Outline for Today Wednesday, Nov. 7 Chapter 8: Chemical Bonding - - PowerPoint PPT Presentation

Outline for Today

Wednesday, Nov. 7

• Chapter 8: Chemical Bonding
• Lattice Energy
• Covalent Bonding
• Lewis Structure
• Electronegativity, Bond Polarity, Dipole Moments
• Resonance Structures

1

Chapter 8: Basic Concepts in Chemical Bonding

• Chemical Bond: When atoms or ions are strongly

attracted.

• Ionic Bond: Electrostatic forces (opposite charges) attract.

Usually metal cations ionic bond to nonmetal anions.

• Covalent Bond: Sharing of electrons, usually between

nonmetals.

• Metallic Bond: Each metal atom is bonded to many
• neighbors. Electrons are relatively free to move in the metal.

2

Valence Electrons form Bonds between Atoms

• Lewis Electron Dot Symbols help us keep

track of valence electrons.

• Octet Rule: Atoms will tend to gain or lose

electron to become isoelectronic with the nearest noble gas.

3

Ionic Bonding and Lattice Energy

∆Hfo=-788 kJ/mol ∆Hrxn=+147 kJ/mol

4

Trends in Lattice Energy

Eelectrostatic = κQ1Q2 d

Li+: 0.9 Å Na+: 1.16 Å K+: 1.52 Å Rb+: 1.66 Å Cl-: 1.67 Å

5

Procedure To Draw Lewis Structures

• 1. Count the total valence electrons. Include the
• verall charge.
• 2. Write the atomic symbols. Connect with single

bonds.

• A. In polyatomic molecule the central atom is

usually written first.

• 3. Complete octets around all atoms.
• A. Place any leftover electrons on the central atom.
• B. Not enough electrons? Try multiple bonds.

6

Examples: Single bonds and Lone Pairs

• 1. H2
• 2. F2

3.CH4

• 4. Ammonia
• 5. Hydroxide
• 6. Water
• 7. Cl2
• 8. SH2
• 9. PH3
• 10. HCl
• 11. CF4

7

Examples: Double Bonds

• 1. CO2
• 2. N2
• 3. C2H2
• 4. N3-
• 5. CO

6.CS2 7.CH2O

8

Electronegativity and Bond Polarity

• Bond Polarity: Describes the extent of unequal sharing of

electrons in a bond.

• Nonpolar covalent: electrons are shared equally

between atoms

• Polar covalent: One atom attracts the electrons more

than the other.

• Electronegativity: Ability of an atom in a molecule to

attract electrons.

9

Electronegativity Trends in the Periodic Table

10

Bond Polarity

11

Dipole Moment: Separation

• f Charge

Anion Charge:

Q=-1

Cation Charge:

Q=+1

Distance Between Charges:

r

Dipole Moment:

12

Example: Bond Polarity

• 1. CO2
• 2. F2
• 3. C2F2
• 4. NH3
• 5. HF
• 6. CH2Cl2

13

Formal Charge

• Formal Charge: the charge each atom would

have if each bonded electron is shared equally. F .C.= (# Valence Electrons)-(# Assigned Electrons)

14

Examples: Double Bonds and Lone Pairs (Using Formal Charge)

• 1. HCN vs HNC
• 2. SCN
• 3. OCN
• 4. NO2-
• 5. ClO3-

15

Procedure To Draw Lewis Structures

• 1. Count the total valence electrons. Include the overall

charge.

• 2. Write the atomic symbols. Connect with single bonds.
• A. In polyatomic molecule the central atom is usually

written first.

• B. Central atom is usually the least electronegative.
• 3. Complete octets around all atoms.
• A. Place any leftover electrons on the central atom.
• B. Not enough electrons? Try multiple bonds.
• 4. Check formal charges. Can rearranging electrons lead to

formal charges closer to zero?

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On your group’s note card

• 1. Names of group members.
• 2. Write a 2-3 sentence summary of how to draw lewis

structures.

• 3. Draw 1 picture or diagram that is important to

understanding chapter 8.

Examples: Resonance Structures

• 1. Acetate ion
• 2. CO32-
• 3. O3
• 4. C6H6 (benzene)

18

Examples: Exceptions to the Octet Rule

• 1. XeF4
• 2. BH3
• 3. NO
• 4. PCl5
• 5. Phosphate ion
• 6. SO2
• 7. SF4
• 8. I3-

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Bond Strength and Bond Enthalpy

Bond Enthalpy: The enthalpy required to break a one mole

• f a particular bond in the gas state.

∆Hrxn= ∆Hbonds broken - ∆Hbonds formed

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