, the shorter the bond should be.

In general, triple bonds are shorter than double bonds, and double bonds are shorter than single bonds.

Polar Covalent Bonding

The result is a polar covalent bond

1) bond polarity

2) the end with the larger electron density gets a partial negative charge

3) the end that is electron deficient gets a partial positive charge

Covalent bonding between unlike atoms results in unequal sharing of the electrons

1) One atom pulls the electrons in the bond closer to its side.

2) One end of a bond has a larger electron density than the other.

H F • • δ+ δ�

Bond Polarity

Bond Polarity

Most bonds have some degree of sharing and some degree of ion formation to them.

Covalent indicates an amount of transfer insufficient for the material to display the classic properties of ionic compounds.

Polar covalent indicates enough unequal sharing enough to produce a dipole in the bond.

Electronegativity

Increases across period (➡) Decreases down group(⬇)

1) fluorine is the most electronegative element

2) francium is the least electronegative element

3) noble gas atoms are not assigned values

4) opposite of atomic size trend

The ability of an atom to attract bonding electrons to itself

Electronegativity Scale

Electronegativity Difference (∆EN) and Bond Classification

The larger the difference in electronegativity, the more polar the bond.

∆EN Classification

0 pure covalent

0.1-0.4 nonpolar covalent

0.5-1.9 polar covalent

2.0- ionic

ENCl

= 3.0

3.0 − 3.0 = 0

Pure Covalent

ENCl

= 3.0

ENH = 2.1

3.0 – 2.1 = 0.9

Polar Covalent

ENCl

= 3.0

ENNa

= 0.9

3.0 – 0.9 = 2.1

Ionic

Electronegativity Difference (∆EN) and Bond Classification

Bond Dipole MomentsDipole moment, µ, is a measure of bond polarity

1) A dipole is a material with a + and − end

2) Directly proportional to the size of the partial charges Directly proportional to the distance between them

μ = (q)(r) measured in Debyes, D

Generally, the more electrons two atoms share and the larger the atoms are, the larger the dipole moment.

Representative Dipole Momentsin the Gas Phase

Molecule ∆EN Dipole Moment (D)

Cl2 0 0

ClF 1.0 0.88

HF 1.9 1.82

LiF LiF 6.33

Percent Ionic Character

The percent ionic character is the percentage of a bond’s measured dipole moment compared to what it would be if the electrons were completely transferred.

The percent ionic character indicates the degree to which the electron is transferred.

Properties of the Period 3 chlorides.

The Period 3 Halides

Properties of the Period 3 chlorides.

∆EN 2.1 1.8 1.5 1.2 0.9 0.5 0.0

ENCl

= 3.0

3.0 − 3.0 = 0

Pure Covalent

ENCl

= 3.0

3.0 − 3.0 = 0

Pure Covalent

ENCl

= 3.0

3.0 − 3.0 = 0

Pure Covalent

ENCl

= 3.0

3.0 − 3.0 = 0

Pure Covalent

ENCl

= 3.0

3.0 − 3.0 = 0

Pure Covalent

ENCl

= 3.0

3.0 − 3.0 = 0

Pure Covalent

ENCl

= 3.0

3.0 − 3.0 = 0

Pure Covalent

ENCl

= 3.0

3.0 − 3.0 = 0

Pure Covalent

ENCl

= 3.0

3.0 − 3.0 = 0

Pure Covalent

ENCl

= 3.0

3.0 − 3.0 = 0

Pure Covalent

ENCl

= 3.0

3.0 − 3.0 = 0

Pure Covalent

ENCl

= 3.0

3.0 − 3.0 = 0

Pure Covalent

ENCl

= 3.0

3.0 − 3.0 = 0

Pure Covalent

ENCl

= 3.0

3.0 − 3.0 = 0

Pure Covalent

ENCl

= 3.0

3.0 − 3.0 = 0

Pure Covalent

ENCl

= 3.0

ENNa

= 0.9

3.0 – 0.9 = 2.1

Ionic

ENCl = 3.0

ENNa = 0.9

3.0 – 0.9 = 2.1

Ionic EN

Cl = 3.0

ENNa

= 0.9

3.0 – 0.9 = 2.1

Ionic

ENCl

= 3.0

ENNa

= 0.9

3.0 – 0.9 = 2.1

Ionic EN

Cl = 3.0

ENNa

= 0.9

3.0 – 0.9 = 2.1

Ionic EN

Cl = 3.0

ENNa

= 0.9

3.0 – 0.9 = 2.1

Ionic

ENCl

= 3.0

ENNa

= 0.9

3.0 – 0.9 = 2.1

Ionic EN

Cl = 3.0

ENNa

= 0.9

3.0 – 0.9 = 2.1

Ionic

ENCl

= 3.0

ENNa

= 0.9

3.0 – 0.9 = 2.1

Ionic

ENCl

= 3.0

ENNa

= 0.9

3.0 – 0.9 = 2.1

Ionic

ENCl

= 3.0

ENNa

= 0.9

3.0 – 0.9 = 2.1

Ionic

ENCl

= 3.0

ENNa

= 0.9

3.0 – 0.9 = 2.1

Ionic

ENCl

= 3.0

ENNa

= 0.9

3.0 – 0.9 = 2.1

Ionic

ENCl

= 3.0

ENNa

= 0.9

3.0 – 0.9 = 2.1

Ionic

ENCl

= 3.0

ENNa

= 0.9

3.0 – 0.9 = 2.1

Ionic

ENCl

= 3.0

ENNa

= 0.9

3.0 – 0.9 = 2.1

Ionic

Weak“secondary” forces

Strong“secondary” forces

Ionic vs Covalent “Secondary” Forces

Lewis Structures of Molecules

Lewis Structures of Molecules

Predicting the distribution of valence electrons in a molecule

Understanding the bonding in many compounds

Predicting shapes of molecules

Predicting physical properties of molecules

Beware!!

Lewis Theory does not require that atoms have the same number of lone pair electrons they had before bonding.

Some atoms commonly violate the octet rule.

H, He, Li, Be, B, P, S.....

Lewis StructuresGenerally try to follow the common bonding patterns

C = 4 bonds & lone pairs N = 3 bonds & 1 lone pairO = 2 bonds & 2 lone pairs H and halogens = 1 bondBe = 2 bonds & 0 lone pairs B = 3 bonds & 0 lone pairs

Structures that result in bonding patterns different from the common may have formal charges.

B C N O F

Drawing Lewis structures of molecules

1. Write skeletal structure

2. Count valence electrons

3. Attach atoms together with pairs of electrons, and subtract from the total

4. Complete octets, outside-to-inside

5. Give extra electrons to the central atom

6. If central atom does not have octet, bring in electrons from outside atoms to share

CH4 C

HH H

H

C

C,H,H,H,H,

HH

HH

C HH

HH

C2H4

C HH

HH

C

HH H

H

C C,C,H,H,H,H,

C HH

HHCC

P2H4

4(1) + 2(5) = 14 valence electrons

14 - 10 = 4 electrons remaining


H P

H

P

H

H

H P

H

P

H

H

H P

H

P

H

H

HNO3 (HONO2)

H O N

O

O 1 + 3(6) + 5 = 24 valence electrons

H O N

O

O 24 - 8 = 16 electrons remaining

H O N

O

O 16 - 16 = 0 electrons remaining

H O N

O

O

NO2-

2(6) + 5 + 1 = 18 electrons


14 -14 = 0 electrons remaining

N OO

N OO

N OO

N OO

CO2

6 + 4 + 6 = 16 valence electrons



C OO

C OO

C OO

C OO

SeOF2

2(7) + 6 + 6 = 26 valence electrons



Se FF

O

Se FF

O

Se FF

O

H3PO4

3(1) + 4(6) + 5 = 32 electrons


14 -14 = 0 electrons remainingH O P

O

O H

O

H

H O P

O

O H

O

H

H O P

O

O H

O

H

Draw Lewis Structures of the Following

CO2

P2H4

H P

H

P

H

H O N

O

C OO

NO2−

HNO3

H O N

O

O

SO32-

6 + 3(6) + 2 = 26 electrons


20 -20 = 0 electrons remaining

O S

O

O

O S

O

O

O S

O

O

Draw Lewis Structures of the Following

SeOF2 H3PO4

Se FF

O

H O P

O

O H

O

H

O S

O

O

SO32−

Formal Charge

During bonding, atoms may end with more or fewer electrons than the valence electrons they brought in order to fulfill octets.

This results in atoms having a formal charge.

FC = (#valence e−) − (#nonbonding e−) − (½ #bonding e−)

Sum of all the formal charges in a molecule = 0

Sum of all the formal charges in an ion = ionic charge

Writing Lewis Formulas of Molecules

Assign formal charges to the atoms

(fc = valence e− − lone pair e− − ½ bonding e−)

H O N

O

OFor H, fc = 1-0-½(2) = 0

H O N

O

ON

For O, fc = 6-4-½(4) = 0

For N, fc = 5-0-½(8) = +1

For O, fc = 6-4-½(4) = 0

For O, fc = 6-6-½(2) = -1

Common Bonding Patterns

B C N O

C +

N +

O +

C −

N −

O −

B −

F

F +

− F

Practice - Assign formal charges

H P

H

P

H

H C OO

Se FF

O H O P

O

O H

O

H

N OO[ ]-

-1 0 0

0 +1 0

-1

-1

+1 O S

O

O[ ]2-

-1 +1 -1

-1

(fc = valence e− − lone pair e− − ½ bonding e−)

0 0 00 0

6-6-1

6-6-1

6-6-1

5-o-4 6-2-3

6-2-3

Resonance

Lewis theory localizes the electrons between the atoms that are bonding together.

Extensions of Lewis theory suggest that there is some degree of delocalization of the electrons – We call this concept resonance.

Delocalization of charge helps to stabilize the molecule.

Resonance Structures

When there is more than one Lewis structure for a molecule that differ only in the position of the electrons, they are called resonance structures.

The actual molecule is a combination of the resonance forms – a resonance hybrid.

The molecule does not resonate between the two forms, though we often draw it that way.

Rules of Resonance Structures

1. Resonance structures must have the same connectivity (only electron positions can change)

2. Resonance structures must have the same number of electrons

3. Second row elements have a maximum of eight electrons (third row can have expanded octet)

4. Formal charges must total the same

Drawing Resonance Structures

1. First draw Lewis structure that maximizes octets

2. Assign formal charges

3. Move electron pairs from atoms with (−) formal charge toward atoms with (+) formal charge

4. If (+) fc atom 2nd row, only move in electrons if you can move out electron pairs from multiple bond

H O N

O

O

-1

+1

H O N

O

O

-1

+1

H O N

O

O-1

+1

Evaluating Resonance Structures

Better structures have fewer formal charges

Better structures have smaller formal charges

Better structures have the negative formal charge on the more electronegative atom

Drawing Resonance Structures - NO2-





5. If (+) fc atom 3rd row or below, keep bringing in electron pairs to reduce the formal charge, even if get expanded octet

N OO[ ]-

-1 0 0

NO O

0 0 -1

[ ]-

Drawing Resonance Structures - SeOF2





5. If (+) fc atom 3rd row or below, keep bringing in electron pairs to reduce the formal charge, even if get expanded octet Se FF

O

0 0 0

0

Se FF

O

0 +1 0

-1

SeOF2

Identify Structures with Better or Equal Resonance Forms and Draw Them

O N

O

NO2−

Se FF

O

Se FF

O

O

NO

Drawing Resonance Structures - H3PO4






H O P

O

O H

O

H

-1

+1

H O P

O

O H

O

H

00

Drawing Resonance Structures - H2SO4






−1

+2

−1

Drawing Resonance Structures - SO32-






O S

O

O[ ]2-

-1 +1 -1

-1

O S

O

O[ ]2-

-1 0 -1

0

Identifying Structures with Better or Equal Resonance Structures

-1

+2

-1

H2SO4

H O S

O

O H

O

0

0

0H O S

O

O H

O

CO2

C OO

-1 +2 -1

C OO

0 0 0

H3PO4

H O P

O

O H

O

H

O S

O

O

-1

+1 H O P

O

O H

O

H

0

0

-1

+1

-1 -1

O S

O

O

0

0

-1 -1 SO32−

Identifying Structures with Better or Equal Resonance Structures

Resonance Forms of the Sulfite Ion

O S

O

OO S

O

O O S

O

O

Resonance Structures

When there is more than one Lewis structure for a molecule that differ only in the position of the electrons, they are called resonance structures

The actual molecule is a combination of the resonance forms – a resonance hybrid

The molecule does not resonate between the two forms, though we often draw it that way

Bond Energies and Their Uses

Bond Energies

Chemical reactions involve breaking bonds in reactant molecules and making new bonds to create the products.

The ΔH°reaction can be estimated by comparing the cost of breaking old bonds to the income from making new bonds.

The amount of energy it takes to break one mole of a bond in a compound is called the bond energy.

Measured in the gas stateBonds are broken homolytically – each atom gets ½

bonding electrons

Trends in Bond Energies

In general, the more electrons two atoms share, the stronger the covalent bond C≡C (837 kJ) > C=C (611 kJ) > C−C (347 kJ)C≡N (891 kJ) > C=N (615 kJ) > C−N (305 kJ)

In general, the shorter the covalent bond, the stronger the bond Br−F (237 kJ) > Br−Cl (218 kJ) > Br−Br (193 kJ)

In general, bonds get weaker down the column In general, bonds get stronger across the period

Using Bond Energies to Estimate ΔH°rxn

We often use average bond energies to estimate ΔHrxn

Bond breaking is endothermic, ΔH(breaking) = +

Bond making is exothermic, ΔH(making) = −

ΔHrxn = ∑ (ΔH(bonds broken)) + ∑ (ΔH(bonds formed))

Average Bond Energies

Estimate the enthalpy of the following reaction

H C

H

H

H

Cl Cl H C

H

Cl

H

H Cl+ +

ΔHrxn = ∑ (ΔH(bonds broken)) + ∑ (ΔH(bonds made))

Bond breaking1 mole C─H" +414 kJ1 mole Cl─Cl" +243 kJ" total" +657 kJ

Bond making1 mole C─Cl" −339 kJ1 mole Cl─H" −431 kJ" total" −770 kJ

ΔHrxn = (+657 kJ) + (−770 kJ) = −113 kJ

H2(g) + O2(g) ➜ H2O2(g)

Reaction involves breaking 1 mol H–H and 1 mol O=O and making 2 mol H–O and 1 mol O–O

Bonds broken (energy cost) (+436 kJ) + (+498 kJ) = +934 kJ

Bonds made (energy release) 2(−464 kJ) + (−142 kJ) = −1070. kJ

ΔHrxn = (+934 kJ) + (−1070. kJ) = −136 kJ

(Appendix ΔH°f = −136.3 kJ/mol)

Estimate the enthalpy of the following reactionHO O

H

Trends in Bond Lengths

In general, the more electrons two atoms share, the shorter the covalent bond C≡C (120 pm) < C=C (134 pm) < C−C (154 pm)C≡N (116 pm) < C=N (128 pm) < C−N (147 pm)

In general, as bonds get longer, they also get weaker

Generally bond length decreases from left to right across period

C−C (154 pm) > C−N (147 pm) > C−O (143 pm)

Generally bond length increases down the column

F−F (144 pm) > Cl−Cl (198 pm) > Br−Br (228 pm)

Average Bond Lengths

Metallic Bonds

Metallic BondsThe low ionization energy of metals

allows them to lose electrons easily.

Metal atoms release their valence electrons to be shared by all to atoms/ions in the metal. a “sea” of electrons

electrons delocalized throughout the metal structure

Bonding results from attraction of the cations for the delocalized electrons.

Metallic Bonding-Model vs. Reality

Because the electrons are delocalized, they are able to move through the metallic crystal.

Because electrical conductivity takes place when charged particles (such as electrons) are able to move through the structure, this model predicts metallic solids should conduct electricity well.

Metallic solids do conduct electricity well.


This theory implies heating will cause the metal ions to vibrate faster.

Heating will therefore make it

more difficult for the electrons to travel through the crystal.

This theory predicts the conductivity of a metal should decrease as its temperature increases.

As temperature increases, the electrical conductivity of metals decreases.

Resistivity vs Temperature


Small, light electrons moving through the metallic solid can transfer kinetic energy quicker than larger particles locked into position.

Metallic solids do conduct heat well.

Delocalized electrons will share a set of orbitals that belong to the entire metallic crystal.

Delocalized electrons on the surface can absorb the outside light and then emit it at the same frequency.

Metallic solids do reflect light.

The attractive forces that hold the metal structure together result from the attraction of the metal atom cores for the delocalized electrons.

This model implies the attractive forces should not break if positions of the atom cores shift.

Metallic solids are malleable and ductile.



The attractions of the core atoms for the delocalized electrons is strong because it involves full charges.

In order to melt or boil, attractions holding the metallic crystal together must be broken.

Metals generally have high melting points and boiling points.

Metallic Bonding-Model vs. RealityThe attractions of the atom cores

for the delocalized electrons will be stronger when there are more delocalized electrons and when the charge on the atom core is larger.

Melting points generally increase left-to-right across period

Na (97.72 ºC) Mg (650 ºC) Al (660.32 ºC)

Metal ions get larger as you traverse down a column. The attractions of the atom cores for the delocalized electrons will be stronger when the atom cores are smaller

Melting points generally decrease down column

Li (180.54 ºC) Na (97.72 ºC) K (63.38 ºC)

Melting Points of the Group 1A(1) and Group 2A(2) Elements

Documents

Lewis Theory of Covalent Bondingprofkatz.com/courses/wp-content/uploads/2014/03/CH1710-CH9-Chemical-Bonding-II.pdfLewis theory implies that some combinations should be stable. (“octets”)