Chapters 1 & 2

MolecularRepresentations

General Chemistry ReviewElectronic Structure

and Bonding

• 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 1.1

• 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.

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

Worksheet 1: Bonds, Formulas, Structures & Shapeshttp://chemconnections.org/organic/chem226/226assign-12.html#Worksheets

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.

• For simple Lewis structures:1. Draw the individual atoms using dots to represent the

valence electrons.2. Put the atoms together so they share PAIRS of

electrons to make complete octets. • Take NH3, for example:

• Practice with SKILLBUILDER 1.3.

Covalent Bonding – Simple Lewis Structures

Question 1.2

• Select the correct Lewis structure for methyl fluoride (CH3F).

• A) B)

•• C) D)

Important Bond Numbers(Neutral Atoms / Normal electron distribution / Free electron pairs

not shown)H F ICl Brone bond

Otwo bonds

Nthree bonds

Cfour bonds

Question 1.3

• What is the best Lewis structure for formaldehyde (H2CO)?

• A) B)

• C) D)

Question 1.4

• 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.

Klein: 2.4, 2.9

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 [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.  

Formal Charge = # of valence e-s -[1/2(# of bonding e-s ) + # of non-bonding e-s ]

CH3

N

C H3

C H3

C H3

CH3

N

C H3

..

CH3

N

C H3

C H3

O

..

: :

:

CH3

O

O

: :

..

:

:::

:

H O

..

:

:

C H3

C H3

CH3

C

C H3

C H3

CH3

C :

CH3

C H3

C H3

O: :

..

:

H O H

H

:

CH3

O H

C H3

..

CH3

O

C H2

: :

..

:

CH3

O

..

:

:

CH3

O

C H3

H

:

C H2

= ?

= ?

Worksheet #1 & Worksheet #4http://chemconnections.org/organic/chem226/226assign-12.html#Worksheets

Question 1.5 • What is the formal charge of the carbon

atom in the Lewis structure? • A) -1• B) 0• C) +1• D) +2 C

Question 1.6 • What is the formal charge of the oxygen

atom in the Lewis structure? • A) -1• B) 0• C) +1• D) +2

Resonance

Klein: 2.7 – 2.12

Resonance

Eg. SO2

Bond order 1.5

Bond length > double bond; Bond length < single bond

Resonance is a very important intellectual concept that was introduced by Linus Pauling in 1928 to explain experimental observations.

TUTORIALhttp://www.nku.edu/~russellk/tutorial/reson/resonance.html

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.

• 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 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.

Rules of Resonance

N

H

H

C

O

H

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

• Example:These structures have formal charges.

Rules 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 Lewis

structure

less stable Lewis

structure

....

C O N OH

H

H

.. :..+ –

....

C O N OH

H

H

..:..

“Pushing” Electrons

•same atomic positions

•differ in electron positions only

more stable Lewis

structure

less stable Lewis

structure

....

C O N OH

H

H

.. :..+ –

....

C O N OH

H

H

..:..

“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 long bond

Reality: bonds are of equal length (128 pm)

Resonance Example

O O••O••••••

••••–+

•Ozone (O3)

–Lewis structure of ozone shows one double bond and one single bond

Resonance:

O O••O••••••

••••–+

O O••O••••••

••••–+

O OO••••••

••••– +

••

Resonance Example

•Ozone (O3)

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

O O••O••••••

••••–+

O OO••••••

••••– +

••

Resonance Example

Detailed Resonance Examples

Question 1.7

• Which resonance structure contributes more to the hybrid?

• A) B)

VSEPR

Klein: 1.10

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

Trigonal Planar Trigonal Planar

Trigonal Planar Bent

Tetrahedral Tetrahedral

Tetrahedral Trigonal Pyramidal

Tetrahedral Bent

Trigonal Bipyramidal Trigonal Bipyramidal

Trigonal Bipyramidal Seesaw

Trigonal Bipyramidal T-shape

Trigonal Bipyramidal Linear

Octahedral Octahedral

Octahedral Square Pyramidal

Octahedral Square Planar

Orbital Geometry

Molecular Geometry Bond Angle

0

0

1

0

1

2

0

1

2

3

0

1

2

# of lone pairsChem 226

Practice: SKILLBUILDER 1.8.

Molecular Geometry –Summary

Lewis Structures / VSEPR / Molecular Models

• Computer Generated Models

Ball and stick models of ammonia, water and methane.

http://chemconnections.org/COT/organic1/VSEPR/

Worksheet 1: Bonds, Formulas, Structures & Shapeshttp://chemconnections.org/organic/chem226/226assign-12.html#Worksheets

Covalent Bond PolarityMolecular Polarity

Dipole Moment

Klein: 1.11

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)

• Electrons tend to shift away from lower electronegativity atoms to higher electronegativity atoms.

• The greater the difference in electronegativity, the more polar the bond.

Polar Covalent Bonds

Practice: SKILLBUILDER 1.5

http://chemconnections.org/organic/Movies%20Org%20Flash/en-polarity.swf

Question 1.8

• 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 = m = e x d(e) : magnitude of the charge on the atom(d) : distance between the two charges

Bond Dipole & Dipole Moment

1 debye = 10-18 esu ∙ cm An electrostatic unit of charge (esu) is a unit of charge. One electron has a charge of 4.80 x 10-10 esu.

Question 1.9

• Which of the following bonds have the greatest dipole moment (m)?

• A) B)

• C) D)

Bond PolarityA 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 1.10

• 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

Question 1.11

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

• A) CH4

• B) NH3

• C) SiH4

• D) H2O

Molecular PolarityElectrostatic Potential Maps

Models that visually portray polarity and dipoles

Hydrogen Halides

When identical polar bonds point in opposite directions,the effects of their polaritiescancel, giving no net dipolemoment. When they do notpoint in opposite directions,there is a net effect and a netmolecular dipole moment,designated .

Molecular Polarity & Dipole Moment

The vector sum of the magnitude and the direction of the individual bond dipole determines the overall dipole moment of a molecule

Molecular Dipole Moment

An electrically charged rod attracts a streamof chloroform buthas no effect on a stream of carbontetrachloride.

• Consider the resultant dipole for CH3Cl• Dipole moment (μ) = charge (e) x

distance (d)e = 1.056 x 10-10 esu;d = 1.772 x 10-8 cm The amount of charge separation is less

than if it were complete charge separation (4.80 x 10-10 esu).

• μ = 1.87 x 10-18 esu ∙ cm Conversion to debye

• μ = 1.87 D

Molecular Polarity

Visual depictions of polarity:

Molecular PolarityElectrostatic Potential Maps

Ammonia and in the Ammonium Ion

Water

• Resultant Molecular Dipoles > 0• Solubility: Polar molecules that

dissolve or are dissolved in like molecules

Polarity & Physical PropertiesOzone and Water

0.1278 nm

• The Lotus flower• Water & dirt repellancy

Worksheet 4: Functions, Polarity, Formal Chargehttp://chemconnections.org/organic/chem226/226assign-12.html#Worksheets

• Generally likes dissolves like.• Polar compounds with other polar compounds:

If compounds that are mixed are capable of H-bonding and/or strong dipole–dipole interactions, then they should dissolve.

• Nonpolar compounds with other nonpolar compounds: If a compound has no strong intermolecular attraction

with itself, then there are no strong forces which would have to be overcome to have it dissolve with a similar compound.

Solubility

The “Lotus Effect” Biomimicry

http://bfi.org/biomimicry

• Lotus petals have non-polar wax on its micrometer-scale rough surface, resulting in water contact angles up to 170°

• See the middle image above.• When it rains, water dissolves any polar materials on the

surface and gravity takes care of any insoluble, non-polar “dirt” on the lotus surface, much like a snow ball.

Wax

The “Lotus Effect” Biomimicry

http://www.sciencemag.org/cgi/content/full/299/5611/1377/DC1

• Isotactic polypropylene (i-PP) melted between two glass slides and subsequent crystallization provided a smooth surface. Atomic force microscopy tests indicated that the surface had root mean square (rms) roughness of 10 nm.

• A) The water drop on the resulting surface had a contact angle of 104° ± 2

• B) the water drop on a superhydrophobic i-PP coating surface has a contact angle of 160°.

Science, 299, (2003), pp. 1377-1380, H. Yldrm Erbil, A. Levent Demirel, Yonca Avc, Olcay Mert