3. Formal Charge-Resonance Excellent

Chemical Bonding I:

Lewis Theory

2008, Prentice Hall

Chemistry: A Molecular Approach, 1st Ed.Nivaldo Tro

Roy KennedyMassachusetts Bay Community College

Wellesley Hills, MA

Tro, Chemistry: A Molecular Approach 2

Bonding Theories• explain how and why atoms attach together• explain why some combinations of atoms are stable

and others are notwhy is water H2O, not HO or H3O

• one of the simplest bonding theories was developed by G.N. Lewis and is called Lewis Theory

• Lewis Theory emphasizes valence electrons to explain bonding

• using Lewis Theory, we can draw models – called Lewis structures – that allow us to predict many properties of moleculesaka Electron Dot Structuressuch as molecular shape, size, polarity


Why Do Atoms Bond?• processes are spontaneous if they result in a system

with lower potential energy• chemical bonds form because they lower the potential

energy between the charged particles that compose atoms

• the potential energy between charged particles is directly proportional to the product of the charges

• the potential energy between charged particles is inversely proportional to the distance between the charges


Potential Energy Between Charged Particles

• 0 is a constant = 8.85 x 10-12 C2/J∙m

• for charges with the same sign, Epotential is + and the magnitude gets less positive as the particles get farther apart

• for charges with the opposite signs, Epotential is and the magnitude gets more negative as the particles get closer together

• remember: the more negative the potential energy, the more stable the system becomes

r

qq 21

0potential 4

1E


Potential Energy BetweenCharged Particles

The repulsion between like-charged particles increases as the particles get closer together. To bring them closer requires the addition of more energy.

The attraction between opposite-charged particles increases as the particles get closer together. Bringing them closer lowers the potential energy of the system.


Bonding• a chemical bond forms when the potential

energy of the bonded atoms is less than the potential energy of the separate atoms

• have to consider following interactions: nucleus-to-nucleus repulsionelectron-to-electron repulsionnucleus-to-electron attraction


Types of Bonds

Types of Atoms Type of BondBond

Characteristic

metals to nonmetals

Ionicelectronstransferred

nonmetals tononmetals

Covalentelectrons shared

metal tometal

Metallicelectronspooled

8

Types of Bonding


Ionic Bonds

• when metals bond to nonmetals, some electrons from the metal atoms are transferred to the nonmetal atomsmetals have low ionization energy, relatively easy to

remove an electron fromnonmetals have high electron affinities, relatively

good to add electrons to


Covalent Bonds• nonmetals have relatively high ionization energies, so it

is difficult to remove electrons from them• when nonmetals bond together, it is better in terms of

potential energy for the atoms to share valence electronspotential energy lowest when the electrons are between the

nuclei• shared electrons hold the atoms together by attracting

nuclei of both atoms


Determining the Number of Valence Electrons in an Atom

• the column number on the Periodic Table will tell you how many valence electrons a main group atom hasTransition Elements all have 2 valence electrons; Why?

1A/1 2A/2 3A/13 4A/14 5A/15 6A/16 7A/17 8A/10

Li Be B C N O F Ne

1 e-1 2 e-1 3 e-1 4 e-1 5 e-1 6 e-1 7 e-1 8 e-1


Lewis Symbols of Atoms• electron dot symbols• use symbol of element to represent nucleus and

inner electrons• use dots around the symbol to represent valence

electronspair first two electrons for the s orbitalput one electron on each open side for p electrons then pair rest of the p electrons

Li Be

B

C

N

O

F

Ne


Lewis Symbols of Ions• Cations have Lewis symbols without

valence electronsLost in the cation formation

• Anions have Lewis symbols with 8 valence electronsElectrons gained in the formation of the anion

Li• Li+1

F

1

F


What We Know

• the noble gases are the least reactive group of elements

• the alkali metals are the most reactive metals and their atoms almost always lose 1 electron when they react

• the halogens are the most reactive group of nonmetals and in a lot of reactions they gain 1 electron


Stable Electron ArrangementsAnd Ion Charge

• Metals form cations by losing enough electrons to get the same electron configuration as the previous noble gas

• Nonmetals form anions by gaining enough electrons to get the same electron configuration as the next noble gas

• The noble gas electron configuration must be very stable

Atom Atom’s Electron Config

Ion Ion’s Electron Config

Na [Ne]3s1 Na+1 [Ne]

Mg [Ne]3s2 Mg+2 [Ne]

Al [Ne]3s23p1 Al+3 [Ne]

O [He]2s22p4 O-2 [Ne]

F [He]2s22p5 F-1 [Ne]


Octet Rule• when atoms bond, they tend to gain, lose, or share electrons to

result in 8 valence electrons• ns2np6

noble gas configuration• many exceptions

H, Li, Be, B attain an electron configuration like He He = 2 valence electrons Li loses its one valence electron H shares or gains one electron

though it commonly loses its one electron to become H+ Be loses 2 electrons to become Be2+

though it commonly shares its two electrons in covalent bonds, resulting in 4 valence electrons

B loses 3 electrons to become B3+

though it commonly shares its three electrons in covalent bonds, resulting in 6 valence electrons

expanded octets for elements in Period 3 or below using empty valence d orbitals


Lewis Theory• the basis of Lewis Theory is that there are

certain electron arrangements in the atom that are more stableoctet rule

• bonding occurs so atoms attain a more stable electron configurationmore stable = lower potential energyno attempt to quantify the energy as the calculation

is extremely complex


Covalent Bonding:Bonding and Lone Pair Electrons

• Covalent bonding results when atoms share pairs of electrons to achieve an “octet”

• Electrons that are shared by atoms are called bonding pairs

• Electrons that are not shared by atoms but belong to a particular atom are called lone pairsaka nonbonding pairs

Lone PairsBonding Pairs O S O•• ••••••••

•• ••••••


Single Covalent Bonds• two atoms share a pair of electrons

2 electrons

• one atom may have more than one single bond

F••

••

•• • F•••••••

F••

••

•• ••

••F•••• HH O

•• ••••

••

H•H• O••

• •

••

F F


Double Covalent Bond

• two atoms sharing two pairs of electrons4 electrons

O••••O••

••••••

O••

• •

••O••

• •

••

O O······ ··


Triple Covalent Bond

• two atoms sharing 3 pairs of electrons6 electrons

N••

• •

•N••

• •

•

N•••••••••• N

N N····


Covalent BondingPredictions from Lewis Theory

• Lewis theory allows us to predict the formulas of molecules

• Lewis theory predicts that some combinations should be stable, while others should notbecause the stable combinations result in “octets”

• Lewis theory predicts in covalent bonding that the attractions between atoms are directional the shared electrons are most stable between the bonding atoms resulting in molecules rather than an array


Covalent BondingModel vs. Reality

• molecular compounds have low melting points and boiling pointsMP generally < 300°Cmolecular compounds are found in all 3 states at room

temperature• melting and boiling involve breaking the attractions

between the molecules, but not the bonds between the atoms the covalent bonds are strong the attractions between the molecules are generally weak the polarity of the covalent bonds influences the strength of

the intermolecular attractions


Intermolecular Attractions vs. Bonding


Ionic BondingModel vs. Reality

• some molecular solids are brittle and hard, but many are soft and waxy

• the kind and strength of the intermolecular attractions varies based on many factors

• the covalent bonds are not broken, however, the polarity of the bonds has influence on these attractive forces


Ionic BondingModel vs. Reality

• molecular compounds do not conduct electricity in the liquid state

• molecular acids conduct electricity when dissolved in water, but not in the solid state

• in molecular solids, there are no charged particles around to allow the material to conduct

• when dissolved in water, molecular acids are ionized, and have the ability to move through the structure and therefore conduct electricity


Bond Polarity• covalent bonding between unlike atoms results in

unequal sharing of the electronsone atom pulls the electrons in the bond closer to its

sideone end of the bond has larger electron density than the

other• the result is a polar covalent bond

bond polaritythe end with the larger electron density gets a partial

negative chargethe end that is electron deficient gets a partial positive

charge


HF

H F••+d -d

FH

EN 2.1 EN 4.0


Electronegativity• measure of the pull an atom has on bonding

electrons• increases across period (left to right) and• decreases down group (top to bottom)

fluorine is the most electronegative elementfrancium is the least electronegative element

• the larger the difference in electronegativity, the more polar the bondnegative end toward more electronegative atom


Electronegativity Scale

31

Electronegativity and Bond Polarity• If difference in electronegativity between bonded atoms

is 0, the bond is pure covalentequal sharing

• If difference in electronegativity between bonded atoms is 0.1 to 0.4, the bond is nonpolar covalent

• If difference in electronegativity between bonded atoms 0.5 to 1.9, the bond is polar covalent

• If difference in electronegativity between bonded atoms larger than or equal to 2.0, the bond is ionic

“100%”

0 0.4 2.0 4.0

4% 51%Percent Ionic Character

Electronegativity Difference


Bond Polarity

ENCl = 3.03.0 - 3.0 = 0

Pure Covalent

ENCl = 3.0ENH = 2.1

3.0 – 2.1 = 0.9Polar Covalent

ENCl = 3.0ENNa = 0.9

3.0 – 0.9 = 2.1Ionic



Bond Dipole Moments• the dipole moment is a quantitative way of describing the

polarity of a bonda dipole is a material with positively and negatively charged endsmeasured

• dipole moment, m, is a measure of bond polarity it is directly proportional to the size of the partial charges and

directly proportional to the distance between them m = (q)(r)not Coulomb’s Lawmeasured in Debyes, D

• the percent ionic character is the percentage of a bond’s measured dipole moment to what it would be if full ions


Dipole Moments


Water – a Polar Molecule

stream of water attracted to a charged glass rod

stream of hexane not attracted to a charged glass rod


Example 9.3(c) - Determine whether an N-O bond is ionic, covalent, or polar covalent.

• Determine the electronegativity of each elementN = 3.0; O = 3.5

• Subtract the electronegativities, large minus small(3.5) - (3.0) = 0.5

• If the difference is 2.0 or larger, then the bond is ionic; otherwise it’s covalent

difference (0.5) is less than 2.0, therefore covalent• If the difference is 0.5 to 1.9, then the bond is

polar covalent; otherwise it’s covalentdifference (0.5) is 0.5 to 1.9, therefore polar covalent


Lewis Structures of Molecules

• shows pattern of valence electron distribution in the molecule

• useful for understanding the bonding in many compounds

• allows us to predict shapes of molecules• allows us to predict properties of molecules and

how they will interact together


Lewis Structures• use common bonding patterns

C = 4 bonds & 0 lone pairs, N = 3 bonds & 1 lone pair, O= 2 bonds & 2 lone pairs, H and halogen = 1 bond, Be = 2 bonds & 0 lone pairs, B = 3 bonds & 0 lone pairs

often Lewis structures with line bonds have the lone pairs left off their presence is assumed from common bonding patterns

• structures which result in bonding patterns different from common have formal charges

B C N O F


Writing Lewis Structures of Molecules HNO3

1) Write skeletal structure H always terminal

in oxyacid, H outside attached to O’s

make least electronegative atom central N is central

2) Count valence electrons sum the valence electrons for each

atom add 1 electron for each − charge subtract 1 electron for each + charge

ONOH

O

N = 5H = 1O3 = 3∙6 = 18Total = 24 e-



3) Attach central atom to the surrounding atoms with pairs of electrons and subtract from the total

ONOH

O

———

ElectronsStart 24Used 8Left 16



4) Complete octets, outside-in H is already complete with 2

1 bond

and re-count electrons

:

::

——— ONOH

O

N = 5H = 1O3 = 3∙6 = 18Total = 24 e-





5) If all octets complete, give extra electrons to central atom.

elements with d orbitals can have more than 8 electrons Period 3 and below

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

follow common bonding patterns if possible

:

::

—— ONOH|

O


Practice - Lewis Structures

• CO2

• SeOF2

• NO2-1

• H3PO4

• SO3-2

• P2H4


Practice - Lewis Structures

• CO2

• SeOF2

• NO2-1

• H3PO4

• SO3-2

• P2H4

:O::C::O:

::

O P

O

O

O

HH

H

••

••

••

••

••

••

••

••

••

F Se

O

F

••

••

•• •

•••

••

••

••

••

••

O S

O

O

••

••

•• •

•••

••

••

••

••

••

O N O ••

••

••

••

••••

16 e-

26 e-

18 e-

26 e-

32 e-

14 e-H P P H

HH

•• ••


Formal Charge• during bonding, atoms may wind up with more

or less electrons in order to fulfill octets - this results in atoms having a formal charge

FC = valence e- - nonbonding e- - ½ bonding e-

left O FC = 6 - 4 - ½ (4) = 0

S FC = 6 - 2 - ½ (6) = +1

right O FC = 6 - 6 - ½ (2) = -1• sum of all the formal charges in a molecule = 0

in an ion, total equals the charge

•• •• ••••••••

••O S O••••


Writing Lewis Formulas of Molecules (cont’d)

7) Assign formal charges to the atoms a) formal charge = valence e- - lone pair e- - ½ bonding e-

b) follow the common bonding patterns

OSO

H

|

HOCCH

|||

OH

0 +1 -1

all 0


Common Bonding Patterns

B C N O

C+

N+

O+

C-

N-

O-

B-

F

F+

-F


Practice - Assign Formal Charges

• CO2

• SeOF2

• NO2-1

• H3PO4

• SO3-2

• P2H4

O P

O

O

O

HH

H

••

••

••

••

••

••

••

••

••

F Se

O

F

••

••

•• •

•••

••

••

••

••

••

O S

O

O

••

••

•• •

•••

••

••

••

••

••

O N O ••

••

••

••

••••H P P H

HH

•• ••


Practice - Assign Formal Charges

• CO2

• SeOF2

• NO2-1

• H3PO4

• SO3-2

• P2H4

O P

O

O

O

HH

H

••

••

••

••

••

••

••

••

••

F Se

O

F

••

••

•• •

•••

••

••

••

••

••

O S

O

O

••

••

•• •

•••

••

••

••

••

••

O N O ••

••

••

••

••••H P P H

HH

•• ••

all 0

-1

P = +1rest 0

S = +1Se = +1

-1

-1all 0

-1

-1-1


Resonance• 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 hybridit does not resonate between the two forms,

though we often draw it that way• look for multiple bonds or lone pairs

•••• •• ••••••••

•• ••O S O O S O•••••• ••••

••••

••••


Resonance


Ozone Layer


Rules of Resonance Structures• Resonance structures must have the same connectivity

only electron positions can change• Resonance structures must have the same number of

electrons• Second row elements have a maximum of 8 electrons

bonding and nonbonding third row can have expanded octet

• Formal charges must total same• Better structures have fewer formal charges• Better structures have smaller formal charges• Better structures have − formal charge on more

electronegative atom


O N

O

O·· ··

········

··

··

Drawing Resonance Structures1. draw first Lewis structure that

maximizes octets2. assign formal charges3. 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

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

-1

-1

+1

O N

O

O

·· ····

····

······

-1

-1 +1


Exceptions to the Octet Rule

• expanded octetselements with empty d orbitals can have more

than 8 electrons• odd number electron species e.g., NO

will have 1 unpaired electronfree-radicalvery reactive

• incomplete octetsB, Al


Drawing Resonance Structures1. draw first Lewis structure that

maximizes octets2. assign formal charges3. 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

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

O S

O

O

O

HH

·· ··

········

··

······

-1

-1

+2

O S

O

O

O

HH

··

······

··

······

0

0

0


Practice - Identify Structures with Better or Equal Resonance Forms and Draw Them

• CO2

• SeOF2

• NO2-1

• H3PO4

• SO3-2

• P2H4

O P

O

O

O

HH

H

••

••

••

••

••

••

••

••

••

F Se

O

F

••

••

•• •

•••

••

••

••

••

••

O S

O

O

••

••

•• •

•••

••

••

••

••

••

O N O ••

••

••

••

••••H P P H

HH

•• ••

all 0

-1

P = +1

S = +1Se = +1

-1

-1all 0

-1

-1-1


Practice - Identify Structures with Better or Equal Resonance Forms and Draw Them

• CO2

• SeOF2

• NO2-1

• H3PO4

• SO3-2

• P2H4

O P

O

O

O

HH

H

••

••

••

••

••

••

••

••

••

O P

O

O

O

HH

H

••

••

••

••

••

••

••

••

F Se

O

F

••

••

•• •

•••

••

••

••

••

••

F Se

O

F

••

•• •

•••

••

••

••

••

••O S

O

O

••

••

•• •

•••

••

••

••

••

••

O S

O

O

••

••

•• •

• O S

O

O

••

••

•• •

•

O S

O

O

••

••

•• •

•

••

••

••

••

••

••

••••

••

••

••

••

••

••

••

O N O ••

••

••

••

••••O N O •

•••

••

••

••••

H P P H

HH

•• ••

none

-1

-1

-1

+1

all 0

+1

all 0

-1

none

S = 0in allres. forms


Bond Energies• chemical reactions involve breaking bonds in reactant

molecules and making new bond to create the products• the DH°reaction can be calculated by comparing the cost

of breaking old bonds to the profit 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 in the gas statehomolytically – each atom gets ½ bonding electrons


Trends in Bond Energies• the more electrons two atoms share, the stronger

the covalent bondC≡C (837 kJ) > C=C (611 kJ) > C−C (347 kJ)C≡N (891 kJ) > C=N (615 kJ) > C−N (305 kJ)

• the shorter the covalent bond, the stronger the bondBr−F (237 kJ) > Br−Cl (218 kJ) > Br−Br (193 kJ)bonds get weaker down the column


Using Bond Energies to Estimate DH°rxn

• the actual bond energy depends on the surrounding atoms and other factors

• we often use average bond energies to estimate the DHrxn

works best when all reactants and products in gas state

• bond breaking is endothermic, DH(breaking) = +• bond making is exothermic, DH(making) = −

DHrxn = ∑ (DH(bonds broken)) + ∑ (DH(bonds formed))

63


Using Bond Energies to Estimate DH°rxn

DH°rxn NaCl

65

Estimate the Enthalpy of the Following Reaction

H H + O O H O O H


Estimate the Enthalpy of the Following Reaction

H2(g) + O2(g) ® H2O2(g)

reaction involves breaking 1mol 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

DHrxn = (+934 kJ) + (-1070. kJ) = -136 kJ

(Appendix DH°f = -136.3 kJ/mol)


Bond Lengths

• the distance between the nuclei of bonded atoms is called the bond length

• because the actual bond length depends on the other atoms around the bond we often use the average bond lengthaveraged for similar bonds from

many compounds


Trends in Bond Lengths• the more electrons two atoms share, the shorter the

covalent bondC≡C (120 pm) < C=C (134 pm) < C−C (154 pm)C≡N (116 pm) < C=N (128 pm) < C−N (147 pm)

• decreases from left to right across periodC−C (154 pm) > C−N (147 pm) > C−O (143 pm)

• increases down the columnF−F (144 pm) > Cl−Cl (198 pm) > Br−Br (228 pm)

• in general, as bonds get longer, they also get weaker


Bond Lengths

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3. Formal Charge-Resonance Excellent