Outline for Today Monday, Nov. 12 Chapter 8: Chemical Bonding Bond - - PowerPoint PPT Presentation

Outline for Today

Monday, Nov. 12

• Chapter 8: Chemical Bonding
• Bond Enthalpies
• Chapter 9: Theories of Bonding
• VSEPR (Valence Shell Electron Pair Repulsion) Theory
• Valence Bond
• Orbital Hybridization
• Molecular Orbital

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← ⃪ Wednesday ← ⃪ Friday

Exceptions to the Octet Rule

Odd number of electrons? Use resonance structures and formal charge to guide your decisions.

• 1. NO
• 2. NO2
• 3. Superoxide: O2-

Chapter 9 Spoiler Alert! Lewis Structures aren’t great at describing radicals! We’ll learn about a better model next week called Molecular Orbital Theory!

Procedure To Draw Lewis Structures

• 1. Count the total valence electrons. Include overall charge.
• 2. Write the atomic symbols. Connect with single bonds.
• A. The central atom is usually written first.
• B. Central atom is usually the least electronegative.
• 3. Complete octets around all atoms using lone pairs or multiple
• bonds. Make sure you don’t exceed total valence electrons!
• 4. Check formal charges. Can rearranging electrons lead to

formal charges closer to zero?

• 5. Consider breaking the octet rule for elements row 3 and

below if it leads to better formal charges.

• 6. Multiple possible electron configurations? Consider

resonance structures!

Bond Strengths and Bond Enthalpies

• Bond Enthalpy: The energy it takes to BREAK a bond.
• Related to bond strength and bond length.
• As the number of bonds between atoms increase, the

bond becomes shorter and stronger. It takes more energy to break it.

Bond Strengths and Bond Energies

Energy to Break a Bond

Using Bond Enthalpies to Estimate Enthalpy of a Reaction

∆Hrxn= 𝚻∆Hbonds broken — 𝚻∆Hbonds formed

Bond Bond Enthalpy (kJ/mol)

C—H 413 C—C 348 C—O 358 O—O 146 O—H 463 N—H 391 C—N 293

Bond Bond Enthalpy (kJ/mol)

C=C 614 O=O 495 C=O 799 N=N 418 C≡O 1072 C≡N 891 N≡N 941

Selected Values from Table 8.4 in your text

Example Problem: Bond Enthalpies

Use bond enthalpies to estimate the ∆H for the combustion reaction of CH4 (methane).

• 1. Balance the Reaction
• 2. Draw out Lewis Structures for all molecules
• 3. Look up Bond Enthalpies
• 4. Use ∆Hrxn= 𝚻∆Hbonds broken — 𝚻∆Hbonds formed

On your note card…

• 1. Your Name
• 2. On one side, draw a picture or diagram that is

important to your understanding of drawing Lewis structures.

• 3. On the other side, write a 2-3 sentence summary of

how to draw lewis structures.

Molecules are Three Dimensional Objects

How do we go from two dimensional Lewis structures to three dimensional shapes?

Chapter 9: Molecular Bonding Theories

• Guiding Questions about each bonding theory:
• How does it work?
• What assumptions are made?
• What are the strengths and weaknesses?
• Theories:
• 1. Valence Shell Electron Pair Repulsion Theory

(VSEPR)

• 2. Valence Bonding
• 3. Hybridized Orbitals
• 4. Molecular Orbital Theory

VSEPR Shapes are based on how many Electron Groups are around an Atom

2 Electron Groups 3 Electron Groups 4 Electron Groups

Linear Trigonal Planar Tetrahedral 180o 120o 109.5o

109.5o 107o 104.5o

Be sure to study Tables 9.2 and 9.3!

Putting it together: Molecular Geometry for Polyatomic Molecules

O C C O C N H H H H H H H

Alanine, an amino acid:

What is the electron domain geometry of each atom indicated with an arrow?

• 1. C
• 2. O
• 3. N
• 4. C

Putting it together: Molecular Geometry for Polyatomic Molecules

O C C O C N H H H H H H H

What is the molecular geometry of each atom indicated with an arrow?

Alanine, an amino acid:

• 1. C
• 2. O
• 3. N
• 4. C

Putting it together: Molecular Geometry for Polyatomic Molecules

O C C O C N H H H H H H H

What is the approximate atomic bond angles of each atom indicated with an arrow?

Alanine, an amino acid:

• 1. C
• 2. O
• 3. N
• 4. C

VSEPR Shapes are based on how many Electron Groups are around an Atom

5 Electron Groups 6 Electron Groups

Trigonal Bipyramidal Octahedral 90o, and 120o 90o

Dipoles of Three Dimensional Molecules

Valence Bond Theory

According to Valence Bond theory: Covalent bonds are a result of the

• verlap of atomic orbitals of the

valence electrons Strengths: Does an orbital overlap does a good job at explaining why bonds form between 2 atoms. Weaknesses: Does not work well for predicting 3D shapes of molecules with 3 or more atoms.