Chapter 1

Electronic Structure and

Bonding

Acids and Bases

• 1 nm = 10 Å• An atom vs. a nucleus ~10,000 x larger

~ 0.1 nm

Nucleus =1/10,000of the atom

Anders Jöns Ångström(1814-1874)

1 Å = 10 picometers = 0.1 nanometers =10-4 microns = 10-8 centimeters

Question 1

• What is the electronic configuration ofcarbon?

• A) 1s2 2s2 2px2

• B) 1s2 2s2 2px1 2py

12pz0

• C) 1s2 2s2 2px12py

12pz1

• D) 1s2 1px1 1py

12s2

Electron ConfigurationsNoble Gases and The Rule of Eight

• When two nonmetals react to form acovalent bond: They share electrons toachieve a Noble gas electronconfiguration.

• When a nonmetal and a metal react toform an ionic compound: Valenceelectrons of the metal are lost and thenonmetal gains these electrons.

G.N. LewisPhoto Bancroft Library, University of California/LBNL Image Library

Notes from LewisNotes from Lewisʼ̓s notebook and his s notebook and his ““LewisLewis”” structure. structure.

Footnote:G.N. Lewis, despite his insight and contributionsto chemistry, was never awarded the Nobel prize.

http://chemconnections.org/organic/Movies%20Org%20Flash/LewisDotStructures.swf

• Ionic compounds are formed when electron(s) aretransferred.• Electrons go from less electronegative element to themore electronegative forming ionic bonds.

Ionic Compounds Covalent Compounds•Share electrons.•1 pair = 1 bond.•Octet rule (“duet” for hydrogen)•Lewis structures:

Notice the charges: In one case they balance, can you name the compound?In the other they do not, can you name the polyatomic ion?

More about “formal” charge to come.

Question 2

• Select the correct Lewis structure formethyl fluoride (CH3F).

• A) B)

•• C) D)

Important Bond Numbers(Neutral Atoms / Normal electron distribution)

H F ICl Brone bond

Otwo bonds

Nthree bonds

Cfour bonds

Question 3

• What is the correct Lewis structure offormaldehyde (H2CO)?

• A) B)

• C) D)

Question 4

• Which of the following contains a triplebond?

• A) SO2• B) HCN• C) C2H4• D) NH3

Formal Charge

Formal charge is the charge of an atom ina Lewis structure which has a differentthan normal distribution of electrons.

Important Bond Numbers(Neutral Atoms / Normal electron distribution)

H F ICl Brone bond

Otwo bonds

Nthree bonds

Cfour bonds

Important Bond Numbers(Neutral Atoms / Normal electron distribution)

Organic Chemistry

C H O N

# of Valence e- s

4

1

6

5

Total # of Bonds

(neutral atom)

4

1

2

3

Combinations of bonds

(neutral atom):

# of single bonds

4

2

1

1

2

0

3

1

0

# of double bonds

0

1

0

0

0

1

0

1

0

# of triple bonds

0

0

1

0

0

0

0

0

1

Total Bonds

4

4

4

1

2

2

3

3

3

# of Free Pairs of

electrons

0

0

0

0

2

2

1

1

1

Formal Charge

• Equals the number of valence electrons(Group Number of the free atom) minus [thenumber of unshared valence electrons in themolecule + 1/2 the number of sharedvalence electrons in the molecule].

• Moving/Adding/Subtracting atoms andelectrons.

Formal charge = number of valence electrons –(number of lone pair electrons +1/2 number of bonding electrons)

HNO3 Nitric Acid

Complete the following table. It summarizes the formal charge on a(“central”) atom for the most important species in organic chemistry.

Question 5

• What is the formal charge of the carbonatom in the Lewis structure?

• A) -1• B) 0• C) +1• D) +2 C

Question 6

• What is the formal charge of the oxygenatom in the Lewis structure?

• A) -1• B) 0• C) +1• D) +2

Resonance

Resonance

Eg. SO2

Bond order ≅ 1.5

Bond length > double bond; < single bond

Resonance is a very important intellectual concept thatwas introduced by Linus Pauling in 1928 to explainexperimental observations.

TUTORIAL

Resonance

•Two or more Lewis structures may belegitimately written for certain compounds(or ions) that have double bonds and/orfree pairs of non-bonded electrons•It is a mental exercise in “pushing” ormoving electrons.•Refer to Table 1.6

• Step 1:The atoms must stay in the same position. Atomconnectivity is the same in all resonancestructures. Only electrons move.

• NON-Example: The Lewis formulas below arenot resonance forms. A hydrogen atom haschanged position.

Rules of Resonance

N

H

H

C

O

H

N

H

C

OH

H

• Step 2:Each contributing structure must havethe same total number of electrons andthe same net charge.

• Example:All structures have 18 electrons and anet charge of 0.

Rules of ResonanceRules of Resonance

N

H

H

C

O

H

N

H

H

C

O

H

N

H

H

C

O

H

• Step 3:Calculate formal charges for each atomin each structure.

• Example:None of the atoms possess a formalcharge in this Lewis structure.

Rules of ResonanceRules of Resonance

N

H

H

C

O

H

• Step 4:Calculate formal charges for the secondand third structures.

• Example:These structures have formal charges.

Rules of ResonanceRules of Resonance

N

H

H

C

O

H

N

H

H

C

O

H

NOTE: They are less favorable Lewis structures.

•same atomic positions

•differ in electron positions

more stable more stable Lewis Lewis

structurestructure

less stable less stable Lewis Lewis

structurestructure

........

CC OO NN OOHH

HH

HH

.... ::....++ ––

........

CC OO NN OOHH

HH

HH

....::....

“Pushing” Electrons

•same atomic positions

•differ in electron positions only

more stable more stable Lewis Lewis

structurestructure

less stable less stable Lewis Lewis

structurestructure

........

CC OO NN OOHH

HH

HH

.... ::....++ ––

........

CC OO NN OOHH

HH

HH

....::....

“Pushing” Electrons Why use Resonance Structures?

•Delocalization of electrons and charges betweentwo or more atoms helps explain energeticstability and chemical reactivity.

•Electrons in a single Lewis structure areinsufficient to show electron delocalization.

•A composite of all resonance forms moreaccurately depicts electron distribution. (HYBRID)NOTE: Resonance forms are not always evenlyweighted. Some forms are better than others.

•Ozone (O3)–Lewis structure of ozone shows one double bond and one single bond

Expect: one short bond and one Expect: one short bond and one long bondlong bond

Reality: bonds are of equal length Reality: bonds are of equal length (128 pm)(128 pm)

Resonance Example

OO OO••••OO••••••••••••

••••••••––++

•Ozone (O3)–Lewis structure of ozone shows one double bond and one single bond

Resonance:Resonance:

OO OO••••OO••••••••••••

••••••••––++

OO OO••••OO••••••••••••

••••••••––++

OO OOOO••••••••••••

••••••••–– ++

••••

Resonance Example

•Ozone (O3)

–Electrostatic potentialmap shows both endcarbons are equivalentwith respect to negativecharge. Middle carbonis positive.

OO OO••••OO••••••••••••

••••••••––++

OO OOOO••••••••••••

••••••••–– ++

••••

Resonance Example Detailed ResonanceExamples

Question 7

• Which resonance structure contributesmore to the hybrid?

• A) B)

VSEPR ModelValence Shell Electron Pair Repulsion

VSEPR Model

The molecular structure of a given atom is determinedprincipally by minimizing electron pair (bonded &free)repulsions through maximizing separations.

Some examples of minimizing interactions.

Predicting a VSEPR Structure• 1. Draw Lewis structure.• 2. Put pairs as far apart as possible.• 3. Determine positions of atoms from the

way electron pairs are shared.• 4. Determine the name of molecular

structure from positions of the atoms.

Linear Linear LinearLinear

Trigonal Trigonal PlanarPlanar Trigonal Trigonal PlanarPlanar

Trigonal Trigonal Planar Planar BentBent

TetrahedralTetrahedral TetrahedralTetrahedral

TetrahedralTetrahedral Trigonal Trigonal PyramidalPyramidal

TetrahedralTetrahedral BentBent

Trigonal BipyramidalTrigonal Bipyramidal Trigonal BipyramidalTrigonal Bipyramidal

Trigonal BipyramidalTrigonal Bipyramidal SeesawSeesaw

Trigonal BipyramidalTrigonal Bipyramidal T-shapeT-shape

Trigonal BipyramidalTrigonal Bipyramidal LinearLinear

OctahedralOctahedral OctahedralOctahedral

OctahedralOctahedral Square PyramidalSquare Pyramidal

OctahedralOctahedral Square PlanarSquare Planar

Orbital Orbital GeometryGeometry

Molecular Molecular GeometryGeometry Bond AngleBond Angle

00

00

11

00

11

22

00

11

22

33

00

11

22

# of lone pairs# of lone pairsChem 226

Lewis Structures / VSEPR /Molecular Models

• Computer Generated Models

Ball and stick models of ammonia, water andmethane.

http://chemconnections.org/organic/chem226/Labs/VSEPR/

Worksheet 1: Bonds, Formulas, Structures & Shapeshttp://chemconnections.org/organic/chem226/226assign-09.html#Worksheets

Covalent Compounds•Equal sharing of electrons: nonpolar covalentbond, same electronegativity (e.g., H2)• Unequal sharing of electrons between atoms ofdifferent electronegativities: polar covalent bond(e.g., HF)

Question 8

• Which of the following bonds is the mostpolar?

•• A) B)•• C) D)

• Dipole moments are experimentally measured.• Polar bonds have dipole moments.

dipole moment (D) = µ = e x d(e) : magnitude of the charge on the atom(d) : distance between the two charges

Bond Dipole & Dipole Moment Question 9

• Which of the following bonds have thegreatest dipole moment (µ)?

• A) B)

• C) D)

Bond PolarityA molecule, such as HF, that has a centerof positive charge and a center of negativecharge is polar, and has a dipole moment.The partial charge is represented by δ andthe polarity with a vector arrow.

!+ !"

FH

Question 10

• In which of the compounds below is the δ+for H the greatest?

• A) CH4• B) NH3• C) SiH4• D) H2O

Question 11

• In which of the following is oxygen thepositive end of the bond dipole?

• A) O-F• B) O-N• C) O-S• D) O-H