CHAPTER 3 CHEMICAL BONDING

NUR FATHIN SUHANA BT AYOB

SMK SULTAN ISMAIL, JB

LEARNING OUTCOMES (ionic bonding)

1. Describe ionic (electrovalent) bonding such as NaCl

and Mg𝐶𝑙2

LEARNING OUTCOMES (metallic bonding)

1. Explain metallic bonding in terms of electron sea

model

LEARNING OUTCOMES (intermolecular

forces : van der Waals forces and hydrogen bonding)

1. Describe the hydrogen bonding and van der Waals

forces (permanent, temporary and induced dipole)

2. Deduce the effect of van der Waals between

molecules on the physical properties of substances

3. Deduce the effect of hydrogen bonding

(intermolecular and intramolecular)on the physical

properties of substances

TYPES OF CHEMICAL BONDING

1. Metal and non-metal :

Electron transfer and ionic bonding

Metal atom (low IE) loses its valence electrons, non metal (high negative EA) gains electrons

2. Non-metal with non-metal :

Electron sharing and covalent bonding

A shared electron pair is considered to be localized between the two atoms

3. Metal with metal :

Electron pooling and metallic bonding

Electron sea model

LEWIS STRUCTURE

Subtopic 4.1

LEWIS SYMBOLS

1. When atoms react to form chemical bonds, only the electrons in the outermost valence shells are involved

2. Valence shell electrons of an atom represented either by cross ( X ) or a dot (

3. It known as Lewis structures or electron-dot structures

4. E.g

• )

LEWIS SYMBOLS

RELATIONSHIP BETWEEN GROUP AND VALENCE ELECTRON

Group no. Example Electronic configuration

Lewis diagram

1 Sodium

2 Magnesium

13 Aluminium

14 Silicon

15 Phosphorus

16 Sulphur

17 Chlorine

18 Argon

HOW TO WRITE LEWIS SYMBOL

Eg. N (Z = 7)

Electron configuration : 1𝑠2 2𝑠2 2𝑝3 (valence

electron = 5)

1. Identify no. of valence electron

2. Place one dot at a time on the four side (top,

bottom, right, left)

3. Pair up the dots until all are used

KEEP IN MIND !

1. Lewis symbols do not show the electron

configuration of the valence electron

2. E.g C (Z =6)

Electron configuration : 1𝑠2 2𝑠2 2𝑝2

3. C has 4 unpaired dots because it form 4 bonds

KEEP IN MIND ! 4. Element in the same group :

Similar valence electron configuration

Similar Lewis dot symbols

5. E.g.

N gains three electron to form 𝑁3− (-3 charge)

N can form three covalent bonds

Cl gains one electron to form 𝐶𝑙− (-1 charge)

Cl can form one covalent bond

EXERCISE 1

Write Lewis dot symbols for the following atoms:

(a) K (b) Ca (c) Be

(d) Ga (e) O (f) Br

(g) N (h) I (i) As

(j) F (k) Mg (l) S

LEWIS SYMBOL

EXERCISE 2

14

Write Lewis dot symbols for the species according to

the following electronic configuration:

(a)P : 1s2 2s2 2p6 3s1

(b) Q : 1s2 2s2 2p6 3s2 3p6 4s2

(c) R : 1s2 2s2 2p6 3s2 3p4

(d) S : 1s2 2s2 2p3

(e) T : 1s22s22p1

(f) U : 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6

(g) V : 1s2 2s2 2p6 3s2 3p2

(h) W : 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5

LEWIS SYMBOL

OCTET RULE 1. An atom other than H tends to form bonds (by losing

or gaining or sharing electron) until it is surrounded by eight

valence electron

2. It can be achieved by ;

E.g : Transfer of electrons

Li (Z=3) + F (Z =9)

1𝑠2 2𝑠1 1𝑠2 2𝑠2 2𝑝5

1𝑠2 1𝑠2 2𝑠2 2𝑝6

𝐿𝑖 + 𝐹 −

OCTET RULE

E.g. sharing of electron

ELECTRON CONFIGURATION OF IONS

Noble gas configuration ( 8 valence electron)

Eg. 1𝑠2 2𝑠2 2𝑝6 (Ne)

Pseudonoble gas configuration

eg. [Kr] 4𝑑10

Half-filled orbitals

e.g 3𝑑5

Form stable ions (duplet/octet)

STABILITIES OF IONS

18

Noble Gas

Configuration

Valence electronic

configuration: ns2np6

Example:

1) Na → Na+ + e-

1s22s22p63s1 _____________

2) F + e- → F-

1s22s22p5 _____________

NOBLE GAS CONFIGURATION

1. Atoms may lose or gain enough electron so as to forms

stable ion with octet (or duplet) configuration (𝑛𝑠2 𝑛𝑝6)

2. Eg.

𝑁𝑎 → 𝑁𝑎+ + 𝑒−

+ 𝑒− →

1𝑠2 2𝑠2 2𝑝6 3𝑠1 1𝑠2 2𝑠2 2𝑝6 = [Ne]

1𝑠2 2𝑠2 2𝑝6 3𝑠2 3𝑝5 1𝑠2 2𝑠2 2𝑝6 3𝑠2 3𝑝6 = [Ar]

A completely filled orbital but

not the noble gas configuration

STABILITIES OF IONS

20

Valence electronic configuration:

ns2np6nd10 or ns2np6nd10nf14

Example:

1) 31Ga : 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p1

Ga3+: ______________________

2) 29Cu : 1s2 2s2 2p6 3s2 3p6 3d10 4s1

Cu+: 1____________________

Pseudo

Noble Gas

Configuration

PSEUDO-NOBLE GAS CONFIGURATION

1. The (n-1) 𝑑10 configuration of a p-block metal atom that

empties its outer level

2. Eg.

𝑆𝑛4+ + 𝑒−

[Kr] 4𝑑10 5𝑠2 5𝑝2 [Kr] 4𝑑10

Zn : 1s22s22p63s23p64s23d10

Zn2+ : 1s22s22p63s23p63d10

(pseudonoble gas configuration )

A special stability of half-filled d orbital

STABILITIES OF IONS

22

Valence electronic

configuration: nd5

Example:

1) 25Mn : 1s2 2s2 2p6 3s2 3p6 3d5 4s2

Mn2+: ____________________

2)26Fe : 1s2 2s2 2p6 3s2 3p6 3d6 4s2

Fe3+: ____________________

Half-filled

Orbital

Configuration

HALF-FILLED ORBITALS

1. Some transition metal atoms form cations that have

electron configuration associated with half-filled d orbital

(𝑑5)

2. E.g. Mn → 𝑀𝑛2+ 2𝑒−

[Ar] 3𝑑5 4𝑠2 [Ar] 3𝑑5

Fe : 1s22s22p63s23p64s23d6

Fe3+ : 1s22s22p63s23p63d5

(stability of half-filled 3d orbital)

IONIC BONDING

Electron transfer theory

Strength of ionic bond

Properties of ionic compound

IONIC BONDING

1. Attractive electrostatic force between positive

and negative ions

2. Sometimes called : electrovalent bond

3. Ionic bonds are usually formed between metallic

elements (Group 1, 2 and 13) and non-metallic

elements (Group 15,16 and 17)

IONIC BONDING

Metal atom (more electropositive)

Non-metal atom (more electronegative)

FORMATION OF IONIC BONDS 1. By transferring electrons

2. Total number of electron lost by metal atoms = total number of electron gained by the non-metal atoms

3. E.g.

******************************************

𝐿𝑖 → 𝐿𝑖+ + 𝑒−

HOW TO SHOW ELECTRON TRANSFER

1. Three ways :

a. Electron configurations

Eg. 𝐿𝑖 + 𝐹

1𝑠2 2𝑠1 1𝑠2 2𝑠2 2𝑝5

𝐿𝑖+ 𝐹−

1𝑠2 1𝑠2 2𝑠2 2𝑝6

HOW TO SHOW ELECTRON TRANSFER

1. Three ways :

b. Orbital diagram

Eg. 𝐿𝑖 + 𝐹

𝐿𝑖+ 𝐹−

HOW TO SHOW ELECTRON TRANSFER

1. Three ways :

b. Lewis e-dot symnbol

Eg. 𝐿𝑖𝐹

+ F [Li]+ F

- [ ]

Exercises 3:

By using Lewis structure, show how the

ionic bond is formed in the compounds

below.

( a ) KF

( b ) BaO

( c ) Na2O

EXERCISE 4

32

The element X has one electron and the element Y

has six electrons in their outermost shell

respectively.

(a) What is the formula of the compound

formed between the elements X and Y?

(b) Name the type of bond formed in (a)

(c) Draw the Lewis structure to show the

formation of compound in (a)

LEWIS SYMBOL

Ionic bond is very strong, therefore ionic

compounds:

1. Have very high melting and boiling

points

2. Hard and brittle

3. Can conduct electricity when they are in

molten form or aqueous solution because

of the mobile ions

LEWIS STRUCTURE 1. Two dimensional structural formula consists of e-dot

symbols that show each atom and its neighbors, the

bonding pairs and the lone pairs that fill each atom’s

outer shell

2. E.g

WRITING LEWIS STRUCTURE 1. Step 1

Draw skeletal structure of compound showing what

atoms are bonded to each other

Put at least negative electron negative atom in the

center F

E.g

𝐹 − 𝑁 − 𝐹

Electronegativity

N = 3.0

F = 4.0

N = Central atom

F = Surrounding

atom

WRITING LEWIS STRUCTURE

2. Step 2 :

Count total no. of valence electrons

E.g. N𝐹3 𝐹

𝐹 − 𝑁 − 𝐹

Atom No. of valence electron

N X 1 5𝑒− x 1 = 5𝑒−

F X 3 7𝑒− x 3 = 21𝑒−

Total 26 𝑒−

WRITING LEWIS STRUCTURE 2. Step 2 :

For polynomials ions, add one electron for negative

charge of the ion, or subtract one electron for each

positive charge

E.g N𝐻4+

Atom No. of valence electron

N X 1 5𝑒− x 1 = 5𝑒−

H X 4 1𝑒− x 4 = 4𝑒−

+ 1 charge 1𝑒−

Total 8 𝑒−

WRITING LEWIS STRUCTURE 3. Step 3 :

complete an octet (8 𝑒−) for all atoms except H (2 𝑒− )

Complete the surrounding atoms first

electrons not involved in bonding shown as lone pairs

E.g

Check :

8 𝑒− x 3 = 24 𝑒−

+ 2 𝑒−

26 𝑒−

CH4

1 Determine central atom & count valence e-

Central atom :

2 Draw single bond & calculate the non-bonding e- (NBe-)

3 Complete the octet of the terminal atom

4 Place any remaining e- at the central atom

5 Form double or triple bond if octet rule is not satisfied

DRAWING LEWIS STRUCTURE

39

NO2 +

1 Determine central atom & count valence e-

Central atom :

2 Draw single bond & calculate the non-bonding e- (NBe-)

3 Complete the octet of the terminal atom

4 Place any remaining e- at the central atom

-

5 Form double or triple bond if octet rule is not satisfied

40 1 2

DRAWING LEWIS STRUCTURE

WRITING LEWIS STRUCTURE 4. Step 4 :

If a central atom does have an octet, make a multiple

bond by changing a lone pair from one of the surrounding

atoms into a bonding pair to the central atom.

E.g

𝑁 = 𝑁 𝑁 ≡ 𝑁

STRENGTH OF IONIC BONDS 1. The strength of an ionic bond is a measure of the

electrostatic attraction between the ions

2. 𝐹 ∝𝑄+𝑄−

𝑑2

3. The smaller the ions and/or the higher charge on ions

> the stronger attraction between ions > the stronger

the ionic bond

4. E.g

𝑄+ = charge + ve ion 𝑄− = charge – ve ion d = distance between the ions F = force of attraction

Compound NaCl NaBr

Melting point/℃ 801 750

STRENGTH OF IONIC BONDS 4. The melting point of sodium chloride is higher than that

of sodium bromide. This shows that the ionic bond in

NaCl is stronger than that in NaBr

5. This is because the 𝐶𝑙− ion is smaller than that of 𝐵𝑟−

ion

6. Electrostatic attraction between 𝑵𝒂+ and 𝑪𝒍− is

stronger.

Ion 𝑵𝒂+ 𝑪𝒍− 𝐵𝑟−

Ionic radius/nm 0.095 0.181 0.195

STRENGTH OF IONIC BONDS 1. The melting point of sodium chloride and magnesium

chloride are :

2. The 𝑀𝑔2+ ion is smaller in size than in 𝑁𝑎+ ion. On top of

that, 𝑀𝑔2+ has higher charge

3. As a result, the ionic bond in Mg𝑪𝒍𝟐 is stronger than that in

NaCl. This accounts for the higher melting of Mg𝑪𝒍𝟐.

Compound NaCl Mg𝑪𝒍𝟐

Melting point/℃ 801 987

Cation radius/nm 0.095 0.065

BOND LENGTH 2. For a given pair of atoms,

Bond length : single > double > triple

E.g.

𝐶 − 𝐶 > 𝐶 = 𝐶 > 𝐶 ≡ 𝐶 (154 pm) (134 pm) (121 pm)

Bond order increase

Shorter bond

Stronger bond

* Lebih panjang > lebih mudah break

As the number of bonds between the carbon increase, the bond length decreases because C are held more closely and tightly together

As the number of bonds between two atoms increases, the bond grows shorter and stronger

LEARNING OUTCOMES (covalent bonding)

1. Draw the Lewis structure of covalent molecules (octet rule such as

N𝐻3, 𝐶𝐶𝑙4, 𝐻2O, 𝐶𝑂2, 𝑁2𝑂4, and exception to the octet rule such

as 𝐵𝐹3, NO, 𝑁𝑂2, 𝑃𝐶𝑙5, 𝑆𝐹6)

2. Explain the concept of overlapping and hybridisation of the s and

p orbitals such as 𝐵𝑒𝐶𝑙2, 𝐵𝐹3, 𝐶𝐻4, 𝑁2, HCN, 𝑁𝐻3, 𝐻2O

molecules

3. Predict and explain the shapes of and bond angles in molecules

and ions using the principle of valence valence shell electron pair

repulsion, e.g. linear, trigonal planar, tetrahedral, trigonal

bipyramid, octahedral, v-shaped, seesaw and pyramidal

4. Explain the existence of polar and non-polar bonds (including C-

Cl, C-N, C-O, C-Mg) resulting in polar or/and non-polar molecules

LEARNING OUTCOMES (covalent bonding)

5. Relate bond lengths and bond strengths with respect to single,

double and triple bonds

6. Explain the inertness of nitrogen molecule in terms of its strong triple

bond and nonpolarity

7. Describe typical properties associated with ionic and covalent

bonding in terms of bond strength, melting point and electrical

conductivity

8. Explain the existence of covalent character in ionic compounds such

as 𝐴𝑙2𝑂3, 𝐴𝑙𝑙3, and Lil

9. Explain the existence of coordinate (dative covalent) bonding such as

𝐻3𝑂+, N𝐻4+

, 𝐴𝑙2𝐶𝑙6 , and [Fe (𝐶𝑁)6]³ˉ

COVALENT BOND

1. Covalent bond is force of attraction between two adjacent

nuclei and the electrons that are shared together between

them

2. The covalent bond is usually formed between non-metallic

elements

3. There are some exceptions. For example, beryllium and

aluminium are metals, but they form covalent bonds with

chlorine. E.g. 𝐵𝑒𝐶𝑙2, 𝐴𝑙𝐶𝑙3

COVALENT BOND

Bonding pair electron

Lone pair

electron

Covalent compounds:

Compounds may have these covalent bonds:

i. Single bond

ii. Double bond

iii. Triple bond.

8e-

H H O + + O H H

O H H

or

2e- 2e-

Lewis structure of water

single covalent bonds

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 N

8e- 8e-

N N or

triple bond triple bond

RESONANCE STRUCTURE

1. Two or more Lewis structure for a single molecule that

cannot be represented accurately by only one Lewis

structure

2. E.g. Ozone (𝑂3) 6 𝑒− X 3 = 18𝑒−

FORMAL CHARGE

1. Difference between the valence electron in an isolated

atom and the number of electron assigned to that atom

in a Lewis structure

2. Formal charge of atom :

No. of valence electron – [ No. of lone pair electron + half of bonding

electron]

No. of valence electron – [ No. of lone pair electron + No. of bonds]

FORMAL

CHARGE

(FC)

Is used to find the most

stable Lewis structure

The sum of the

FC of the atoms

must equal the

charge on the

molecule or ion

FC should be as

small as possible

Negative FC --

on more

electronegative

atom

Positive FC --

on more

electropositive

atom

FORMAL CHARGE

56

FORMAL CHARGE

1. Formal charge : O of O−𝑂 = 6 − 6 − 1 = −1

: O of O = 𝑂 = 6 − 4 − 2 = 0

: middle O = [6 – 2 – 3 ] = +1

0

+ 1

- 1 - 1

+ 1

0

SELECTING THE BEST RESONANCE STRUCTURE

1. Select the structure with :

All zero formal charge

Small formal charge

Negative formal charges are placed on the more

electronegative atoms

EXAMPLE

59

Calculate the formal charge for each atom of the

following compounds:

(a) (b)

(c) (d)

(e) (f)

H

I

H – C – H

I

H

H – N – H

I

H

H – O – H O

O O

O N O

O

-

S

O O

1) Draw all the possible Lewis structure

of COCl2.

2) Predict the most plausible structure.

EXAMPLE

SOLUTION

The most plausible structure is (2)

Formal charge is determined before

completing a Lewis structure to predict

the most stable structure because

formal charge closest to zero.

1) 2)

EXCEPTION TO OCTET RULE

1. Molecular species that do not follow the octet

rule fall under two categories ;

Molecules in which atom has less than an octet

(Incomplete octet)

Molecules with an odd number of electrons

Molecules in which an atom has more than an octet

(Expanded octet)

EXCEPTION TO OCTET RULE 1. Incomplete octet:

Electron - deficient molecules

The central atoms have fewer than eight electrons around them

E.g Be𝐻2

Other examples : Be𝐶𝑙2 , B𝐶𝑙3 , B𝐹3

𝐻 − 𝐵𝑒 − 𝐻

Be 1 X 𝟐𝒆− = 𝟐𝒆−

2H 2 x 𝟏𝒆− = 𝟐𝒆−

Total 4𝒆−

*Be, B, Al – incomplete

EXCEPTION TO OCTET RULE 2. Odd electron molecules : free radicals

Contain an unpaired electron

The central atoms have fewer than eight electrons around them

E.g NO

Other example : N𝑂2

N 1 X 𝟓𝒆− = 𝟓𝒆−

O 1 x 𝟔𝒆− = 𝟔𝒆−

Total 11𝒆−

Most odd electrons molecules have a central atom from an-odd

numbered group, such as N (Group 15) and Cl (Group 17)

EXCEPTION TO OCTET RULE

3. Expanded octet

Central atoms have more than eight electrons around them

Central atoms are normally elements of Period 3 or higher : d orbital available

E.g S𝐹6

Other examples : PC𝑙5 , 𝑆𝑂42−

S 1 X 𝟔𝒆− = 𝟔𝒆−

F 6 x 𝟕𝒆− = 𝟒𝟐𝒆−

Total 48𝒆−

EXERCISE

66

Draw the Lewis structure of the following molecules

and state the special features at the central atoms.

(a) NO

(b) TeCl4

(c) AlBr3

(d) XeF2

Coordinate Covalent Bond (Dative Bond)

1. Coordinate bond is formed when one of the atom donates

both electron (lone pair electron)

2. Known as dative covalent bond or dative bond

3. The atom contributes two electrons to form the coordinate

bond is called the donor atom.

4. The atom which accepts the electron pair from the donor

atom is called the acceptor atom

Coordinate Covalent Bond (Dative Bond) Example 1 :

1. Ammonium ion, 𝑁𝐻4+

2. The nitrogen atom in an ammonia molecules has a lone pair

of electrons

3. Thus, the nitrogen atom can act as a donor atom. In contrast,

the hydrogen ion has an empty 1s orbital. Thus, hydrogen ion

(𝐻+) can act as an acceptor atom.

4. The formation of a coordinate bond in the ammonium ion can

be represented by Lewis structure as follows:

N ∶ H H

H

+ 𝐻+ N

H

H

H

H

+

Coordinate Covalent Bond (Dative Bond) Example 2:

1. Oxonium ion, 𝐻3𝑂+

2. The oxygen atom in the water molecule has two lone pairs of

electrons.

3. An oxonium ion is formed when one of the lone pairs of

electrons is used to form a coordinate bond with a hydrogen

ion.

4. The oxonium ion is called the hydroxonium ion or the

hydronium ion

+

+ 𝐻+ H : O

H

H

O H

H

: