Chapter 1 Electronic Structure and Bonding Acids and Bases

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 of carbon?

• A) 1s2 2s2 2px2

• B) 1s2 2s2 2px1 2py

12pz0

• C) 1s2 2s2 2px12py

12pz1

• D) 1s2 1px1 1py

12s2

Electron Configurations Noble Gases and The Rule of Eight

• When two nonmetals react to form a covalent bond: They share electrons to achieve a Noble gas electron configuration.

• When a nonmetal and a metal react to form an ionic compound: Valence electrons of the metal are lost and the nonmetal gains these electrons.

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

Notes from Lewis’s notebook and his “Lewis” structure.Notes from Lewis’s notebook and his “Lewis” structure.

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

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

• Ionic compounds are formed when electron(s) are transferred. • Electrons go from less electronegative element to the more 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 for methyl 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 of formaldehyde (H2CO)?

• A) B)

• C) D)

Question 4

• Which of the following contains a triple bond?

• A) SO2

• B) HCN

• C) C2H4

• D) NH3

Formal Charge

Formal charge is the charge of an atom in a Lewis structure which has a different than normal distribution of electrons.


H F ICl Brone bond

Otwo bonds

Nthree bonds

Cfour bonds


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 [the number of unshared valence electrons in the molecule + 1/2 the number of shared valence electrons in the molecule].

• Moving/Adding/Subtracting atoms and electrons.

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 carbon atom in the Lewis structure?

• A) -1

• B) 0

• C) +1

• D) +2 C

Question 6

• What is the formal charge of the oxygen atom 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 that was introduced by Linus Pauling in 1928 to explain experimental observations.

TUTORIAL

Resonance

•Two or more Lewis structures may be legitimately written for certain compounds (or ions) that have double bonds and/or free pairs of non-bonded electrons

•It is a mental exercise in “pushing” or moving electrons.

•Refer to Table 1.6

• Step 1: The atoms must stay in the same position. Atom connectivity is the same in all resonance structures. Only electrons move.

• NON-Example: The Lewis formulas below are not resonance forms. A hydrogen atom has changed position.

Rules of Resonance

N

H

H

C

O

H

N

H

C

OH

H

• Step 2: Each contributing structure must have the same total number of electrons and the same net charge.

• Example:All structures have 18 electrons and a net 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 atom in each structure.

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


N

H

H

C

O

H

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

• Example:These structures have formal charges.


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 between two or more atoms helps explain energetic stability and chemical reactivity.

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

•A composite of all resonance forms more accurately depicts electron distribution. (HYBRID) NOTE: Resonance forms are not always evenly weighted. 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 potential

map shows both end

carbons are equivalent

with respect to negative

charge. Middle carbon

is positive.

OO OO••••

OO••••

••••••••••••••••––++

OO OOOO••••

••••••••••••••••

–– ++

••••

Resonance Example

Detailed Resonance Examples

Question 7

• Which resonance structure contributes more to the hybrid?

• A) B)

VSEPR ModelValence Shell Electron Pair Repulsion

VSEPR Model

The molecular structure of a given atom is determined principally 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 PlanarTrigonal Planar Trigonal PlanarTrigonal Planar

Trigonal Planar Trigonal Planar BentBent

TetrahedralTetrahedral TetrahedralTetrahedral

TetrahedralTetrahedral Trigonal PyramidalTrigonal Pyramidal

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 pairs

Chem 226

Lewis Structures / VSEPR / Molecular Models

• Computer Generated Models

Ball and stick models of ammonia, water and methane.

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 covalent bond, same electronegativity (e.g., H2)• Unequal sharing of electrons between atoms of different electronegativities: polar covalent bond (e.g., HF)

Question 8

• Which of the following bonds is the most polar?

•

• A) B)

•

• C) D)

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

dipole moment (D) = = 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 the greatest dipole moment ()?

• A) B)

• C) D)

Bond Polarity

A molecule, such as HF, that has a center of positive charge and a center of negative charge is polar, and has a dipole moment. The partial charge is represented by and the polarity with a vector arrow.+ δ−FH

Question 10

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

• A) CH4

• B) NH3

• C) SiH4

• D) H2O

Question 11

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

• A) O-F

• B) O-N

• C) O-S

• D) O-H

Chapter 1 Electronic Structure and Bonding Acids and Bases

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Transcript of Chapter 1 Electronic Structure and Bonding Acids and Bases