CH 9: Ionic and Covalent Bonding Renee Y. Becker Valencia Community College CHM 1045 1.

CH 9: Ionic and Covalent Bonding

Renee Y. Becker

Valencia Community College

CHM 1045

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Covalent Bonds

• Covalent bonds are formed by sharing at least one pair of electrons.

• The attraction (nucleus/electrons) outweighs the repulsions (electron/electron &

nucleus/nucleus)

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Covalent Bonds

•Every covalent bond has a characteristic

length that leads to maximum stability.

bond length3

Strength of Covalent Bonds

•Energy required to break a covalent bond in an

isolated gaseous molecule is called the bond

dissociation energy.

•Same amount of energy released when the

bond forms

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Example 1:

Which of the following is correct?

1. Energy is absorbed to form a bond

2. Energy is released when a bond is formed

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Polar Covalent Bonds

• Bond polarity is due to electronegativity differences between atoms.

• Pauling Electronegativity: is expressed on a scale where F = 4.0

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Pauling ElectronegativitiesPauling Electronegativities

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Electron-Dot Structures

• Using electron-dot (Lewis) structures, the

valence electrons in an element are

represented by dots.

• Lewis symbols

• Valence electrons are those electrons with the

highest principal quantum number (n). 9

.

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Electron-Dot Structures

• The electron-dot structures provide a simple, but useful, way of representing chemical reactions.

• Ionic:

• Covalent:11

Electron-Dot StructuresElectron-Dot Structures

• Single Bonds:

• Double Bonds:

• Triple Bonds:

C

HHHH

C HH

H

H

C

H

HH

H

C C C

C HC C CH

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Drawing Lewis-Dot Structures

Rule 1: Count the total valence electrons.

Rule 2: Draw the structure using single bonds.

Rule 3: Distribute the remaining electron pairs around the peripheral atoms.

Rule 4: Put remaining pairs on central atom.

Rule 5: Share lone pairs between bonded atoms to create multiple bonds.

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Drawing Lewis-Dot Structures

• NH2F Amino Fluoride: In this

molecule, nitrogen is the central atom.

• Rule 1: Number of electrons = 5 + (2 x 1) + 7 = 14 = 7 pairs

NH H

F

NH H

F

NH H

F

Rule 2 Rule 3 Rule 414

Drawing Lewis-Dot Structures

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Drawing Lewis-Dot Structures

• Polyatomic molecules with central atoms below the second row ten:

• In this compound there are 10 valence

electrons on bromine; this is called an

expanded octet. The extra pairs go into

unfilled d orbitals.16

Example 2: Drawing Lewis-Dot Structures

•Draw electron-dot structures for:

C3H8 H2O2 CO2 N2H4

CH5N C2H4 C2H2 Cl2CO

H3S+ HCO3–

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Resonance Structures

• How is the double bond formed in O3?

• The correct answer is that both are correct, but neither is correct by itself.

O O O

O O O

O O O

or

Or from this oxygen?

Move lone pair from this oxygen?

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Resonance Structures

• When multiple structures can be drawn, the actual

structure is an average of all possibilities.

• The average is called a resonance hybrid. A straight

double-headed arrow indicates resonance.

O O O O O O

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Resonance Structures

• The nitrate ion, NO3–, has three

equivalent oxygen atoms, and its

electronic structure is a resonance

hybrid of three electron-dot structures.

Draw them.

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Formal Charge

• Formal Charge: Determines the best resonance structure.

• We determine formal charge and estimate the

more accurate representation.

Formal Charge= # of Valence e-

# of bonding e-

2

# of nonbonding e-

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Example 3: Formal Charge

• Calculate the formal charge and

determine the most favorable of the

following electron dot structures:

SO2 NO3– NCO– N2O O3 CO3

2–

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Example 4:

What is the overall formal charge of the following structure?

1. -2

2. -3

3. -1

4. 0

P

O

O

OO

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Example 5: Ionic Radii of Ions

• Compare ionic radii

– Fe & Fe3+

– Cl & Cl-

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