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Chemical Bonding I:The Covalent Bond

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Chemical Bonding

I. Types of bondsa) Ionicb) Covalent

II. Lewis Dot Structuresa) Octet Ruleb) Multiple Bondsc) Resonanced) Polyatomic Ionse) Formal Charge on Atoms

Chemical Bonding

III. Exceptions to Octet Rulea) Incomplete Octetb) Expanded Octetc) Odd # of Electrons

IV. ElectronegativityV. Polar vs. NonPolar Covalent BondsVI. Use of Lewis Dot Structures

a) Bond Strength/Lengthb) Estimate ΔH

Bonding Generalities

• Unlike Charges Attract• Electrons will Be in Pairs• Only Valence Electrons are Involved

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Two Main Types of Bonds

1. Ionic Bonds- Electrostatic Attraction Between a + and – Ion That Holds Together Ionic Compounds

2. Covalent Bond – Sharing of an Valence Electron Pair Between two Nonmetal Atoms to Form a Molecule

Lewis Dot Symbols

9.2

Li + F Li+ F -

The Ionic Bond

Li Li+ + e-

e- + F F -

F -Li+ + Li+ F -

Fig. 9.6

3

Periodic Trends in Lattice EnergyEnergy of electrostatic force is proportional to:

Cation charge X anion charge________________________Cation radius + anion radius

Proportional to Lattice energy,

lattHΔ

lattHΔ

Fig. 9.7

A covalent bond is a chemical bond in which two or more electrons are shared by two atoms.

Why should two atoms share electrons?

F F+

7e- 7e-

F F

8e- 8e-

F F

F F

Lewis structure of F2

lone pairslone pairs

lone pairslone pairs

single covalent bond

single covalent bond

9.4

4

Fig. 9.12 Octet Rule

• When Nonmetal Atoms Share Valence Electrons to Form a Covalent Bond-– It Will Be to Have the Same Number of

Valence Electrons as the Closest Noble Gas

» H 2 electrons» Everything Else 8 Electrons

Comparison of Ionic and Covalent Compounds

9.4

1. Draw skeletal structure of compound showing what atoms are bonded to each other.

2. Count total number of valence e-. Add 1 for each negative charge. Subtract 1 for each positive charge.

3. Connect surrounding atoms to central atom with single bonds. Add remaining electrons (2 at a time) such that surrounding atoms follow octet rule (H -2 electrons). Stop adding electrons once the number exceeds value calculated in step 2.

Writing Lewis Structures for Molecular Compounds

9.6

5

Writing Lewis Structures(cont)

4. Make sure that every atom satisfies octet rule and the total number of valence electrons in Lewis Structure is correct.

5. If a surrounding atom does not have an octet; move lone pair to bonding position from an adjacent atom to form a double or triple bond. 3 Lone Pairs2 Lone Pairs1 Lone

Pair0 Lone Pairs

1 bond2 bonds3 bonds4 bonds

Group VIIF,Cl, Br, I

Group VIO,S

Group VN, P

Group IVC

BondsLonePairsGroupBonds

#4##8#

−=−=

Double bond – two atoms share two pairs of electrons

O C O or O C O

8e- 8e-8e-double bonds double bonds

Triple bond – two atoms share three pairs of electrons

N N8e-8e-

N N

triple bondtriple bond

or

9.4

Exception to Generality of # Bonds and Lone Pairs for Atom

1. Resonance Structures2. Polyatomic Ions3. Exceptions to the Octet Rule

• Incomplete Octet• Expanded Octet• Odd # Electrons

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Write the Lewis structure of nitrogen trifluoride (NF3).

Step 1 – N is less electronegative than F, put N in center

F N F

F

Step 2 – Count valence electrons N - 5 and F - 7

5 + (3 x 7) = 26 valence electronsStep 3 – Draw single bonds between N and F atoms and complete

octets on N and F atoms.Step 4 - Check, are # of e- in structure equal to number of valence e- ?

3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons

9.6

Resonance Structures

1. RS are Imaginary; the Real Structure is Intermediate All the RS.

2. Only Electrons Move Between RS; Atoms Positions Never Change

3. At Least One Atom in RS Will Have a Non-zero Formal Charge and Won’t Follow Generality About # Bonds and Lone Pairs.

9.7

An atom’s formal charge is the difference between the number of valence electrons in an isolated atom and the number of electrons assigned to that atom in a Lewis structure.

formal charge on an atom in a Lewis structure

=12

total number of bonding electrons( )

total number of valence electrons in the free atom

-total number of nonbonding electrons

-

The sum of the formal charges of the atoms in a molecule or ion must equal the charge on the molecule or ion.

Write the Lewis structure of the carbonate ion (CO32-).

Step 1 – C is less electronegative than O, put C in center

O C O

O

Step 2 – Count valence electrons C - 4 (2s22p2) and O - 6 (2s22p4)-2 charge – 2e-

4 + (3 x 6) + 2 = 24 valence electrons

Step 3 – Draw single bonds between C and O atoms and completeoctet on C and O atoms.

Step 4 - Check, are # of e- in structure equal to number of valence e- ?3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons

9.6

Step 5 - Too many electrons, form double bond and re-check # of e-

2 single bonds (2x2) = 41 double bond = 4

8 lone pairs (8x2) = 16Total = 24

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O C O

O

- -O C O

O

-

-

OCO

O

-

- 9.8

What are the resonance structures of the carbonate (CO3

2-) ion?Check Lewis Dot Structures

1. Correct Number of Valence Electrons in Structure

2. Every Atom Follows the Octet Rule3. Every Atom Follows Generality About #

Bonds and Lone Pairs UNLESS it Follows One of the Exceptions

Exceptions to the Octet Rule(on Central Atom)

The Incomplete Octet

H HBeBe – 2e-

2H – 2x1e-

4e-BeH2

BH3

B – 3e-

3H – 3x1e-

6e-

9.9

Incomplete Octet !4 not 8 e-

B HH

H

Incomplete Octet!6 not 8 e-

Coordinate Covalent Bonds

• In a coordinate covalent bonds, both electrons in the bond come from one of the atoms the bond is between.

• Usually, coordinate covalent bond is between– An atom with incomplete octet and– An atom with a lone pair

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Exceptions to the Octet Rule(on Central Atom)

The Expanded Octet (central atom with principal quantum number n≥3

SF6

S – 6e-

6F – 42e-

48e-S

F

F

F

FF

F

6 single bonds (6x2) = 1218 lone pairs (18x2) = 36

Total = 48

9.9

Exceptions to the Octet Rule

Odd-Electron Molecules

N – 5e-

O – 6e-

11e-NO

* When there is an odd # of valence electrons;One atom will have 7 rather than 8 electrons

* Use the electronegativity of atoms to determinewhich atom has 7 rather than 8 electrons

Electronegativity is the ability of an atom to attract toward itself the electrons in a chemical bond.

Electronegativity - relative, F is highest

9.5 9.5

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Lewis Dot Structure for NO (cont)

• O is more electronegative than N• The atom with a lower electronegativity has

a lower attraction for the electrons in the bond and will be the atom with 7 rather than 8 electrons

N OFree radical – unpaired electron

H F FH

Polar covalent bond or polar bond is a covalent bond with between 2 nonmetal atoms of different electronegativities. The electron pair is closer to the atom with a higher electronegativity

electron richregionelectron poor

region e- riche- poor

δ+ δ-

9.5

NonPolar covalent bond or nonpolar bond is a covalent bond between atoms of the about

the same electronegativity; the electrons pair is equally shared between atoms

H-H

Bond Type Comparison

Bond Type

Ionic Polar Covalent

NonPolarCovalent

Type of Atoms Involved

Metal & Nonmetal

2 Nonmetal Atoms of Different EN

2 Nonmetal Atoms with Similar EN

Example NaCl ICl Cl2

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NonPolar Covalent

share e-

Polar Covalent

partial transfer of e-

or unequal sharing

Ionic

transfer e-

Increasing difference in electronegativity

Classification of bonds by difference in electronegativity

Electronegativity Difference Bond Type

0 NonPolar Covalent≥ 2 Ionic

0 < and <2 Polar Covalent

9.5

Classify the following bonds as ionic, polar covalent, or covalent: The bond in CsCl; the bond in H2S; andthe NN bond in H2NNH2.

Cs – 0.7 Cl – 3.0 3.0 – 0.7 = 2.3 Ionic

H – 2.1 S – 2.5 2.5 – 2.1 = 0.4 Polar Covalent

N – 3.0 N – 3.0 3.0 – 3.0 = 0 NonPolarCovalent

9.5

The enthalpy change required to break a particular bond in one mole of gaseous molecules is the bond energy.

Bond Energies

Single bond < Double bond < Triple bond

9.10

Bond TypeBond

Energy(kJ/mole)

C-C 347

C=C 620

C≡C 812

Bond Type

Bond Length(pm)

C-C 154

C=C 133

C≡C 120

C-N 143

C=N 138

C≡N 116

Lengths of Covalent Bonds

Bond LengthsTriple bond < Double Bond < Single Bond 9.4

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Bond energy reflects electronegativitydifference of bonded atoms

Table 9.2

Table 9.3

Bond length reflects atomic or ionic radius

Bond length decreases with increasing energy

Estimating ΔH for a Reaction from Bond Energies

1. Balance Chemical Equation2. Write Lewis Dot Structures for all

reactants and products.3. Calculate energy needed to break reactant

bonds and make product bonds.• Breaking bonds is endothermic (+)• Making bonds is exothermic (-)

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Use bond energies to calculate the enthalpy change for:H2 (g) + F2 (g) 2HF (g)

Type of bonds broken

Number of bonds broken

Bond energy (kJ/mol)

Energy change (kJ)

H H 1 436.4 436.4F F 1 156.9 156.9Type of

bonds formedNumber of

bonds formedBond energy

(kJ/mol)Energy

change (kJ)

H F 2 568.2 1136.4

ΔH0 = 436.4 + 156.9 – 2 x 568.2 = -543.1 kJ

9.10