Lecture 9. Chemical Bonding

CHEMICAL BONDING

GENERAL CHEMISTRY

LECTURE

9

Skills that you should learn:

A. Draw the Lewis Structures of molecules:

Organic molecules

Polyatomic Ions

C. Determine the Formal Charge atoms

Ionic compound

Covalent compound

D. Determine the Polarity of bonds

B. Draw the Resonance Structures

Concept Check

• What is meant by the term “chemical bond?”

• How do atoms bond with each other to form molecules?

Chemical Bonding

1. Metal with nonmetal:

electron transfer and ionic bonding

2. Nonmetal with nonmetal:

electron sharing and covalent bonding

3. Metal with metal:

electron pooling and metallic bonding

Types of Chemical Bonding

- attractive forces that hold atoms together in compounds.

The three models of chemical bonding

Lewis Electron-Dot Symbols

Example:

N : .

. . :

N . . .

. N : .

. : N . . .

The group number gives the number of valence electrons.

Depicts the element and its valence electrons

Nitrogen, N

The dots represents the number of valence electrons

Lewis electron-dot symbols for elements in Periods 2 and 3.

Chemical Bonding

When atoms bound, they lose, gain, or share electrons to attain a filled outer level of eight (or two) electrons.

Octet Rule

Representative elements usually attain stable noble gas electron configurations in most of their compounds

Ionic Bonding

Present in IONIC COMPOUNDS

Bonding between a METAL + NONMETAL

Metal loses an electron CATION (+ charge)

Nonmetal gains an electron ANION (- charge)

e.g. NaCl, KBr, MgS

Ionic Bonding

Three ways to represent the formation of Li+ and F- through

electron transfer.

Formation of Ionic Compounds

Example: Na2O

2 Na . :

O . . .. O : + 2 Na+ :

: : [ ]2-

1. Write the Lewis-dot structure of the following compounds:

Exercise 1

a. Li3N b. AlCl3 c. MgO

a. Li3N Li

:

N . N : + 3Li+ : :

: [ ]3-

Solution:

: 3

b. AlCl3 Al

:

Cl . .. Cl : + Al3+ :

: : [ ]- : : 3 3

.

c. MgO Mg :

O .. O : + Mg2+ :

: : [ ]2- : :

.

Covalent Bonding

Present in COVALENT COMPOUNDS

Bonding between a NONMETAL + NONMETAL

* Atoms share electrons (outer e-/valence e-)

e.g. H2O, O2, NH3, PCl5

Covalent bonds are formed when atoms share electrons.

4 electrons - double bond formed.

6 electrons - triple bond formed.

2 electrons shared - single covalent bond formed.

Covalent Bonding

• The OCTET rule

• Bonding (or shared) electrons

• Nonbonding (or unshared or lone pairs) of electrons.

Writing Lewis Structure

N = number of needed electrons – N usually has a value of 8 for representative elements. – N has a value of 2 for H atoms.

A = number of available electrons in valence shells

– Correspond to the valence electron. – A is equal to the periodic group number for each element.

S = number of shared electrons in bonds. A-S = number of unshared electrons, lone pairs.

N - A = S rule

Writing Lewis Structure: The Octet Rule

• For ions adjust the number of electrons available, A.

– Add one e- to A for each negative charge.

– Subtract one e- from A for each positive charge.

• The central atom in a molecule or polyatomic ion is determined by:

– The atom with the least number in its molecular formula

– The less electronegative element

– Capable of forming multiple bonds


-> 4 bonds

Example 2: Write the Lewis structure of hydrogen cyanide, HCN.

N = 2 (H) + 8 (C) + 8 (N) = 18

A = 1 (H) + 4 (C) + 5 (N) = 10

S = 8 A-S = 2

H N C

Lone pairs


-> 4 bonds

Example 3: Write the Lewis structure of carbon dioxide, CO2

N = 8 (C) + 2 x 8 (O) = 24

A = 4 (C) + 2 x 6 (O) = 16

S = 8 A-S = 8

O O C

Lone pairs


The formal charge is the hypothetical charge on an atom in a molecule or polyatomic ion.

The formal charge determines the correct Lewis structures.

The best Lewis structures have formal charges on the atoms that are zero or nearly zero.

Writing Lewis Structure: Formal Charge

Rules for Assigning Formal Charge

F.C= valence e- – (# of bonds + # of unshared e-)

• Molecules: The sum of the formal charges of all atoms is equal to zero.

• Polyatomic ion: The sum of the formal charges is equal to the ion’s charge.


-> 3 bonds

Example 4: Consider nitrosyl chloride, NOCl. Draw the most stable Lewis structure of the compound.

N = 8 (N) + 8 (O) + 8 (Cl) = 24

A = 5 (N) + 6 (0) + 7 (Cl) = 18_ S = 6 A-S = 12 lone pairs

Cl O N Cl O N


Cl 7-(2+4) = +1

N 5-(3+2) = 0

O 6-(1+6) = -1

Cl 7-(1+6) = 0

N 5-(3+2) = 0

O 6-(2+4) = 0

Cl O N Cl O N


+1 0 -1 0 0 0

1. The covalent compounds of Be. 2. The covalent compounds of the IIIA Group. 3. Species which contain an odd number of electrons. 4. Species in which the central element must have a

share of more than 8 valence electrons to accommodate all of the substituents.

5. Compounds of the d- and f-transition metals.

Writing Lewis Structure:

Limitations of the Octet Rule

The outer atoms attached to the central atom nearly always attain noble gas configurations.

The central atom does not have a noble gas configuration but may have fewer than 8 (exceptions 1, 2, & 3) or more than 8 (exceptions 4 & 5).

Writing Lewis Structure:

Limitations of the Octet Rule

I. Reduced Octet Ex: BBr3; BeCl2

II. Expanded Octet Ex: PCl5; SF6

III. Odd-electron molecules Ex: NO2

Exemptions to the Octet Rule

Br

N = 8 (B) + 3 x 8 (Br) = 32

A = 3 (B) + 3 x 7 (Br) = 24

S = 8

A-S = 18

Br

-> 6 electrons

*Reduced Octet

B Br

Example 5: Write the Lewis structure of boron tribromide, BBr3.


F

N = 8 (As) + 5 x 8 (F) = 48

A = 5 (As) + 5 x 7 (F) = 40

S = 8

A-S = 30

-> 10 electrons

*Expanded Octet

Example 6: Write the Lewis structure of AsF5.

F

As F

F F


1. Reduced Octet a. BBr3

Br = 7–(1+6)= 0 B = 3–(3+0)= 0

b. BeCl2

Cl = 7–(1+6)= 0 Be = 2–(2+0)= 0

2. Expanded Octet a. PCl5

b. SF6

Cl = 7–(1+6)= 0 P = 5–(5+0)= 0

F = 7–(1+6)= 0 S = 6–(6+0)= 0


Exercises

Draw the Lewis structure of the following molecules and calculate the formal charges of each atom:

XeF2

SF4

Quiz

1. Write the Lewis structure of the following ionic compounds 2pts each

a. SrF2

b. K2O

c. Ca3N2

a. ICl3

b. SbF4-

2. Draw the most stable Lewis structure and calculate the formal charge of each atom on: 4 pts each

Answer

1. Write the Lewis structure of the following salts:

a. SrF2

b. K2O

c. Ca3N2

(2 pts each)

F : Sr2+ :

: : [ ]- 2

O : 2K+ :

: : [ ]2-

N : 3Ca2+ :

: : [ ]3- 2

a) ICl3

N = 8 (I) + 3 x 8 (Cl) = 32

A = 7 (I) + 3 x 7 (Cl) = 28

S = 4

A-S = 28-6 =22

-> 6 electrons

Cl

Cl Cl I


Answer

0

0

0 0

I : 7- (3+4) = 0 Cl : 7 – ( 1+6) =0

a) SbF4-

N = 8 (Sb) + 4 x 8 (F) = 40

A = 5 (Sb) + 4 x 7 (F) + 1 = 34

S = 6

A-S = 34-8 =26

-> 8 electrons

F

F F

Sb


Answer

F -

F : 7- (1+6) = 0 Sb :5 – (4+2)= -1

0 0

0 0

-1

Exercises

Draw the Lewis structure of the following molecules and calculate the formal charges of each atom:

1. BiF52-

2. IBr4+

3. OCCl2

4. XeOF4

5. HONO2

Resonance

Example 8: Write the a Lewis structure of sulfur trioxide, SO3.

N = 8 (S) + 4 x 8 (O) = 32

A = 6 (S) + 3 x 6 (O) = 24

S = 8

A-S = 16 -> 4 bonds

O

O S O

O

O S O

O

O S O

*Three possible structures for SO3.

- more than one valid Lewis structure can be written for a particular molecule.

O

O S O

O

O S O

O

O S O

All of the bonds in SO3 are equivalent.

There are no single or double bonds in SO3.

Best representation of the Lewis structure of SO3:

Resonance

Exercises

Draw the Lewis structure of nitrate ion, NO3-

N

O

O

O

N

O

O

O

N

O

O

O

Atom # of bonds

C 4 Tetravelent

N 3 Trivalent

O 2 Divalent

H, X 1 Monovalent

Examples: 1. CH3CH2CH3

2. CH3CH2OH 3. CH3COCH3 4. CH3NH3

1. CH3CH2CH3

2. CH3CH2OH

Molecules with More than One Central Atom

Examples: 1. CH3CH2CH3

2. CH2CH2OH 3. CH3COCH3

4. CH3NH2

3. CH3COCH3

4. CH3NH2

Atom # of bonds

C 4 Tetravelent

N 3 Trivalent

O 2 Divalent

H, X 1 Monovalent

Molecules with More than One Central Atom

N = 8 x 4 (O) + 8 (S) = 40

A = 6 x 4 (O) + 6 (S) + 2 = 32

S = 8

A-S = 24

Example 1: Write the Lewis structure of (NH4)2SO4

N = 4 x 2 (H) + 8 (N) = 16

A = 4 x1 (H) + 5 (N) -1 = 8

S = 8

A-S = 0

NH4+ + SO4

2- (NH4)2SO4

H

H N H

O

O S O

H + O 2-

Lewis Structure of Polyatomic Ion

NH4+ + SO4

2- (NH4)2SO4

H

H N H

H +

O

O S O

O 2-

2


Example 2: Write the Lewis structure of ammonium chloride, NH4Cl

H

H N H

H +

Cl -


O

Draw the correct Lewis Structure of H3PO4

N = 3 x 2 (H) + 8 (P) + 4 x 8 (O) = 46

A = 3 x 1 (H) + 5 (P) + 4 x 6 (O) = 32

S = 14

A-S = 18

-> 7 bonds

O

P

O

O H

H

H H : 1 –(1+0)= 0 3 O : 6- (2+4) = 0 P : 5- (4+0) = 1 O : 6- (1+6) = -1

-1

+1 0

0

0

Concept Check

O

b) H3PO4

O

P

O

O H

H

H H : 1 –(1+0)= 0 3 O : 6- (2+4) = 0 P : 5- (5+0) = 0 O : 6- (2+4) = 0

* Expanded octet

0 0

0

0

0

• Covalent bonds in which the electrons are shared equally are designated as nonpolar covalent bonds.

– Nonpolar covalent bonds have a symmetrical charge distribution.

Examples:

1. H2

2. N2

Bond Polarity

• Covalent bonds in which the electrons are not shared equally are designated as polar covalent bonds

– Polar covalent bonds have an asymmetrical charge distribution

Example:

1.HX (X =F, Cl, Br, I)

bondpolar very 1.9 Difference

4.0 2.1 ativitiesElectroneg

F H

1.9

Bond Polarity

• Polar molecules have a separation of centers of negative and positive charge, an asymmetric charge distribution.

Bond Polarity

Density map of HF

Bond Polarity

• Polar molecules can be attracted by magnetic and electric fields.

Bond Polarity

• The dipole moment has the symbol . (= d x q; unit: debye)

• Polar molecules have dipole moment

• Molecules that have a small separation of charge have a small . Molecules that have a large separation of charge have a large .

• For example, HF and HI:

units Debye0.38 units Debye1.91

I- H F- H

--

Dipole Moment

Bond Order, Bond Length, Bond Energy

Review

1. How were covalent bonds formed?

2. When forming a bond, the atoms obey what rule?

3. Formation of one covalent bond involves how many electrons?

4. In a triple bond, how many electrons are being shared?

5. How will you determine which is the best lewis structure?

Take home quiz

Draw the best Lewis structure of the following molecules and calculate the formal charges of each atom:

a. ICl2-

b. O3

c. Na3PO4

d. CaCO3

e. Glycine, CH3CONH2

Lecture 9. Chemical Bonding

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