SCHEME TUTORIAL CHAPTER 4

SESSION 2022/2023 TOPIC 4: CHEMICAL BONDING
CHEMISTRY SK015

TUTORIAL 4

4.1: Lewis Structure

1. Complete the table below:

Element Li Be B C N O Cl Ne

No. of Valence 1 2 3 4 5 6 7 8
Electron

Lewis Li Be B C N O F Ne
dot symbol

2. a) Define
Octet rule states that atoms tend to form bonds to obtain 8 electrons in the valence
shell.

b) electronic configuration and type of stability
i. O : 1s2 2s2 2p4
O2- : 1s2 2s2 2p6 (noble gas configuration)
ii. Mg : 1s2 2s2 2p6 3s2
Mg2+ : 1s2 2s2 2p6 (noble gas configuration)
iii. Mn : 1s2 2s2 2p6 3s2 3p6 4s2 3d5
Mn2+ : 1s2 2s2 2p6 3s2 3p6 3d5 (half-filled orbital configuration)
iv. Zn : 1s2 2s2 2p6 3s2 3p6 4s2 3d10
Zn2+ : 1s2 2s2 2p6 3s2 3p6 3d10 (pseudo-noble gas configuration)

3. a) i. define
Ionic bond is the strong electrostatic force between positive and negative ions
that hold them together in solid crystals.

ii. formation of ionic bond in Na2O.

Na2O : Na has 1 valence electron, while O has 4 valence electrons
Na loses 1 electron and forms Na+ ion, while O accepts the electrons to form
O2- ion.
Electrostatic force between Na+ ion and O2- ion form ionic bond.

Na + 2-

+O 2 Na O

b) i. Na

covalent bond
Covalent bond is formed by a pair of electrons shared between two atoms.
Atoms in covalent bond are held by electrostatic forces between the shared
electron and the nuclei of the atom involved.

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ii. formation of covalent bond in CF4

CF4 : C has 4 Valence electrons, while F has 7 Valence electrons. Thus C
shared its valence electrons with F in order to form covalent bond.

F F F F
FC
+
F +C+

+
F
F

c) i. Differences

Covalent bond Dative bond
Covalent bond is formed when two a bond in which the pair of shared
atoms share one or more electron electrons is supplied by one of the
pairs. two bonded atoms.

ii. formation of dative bond in NH4+

+H N H H+ H+
HN H
H
H

4. Lewis structure and type of octet

a) BeCl2 d) SO2

Cl Be Cl OSO

Incomplete octet Octet
b) CO2
e) NO2
OCO
-
Octet
c) BF3 ONO
Odd number electron
F
FB f) CH3Cl
Cl
F
H CH
Incomplete octe H

Octet

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g) PCl5 j) XeF4

Cl F
Cl F Xe F

Cl P F
Cl Cl

Expanded octet Expanded octet
h) SF4 k) CO32-

F 2-
FS
F OCO
F O

Expanded octet Octet
i) SF6
l) NO2+
FF
FS F +
FF
ONO

Expanded octet Octet

5. Lewis structure C2H4

C2H6 HH
HH H CC H

H CC H
HH

Bond length
C2H6 has longer C-C bond than C2H4.

6. a) possible structures and formal charge

00 0 -1 0 +1
S CS S CS

AB

most plausible structure.
Structure A is the most plausible structure because the formal charge for each atom is
zero.

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b) possible structures and formal charge

-1 0 0 -2 0 +1 0 0 -1
N CO N CO N CO

AB C

most plausible structure.
Structure C is the most plausible structure because the negative formal charge is on
more electronegative atom, O.

7. a) definition resonance structures
Lewis structure having same arrangement of atoms but differ in the
position of their electrons

b) resonance structures of the O3 and NO3−.
O3

O O+ O O O+ O

NO3− - -
- ONO ONO

ONO O O

O

4.2: Molecular Shape and Polarity

1. a) definition
Valence Shell Electron-pair Repulsion (VSEPR) theory state that electron pairs around
the central atom is located as far as possible from the others in order to minimize the
repulsions.

b) i) All molecules has 4 electron pairs around their central atom. Therefore, Electron pair
arrangement for all molecule: tetrahedral

CH4
• 4 Bonding pair and no lone pair
• Based on VSEPR theory, the valence electron pairs around central atom are

oriented as far as possible to minimize the repulsion between them.
• The strength of electron repulsion of Bonding Pair-Bonding Pair are equals.
• Molecular geometry: tetrahedral
• The H-N-H angle: 109.5o.
NH3
• 3 Bonding pair and 1 lone pair
• Based on VSEPR theory, the valence electron pairs around central atom are

oriented as far as possible to minimize the repulsion between them.
• The strength of electron repulsion are Lone Pair-Bonding Pair > Bonding Pair-

Bonding Pair.
• Molecular geometry: trigonal pyramidal
• The H-N-H angle: 107.3o.
•
H2O

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• 2 Bonding pair and 2 lone pair
• Based on VSEPR theory, the valence electron pairs around central atom are

oriented as far as possible to minimize the repulsion between them.
• The strength of electron repulsion are Lone Pair-Lone Pair > Lone Pair-Bonding

Pair > Bonding Pair-Bonding Pair.
• Molecular geometry: bent
• The H-N-H angle: 104.5o.

ii) SiF4
• 4 electron pairs around central atom.
• Electron pair arrangement: tetrahedral
• 4 Bonding pair and no lone pair
• Based on VSEPR theory, the valence electron pairs around central atom are
oriented as far as possible to minimize the repulsion between them.
• The strength of electron repulsion of Bonding Pair-Bonding Pair are equals.
• Molecular geometry: tetrahedral

SF4
• 5 electron pairs around central atom.
• Electron pair arrangement: trigonal bipyramidal
• 4 Bonding pair and 1 lone pair
• Based on VSEPR theory, the valence electron pairs around central atom are

oriented as far as possible to minimize the repulsion between them.
• The strength of electron repulsion are Lone Pair-Bonding Pair > Bonding Pair-

Bonding Pair.
• Molecular geometry: see-saw

2. a) Definition dipole moment.
Product of the positive charge and distance between the charges.
(Dictionary of chemistry)

b) factors that affect the polarity of a molecule.
• Molecular geometry
• Electronegativity of bonded atom

c) Explanation
• Bond between C-Cl is polar.
• The atoms in CCl4 are arrange symmetrically in tetrahedral shape.
• The bond dipoles of CCl4 may cancel each other.
• Net charge = 0.

3. state molecular geometries, and determine polarity

a) BeCl2

Cl Be Cl

• Molecular geometry: linear

Cl Be Cl

• Cl is more electronegative than Be. Bond between Be-Cl is polar.
• The dipole moment can cancel each other. equal to 0.
• Non-Polar

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b) CO2
OCO

• Molecular geometry: linear
• O is more electronegative than C. Bond between C-O is polar.
• The dipole moment can cancel each other. equal to 0.
• Non-Polar

c) BF3

F
FB

F

• Molecular geometry: trigonal planar

F
FB

F

• F is more electronegative than B. Bond between B-F is polar.
• The dipole moment can cancel each other. equal to 0.
• Non-Polar

d) SO2

OSO

• Molecular geometry: V-shaped

S
OO
• O is more electronegative than S. Bond between S-O is polar.
• The dipole moment cannot cancel each other. not equal to 0.
• Polar

e) NO2

-

ONO

• Molecular geometry: V- shaped

N
O

O

• O is more electronegative than N. Bond between N-O is polar.
• The dipole moment cannot cancel each other. not equal to 0.
• Polar

f) CH3Cl
Cl

H CH

H
• Molecular geometry: tetrahedral

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Cl

HC H
H

• Electronegative of Cl > C > H. Bond between C-H and C-Cl is polar.
• The dipole moment cannot cancel each other. does not equal to 0.
• Polar
g) PCl5

Cl
Cl

Cl P
Cl

Cl

• Molecular geometry: trigonal bipyramidal

Cl
Cl

Cl P
Cl Cl

• Cl is more electronegative than P. Bond between P-Cl is polar.
• The dipole moment can cancel each other. equal to 0.
• Non-Polar

h) SF4

F

FS F

F

• Molecular geometry: see-saw

F
F

S
F

F

• F is more electronegative than S. Bond between S-F is polar.
• The dipole moment can not cancel each other. not equal to 0.
• Polar

i) SF6

FF

FS F

FF

• Molecular geometry: octahedral

F F
F

S

FF

F

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• F is more electronegative than S. Bond between S-F is polar.
• The dipole moment can cancel each other. equal to 0.
• Non-Polar

j) XeF4

F

F Xe F

F

• Molecular geometry: square planar

FF
Xe

FF

• F is more electronegative than Xe. Bond between Xe-F is polar.
• The dipole moment can cancel each other. equal to 0.
• Non-Polar

4.3: Orbital Overlap and Hybridisation

1. a) Define
Valence bond theory state that a covalent bond is formed when neighbouring atomic
orbitals overlaps.

b) formation of sigma and pi bonds.

Sigma bond:
Overlapping of orbitals end to end.




Pi bond:
Overlapping of orbitals side to side.

c) formation of bond in O2.
OO

1 σ bond, 1 π bonds

Valence orbital diagram;
O:

2s 2p

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• 2p orbitals consist of 2 unpaired electrons.

• The 2p orbitals of an oxygen atom will overlap with 2p orbitals of another O atom
to form 1 σ bond and 1 π bonds.

Orbital overlapping diagram;


O O




2. a) definition hybridisation of atomic orbitals
Mixing of atomic orbitals in an atom to generate new set of orbitals.

b) type of hybridisation

HH
2
sp
C C sp2

H sp3 C sp
CN

HH

c) hybridisation of the underlined atoms
i. BF3 sp2
ii. CCl4 sp3

3. Type of hybridisation and orbital overlapping diagram
a) BeH2
H Be H
• 2 Bonding pair and no lone pair
• Type of hybrid: sp

Valence orbital diagram, H (ground state):
Be (ground state):

2s 2p 1s
Be (excited state):

2s 2p
Be (hybrid):

sp 2p

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Orbital overlapping diagram;



sp sp
Be

HH

b) BH3

H BH

H

• 3 Bonding pair and no lone pair
• Type of hybrid: sp2

Valence orbital diagram, H (ground state):
B (ground state):

2s 2p 1s
B (excited state):

2s 2p
B (hybrid):

2 2p

sp

Orbital overlapping diagram;

H


sp2

B sp2 
H  sp2 H

c) PCl5

Cl Cl
Cl P Cl

Cl

• 5 Bonding pair and no lone pair
• Type of hybrid: sp3d

Valence orbital diagram,
P (ground state):

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3s 3p 3d
P (excited state): 3d

3s 3p
P (hybrid):

sp3d 3d
Cl (ground state):

3s 3p

Orbital overlapping diagram;

3p  3p
Cl
Cl


 3 sp3d 

sp d

3p sp3d sp3d 3p
Cl Cl
sp3d P



Cl

3p

d) SeF4 F (ground state):
F Se F

FF
• 4 Bonding pair and 1 lone pair
• Type of hybrid: sp3d

Valence orbital diagram,
Se (ground state):

4s 4p 4d 2s 2p
Se (excited state):

4s 4p 4d

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Se (hybrid): CHEMISTRY SK015

sp3d 4d

Orbital overlapping diagram;

sp3d

F sp3d Se F
2p sp3d 2p

 sp3d sp3d 

F

F 2p

2p

e) SCl6 Cl (ground state):

Cl
Cl Cl

S
Cl Cl

Cl
• 6 Bonding pair and no lone pair
• Type of hybrid: sp3d2

Valence orbital diagram,
S (ground state):

3s 3p 3d 3s 3p
S (excited state):

3s 3p 3d
S (hybrid):

sp3d2 3d

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Orbital overlapping diagram;
Cl (ground state):
3p 3s 3p

Cl

3p  3p
Cl


32 Cl

sp d

sp3d2 sp3d2

 sp3ds2pS3dP2sp3d2 

3p Cl Cl 3p



Cl

f) ICl2− 3p

-
Cl I Cl

• 2 Bonding pair and 3 lone pair
• Type of hybrid: sp3d

Valence orbital diagram,
I− (ground state):

5s 5p 5d
I− (excited state):

5s 5p 5d
I− (hybrid):

sp3d 5d

Orbital overlapping diagram;

 3 sp3d 

sp d

3p sp3d 3 3p
Cl Cl
sp d
sp3d I-

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g) ICl4+
Cl (ground state):
+ 3s 3p

Cl I Cl
Cl Cl

• 4 Bonding pair and 1 lone pair
• Type of hybrid: sp3d

Valence orbital diagram,
I+ (ground state):

5s 5p 5d
I+ (excited state):

5s 5p 5d
I+ (hybrid):

sp3d 5d

Orbital overlapping diagram;

Cl sp3d I+ Cl
2p sp3d sp3d 2p

 3 sp3d 

sp d

Cl

Cl 2p

2p

4.4: Intermolecular Forces

1. a) Attractive force between polar molecule
• Polar molecule attract each other via dipole-dipole forces.
• When partial +ve charge of one molecule nears the partial –ve charge on another
molecule, they attract each other and form dipole-dipole forces.

b) London dispersion forces
• When molecules are very close to each other, the London Dispersion Forces is
significant because electrons repel one another.
• The motion of electrons in one atom influence the motions of electrons in
neighboring atom.
• Thus, the temporary dipole of 1 atom can induce a similar dipole of an adjacent
atom causing the atoms to attract to each other forming London dispersion forces.

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c) type of van der Waals forces
i. Cl2 London Dispersion Forces
i. PCl3 dipole-dipole forces and London Dispersion Forces
iii. H2S dipole-dipole forces and London Dispersion Forces
iv. CCl4 London Dispersion Forces

2. a) factors that influence the strength of van der Waals forces.
Molecular size, and polarity of molecules.

b) i. Boiling point of ICl > Br2
• ICl is polar, Br2 is non-polar.
• Intermolecular forces between Br2 molecules is London Dispersion Forces.
• Intermolecular forces between ICl molecules is dipole-dipole forces.
• dipole-dipole forces stronger than London Dispersion Forces.

ii. Boiling point of I2 > Br2
• Molecular weight of I2 > Br2.
• Strength of van der Waals forces between I2 molecules > Br2.

iii. Boiling point of HBr > CF4
• HBr is polar, CF4 is non-polar.
• Intermolecular forces between CF4 molecules is London Dispersion forces.
• Intermolecular Forces between HBr molecules is dipole-dipole forces.
• dipole-dipole forces stronger than London Dispersion Forces.

3. a) describe HBond and explain factors that influence the hydrogen bond.

• Attractive forces between H which is covalently bonded to highly electronegative

atom (F, O, N) in one molecule and highly electronegative atom (F,O,N) in another

atom.

• Factors that influences the strength of hydrogen bond is electronegativity of F, O,

and N. Higher the electronegativity, the stronger the hydrogen bond.

b) i. Boiling point of H2O > CH4

• H2O can form Hydrogen Bond between molecules.

• CH4 can only form van der Waals forces.

• Hydrogen Bond stronger than van der Waals forces.

ii. Boiling point of H2O > NH3
• O is more electronegative than N.
• Strength of Hydrogen Bond between H2O molecules > between NH3
molecule.

4. Boiling point of ethanol > dimethyl ether.
• Ethanol can form Hydrogen Bond.
• Dimethyl ether can only form van der Waals forces.
• Hydrogen bond stronger than van der Waals forces.

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5. a) The band energy gap is the energy gap between a fully occupied valence band and an empty
conduction band.

b) In conductors, the valence and conduction bands overlap each other. In insulators, there is
large energy gap between valence and conduction bands, while in semiconductors this energy
gap is small, so that electrons can make the jump up to the conductions band.

6. a) Electron sea model

e ee e ee

Al3+ Al3+ Al3+ Al3+
e eee eee e
e e Al3+ Al3+ Al3+
Al3+ e
ee
e eee ee ee e

Al3+ Al3+ Al3+ Al3+
e ee ee e e e

• When Al atoms are arrange closely packed to each other, each of Al atom will
released its valence electrons and forming sea of delocalized electrons and Al3+

ions.
• The attractive forces between Al3+ ions and sea of electrons forming metallic bond.

b) i. Boiling point Be > Mg.
• Size of Be2+ < Mg2+.
• Strength of metallic bond Be > Mg.
• Energy needed to break the metallic bond in Be > Mg.

ii. Boiling point Al > Mg.
• Size of Al3+ < Mg2+.
• Al has greater number of valence electrons than Mg.
• Strength of metallic bond Al > Mg.
• Energy needed to break the metallic bond in Al > Mg.

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MEKA 4 Lewis dot symbol
1. a)
i. H

O

ii.

iii. Al

b) formation bond
i. Al2O3: Al has 3 Valence electrons, while O has 6 Valence electrons.
Al loses 3 electrons and form Al3+ ion, while O accept the electrons to form O2-
ion.
Electrostatic force between Al3+ ion and O2- ion form ionic bond.

O 2-
Al 2 Al 3+3 O

+O

Al

O

ii. H2O: O has 6 Ve, while H has 1 Ve
Oxygen will share its 2 valence electrons with two H. Electrostatic force
between nucleus of atom and electrons shared form covalent bond.

H + O +H HO H

iii.

+H O H+ HO H+
H
H

2. Lewis structure b) AlBr3
a) SeF4
Br
F Se F Br Al

FF Br
Expanded octet
Incomplete octet
c) NO3-
d) CCl2F2
-
O Cl
ON FC F
O
Cl
Octet
Octet
e) NO2+
f) NO
+
NO
ONO Odd number electron
Octet

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3. CO2
a) Possible Lewis structure

OCO OCO

AB

b) Formal Charge 0 00
+1 0 -1
OCO
OCO
B
A

c) Most plausible structure
Structure B is most stable because formal charge of each atom equals to zero.

OCS
a) Possible Lewis structure

OCS OCS OCS

AB C

b) Formal Charge

+1 0 -1 000 -1 0 +1

OCS OCS OCS

A BC
c) Most plausible structure

Structure B is the most stable because formal charge of each atom equals to zero.

N3-

a) Possible Lewis structure

--

NNN NNN

AB

b) Formal Charge

0 +1 -2 - +1 -

NNN NNN

A -1 B -1

c) Most plausible structure
Structure B is most stable because formal charge of each atoms nearer to zero.

4. resonance O OO
a) O3

OO O

b) NO2+ + +
+
ONO ONO
ONO

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5. a) i. Valence Shell Electron-pair Repulsion (VSEPR) theory state that electron
pairs around the central atom is located as far as possible from the others in
order to minimize the repulsions.

ii. The strength of repulsion between lone pair-lone pair > lone pair – bonding
pair > bonding pair – bonding pair.

b) CO2
i. Lewis structure

OCO

iii. electron pair arrangement : linear

iii. number of bonding pair: 2
number of lone pair : 0

iv. Bonding pair-bonding pair repulsion are equal.

v. Molecular geometry: Linear
OCO

vi. Bond angle: 180o

vii. Electronegativity
• O is more electronegative than C
• O-C bond is polar

viii. Polarity
• Bond dipole can cancel each other, μ=0.
• Non-Polar.

HCl
i. Lewis structure

H Cl

ii. Molecular geometry: Linear
H Cl

iii. Electronegativity
• Cl is more electronegative than H
• H-Cl bond is polar

BF3
i. Lewis structure

FBF
F

ii. electron pair arrangement : trigonal planar

iii. number of bonding pair: 3
number of lone pair : 0

iv. Bonding pair-bonding pair repulsion are equal.

v. Molecular geometry: trigonal planar

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FF
B

F

vi. Bond angle: 120o

vii. Electronegativity
• F is more electronegative than B
• B-F bond is polar

viii. Polarity
• Bond dipole can cancel each other, μ=0.
• Non-Polar.

SnCl2 *correction on question 5(b)
i. Lewis structure

Cl Sn Cl

ii. electron pair arrangement : trigonal planar

iii. number of bonding pair: 2
number of lone pair : 1

iv. lone pair-bonding pair repulsion > bonding pair-bonding pair repulsion.

v. Molecular geometry: bent
Cl

Sn

Cl

vi. Bond angle: <120o

vii. Electronegativity
• Cl is more electronegative than Sn
• Sn-Cl bond is polar

viii. Polarity
• Bond dipole cannot cancel each other, μ≠0.
• Polar.

SO2
i. Lewis structure

OS O

ii. electron pair arrangement : trigonal planar

iii. number of bonding pair: 2
number of lone pair : 1

iv. lone pair-bonding pair repulsion > bonding pair-bonding pair repulsion.

v. Molecular geometry: bent
O

S
O

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vi. Bond angle: <120o

vii. Electronegativity
• O is more electronegative than S
• S-O bond is polar

viii. Polarity
• Bond dipole cannot cancel each other, μ≠0.

• Polar.

CCl4 Lewis structure
i. Cl

Cl C Cl
Cl

ii. electron pair arrangement : tetrahedral

iii. number of bonding pair: 4
number of lone pair : 0

iv. Bonding pair-bonding pair repulsion are equal.

v. Molecular geometry: tetrahedral
Cl

Cl C
Cl Cl

vi. Bond angle: 109.5o

vii. Electronegativity
• Cl is more electronegative than C
• C-Cl bond is polar

viii. Polarity
• Bond dipole can cancel each other, μ=0.
• Non Polar.

CH2Cl2
i. Lewis structure

H
Cl C Cl

H

ii. electron pair arrangement : tetrahedral

iii. number of bonding pair: 4
number of lone pair : 0

iv. Bonding pair-bonding pair repulsion are equal.

v. Molecular geometry: tetrahedral

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H

Cl C
Cl H

vi. Bond angle: 109.5o

vii. Electronegativity
• Cl is more electronegative than C
• C-Cl bond is polar

viii. Polarity
• Bond dipole cannot cancel each other, μ≠0.
• Polar.

NF3
i. Lewis structure

FN F
F

ii. electron pair arrangement : tetrahedral

iii. number of bonding pair: 3
number of lone pair : 1

iv. Lone pair-bonding pair repulsion > Bonding pair-bonding pair repulsion.

v. Molecular geometry: trigonal pyramidal
FN
FF

vi. Bond angle: <109.5o

vii. Electronegativity
• F is more electronegative than N
• N-F bond is polar

viii. Polarity
• Bond dipole cannot cancel each other, μ≠0.
• Polar.

H2O
i. Lewis structure

HO H

ii. electron pair arrangement : tetrahedral

iii. number of bonding pair: 2
number of lone pair : 2

iv. lone pair-lone pair repulsion > Lone pair-bonding pair repulsion > Bonding pair-
bonding pair repulsion.

v. Molecular geometry: bent

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O
HH

vi. Bond angle: 104.5o

vii. Electronegativity
• O is more electronegative than H
• O-H bond is polar

viii. Polarity
• Bond dipole cannot cancel each other, μ≠0.

• Polar.

PCl5 Lewis structure
i. Cl Cl

Cl P Cl

Cl

ii. electron pair arrangement : trigonal bipyramidal

iii. number of bonding pair: 5
number of lone pair : 0

iv. Bonding pair-bonding pair repulsion are equals.

v. Molecular geometry: trigonal bipyramidal
Cl

Cl Cl
P

Cl

Cl

vi. Bond angle: 120o, 90o

vii. Electronegativity
• Cl is more electronegative than P
• P-Cl bond is polar

viii. Polarity
• Bond dipole can cancel each other, μ=0.

• Non Polar.

SF4
i. Lewis structure

FF

FSF

ii. electron pair arrangement : trigonal bipyramidal

iii. number of bonding pair: 4
number of lone pair : 1

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iv. lone pair-bonding pair repulsion > Bonding pair-bonding pair repulsion.

v. Molecular geometry: see-saw
F

F
S

F
F

vi. Bond angle: <120o, <90o

vii. Electronegativity
• F is more electronegative than S
• S-F bond is polar

viii. Polarity
• Bond dipole cannot cancel each other, μ≠0.
• Polar.

ICl3
i. Lewis structure

Cl I Cl

Cl

ii. electron pair arrangement : trigonal bipyramidal

iii. number of bonding pair: 3
number of lone pair : 2

iv. lone pair-lone pair repulsion > lone pair-bonding pair repulsion > Bonding pair-
bonding pair repulsion.

v. Molecular geometry: T-shape
Cl

I Cl

Cl

vi. Bond angle: 90o

vii. Electronegativity
• Cl is more electronegative than I
• I-Cl bond is polar

viii. Polarity
• Bond dipole cannot cancel each other, μ≠0.

• Polar.

XeF2 Lewis structure
i. F Xe F

ii. electron pair arrangement : trigonal bipyramidal

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iii. number of bonding pair: 2
number of lone pair : 3

iv. lone pair-lone pair repulsion > lone pair-bonding pair repulsion > Bonding pair-
bonding pair repulsion.

v. Molecular geometry: linear
F

Xe

F

vi. Bond angle: 180o

vii. Electronegativity
• F is more electronegative than Xe
• Xe-F bond is polar

viii. Polarity
• Bond dipole cannot cancel each other, μ≠0.
• Polar.

SF6
i. Lewis structure

FF

FSF

FF

ii. electron pair arrangement : Octahedral

iii. number of bonding pair: 6
number of lone pair : 0

iv. Bonding pair-bonding pair repulsion are equals.

v. Molecular geometry: octahedral
F

FF
S

FF
F

vi. Bond angle: 90o

vii. Electronegativity
• F is more electronegative than S
• S-F bond is polar

viii. Polarity
• Bond dipole can cancel each other, μ=0.
• Non Polar.

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BrF5 Lewis structure
i. FF

Br F
FF

ii. electron pair arrangement : Octahedral

iii. number of bonding pair: 5
number of lone pair : 1

iv. Lone pair-bonding pair repulsion > Bonding pair-bonding pair repulsion.

v. Molecular geometry: square pyramidal
FF
Br
FF

F

vii. Bond angle: <90o

vii. Electronegativity
• F is more electronegative than Br
• Br-F bond is polar

viii. Polarity
• Bond dipole cannot cancel each other, μ≠0.

• Polar.

XeF4 Lewis structure
i. FF

Xe
FF

ii. electron pair arrangement : Octahedral

iii. number of bonding pair: 4
number of lone pair : 2

iv. Lone pair-lone pair repulsion > Lone pair-bonding pair repulsion > Bonding
pair-bonding pair repulsion.

v. Molecular geometry: square planar
FF
Xe
FF

viii. Bond angle: 90o

vii. Electronegativity
• F is more electronegative than Xe
• Xe-F bond is polar

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viii. Polarity
• Bond dipole can cancel each other, μ=0.

• Non Polar.

6. a) i. Valence bond theory state that a covalent bond is formed when neighbouring
atomic orbitals overlapped .

ii. sigma bond is formed when atomic orbitals is overlap end-to-end while pi bond
is formed when atomic orbitals is overlap sideways.

b) Shape of orbital hybrid
i. sp

sp sp

ii. sp2

sp2

sp2 sp2

iii. sp3

sp3

3 sp3 3

sp sp

iv. sp3d

sp3d sp3d
sp3d sp3d

3

sp d

v. sp3d2

sp3d2

sp3d2 sp3d2

sp3ds2p3dP2sp3d2

c) BeCl2

i. & ii. Lewis structure & formal charge

0 0
0
Cl Be Cl

iii. number of bonding pair : 2
number of lone pair:0

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iv. molecular geometry: linear

Cl Be Cl

v. valence orbital diagram
Be (ground state):

2s 2p

Be (excited state):

2s 2p
Be (hybrid):

sp 2p
Cl (ground state):

3s 3p

vi. orbital overlapping diagram


3p sp sp 3p
Cl Be Cl

C2H2 Lewis structure & formal charge
i. & ii.

0 0
HC C H0

0

iii. number of bonding pair : 2
number of lone pair:0

iv. molecular geometry: linear

HC CH

v. valence orbital diagram
C (ground state):

2s 2p
C (excited state):

2s 2p
C (hybrid):

sp 2p

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H (ground state): CHEMISTRY SK015
1s

vi. orbital overlapping diagram


 2p 2p  2p 2p 

sp sp sp C sp
C

HH

AlCl3 Lewis structure & formal charge
i. & ii. 0

Cl 0
0 Cl
Cl Al

0

iii. number of bonding pair : 3
number of lone pair:0

iv. molecular geometry: trigonal planar
Cl

Al Cl
Cl

v. valence orbital diagram
Al (ground state):

3s 3p
Al (excited state):

3s 3p
Al (hybrid):

2 3p

sp

Cl (ground state):

3s 3p

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vi. orbital overlapping diagram

3p Cl  3p
 Cl

2

sp2 sp
Al

sp2



Cl

3p

CH2O Lewis structure & formal charge
i. & ii. 0
O

0 0
HC H

0

iii. number of bonding pair : 3
number of lone pair:0

iv. molecular geometry: trigonal planar

O

C

HH

v. valence orbital diagram
C (ground state):

2s 2p
C (excited state):

2s 2p
C (hybrid):

sp2 2p
O (ground state):

2s 2p
O (excited state):

2s 2p
O (hybrid state):

2 2p

sp

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H (ground state):

1s

vi. orbital overlapping diagram

1s   1s
H sp2 sp2 H

2p
C
sp2

 

sp2
O
2p

sp2 sp2

CCl4 Lewis structure & formal charge
i. & ii.

Cl 0

0

Cl C0 Cl
0

Cl
0

iii. number of bonding pair : 4
number of lone pair:0

iv. molecular geometry: tetrahedral
Cl

C
Cl

Cl Cl

v. valence orbital diagram
C (ground state):

2s 2p
C (excited state):

2s 2p
C (hybrid):

sp3

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Cl (ground state):

3s 3p

vi. orbital overlapping diagram

3p

Cl


sp3

sp3 sp 3 
Psp3 Cl
C 3p

3p Cl  Cl

3p

NF3 Lewis structure & formal charge
i. & ii. 0

F NF
00

F
0

iii. number of bonding pair : 3
number of lone pair:1

iv. molecular geometry: trigonal pyramidal
N
F

FF

v. valence orbital diagram
N (ground state):

2s 2p
N (excited state):

2s 2p
N (hybrid):

3

sp
F (ground state):

2s 2p

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vi. orbital overlapping diagram

sp3

 sp3 
Psp3
sp3 N F
2p

2p F F

2p

H2O Lewis structure & formal charge
i. & ii.

O H0
0

0H

iii. number of bonding pair : 2
number of lone pair:2

iv. molecular geometry: bent
O

HH

v. valence orbital diagram
O (ground state):

2s 2p
O (excited state):

2s 2p
O (hybrid state):

3

sp
H (ground state):

1s
vi. orbital overlapping diagram

3

sp

 sp3 O sp3
Psp3
1s
H  1s H

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PCl5 Lewis structure & formal charge
i. & ii. 00
Cl Cl

0
Cl P0 Cl 0

Cl 0

iii. number of bonding pair : 5
number of lone pair:0

iv. molecular geometry: trigonal bipyramidal
Cl Cl

Cl P Cl

Cl

v. valence orbital diagram
P (ground state):

3s 3p 3d
P (excited state): 3d

3s 3p
P (hybrid):

sp3d 3d
Cl (ground state):

3s 3p

vi. orbital overlapping diagram

3p
Cl 3p
  Cl

 sp3d sp3d 

3p sp3d 3 3p
Cl Cl
sp d

sp3d P



Cl

3p

SF4

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CHEMISTRY SK015

i. & ii. Lewis structure & formal charge
0 00
F SF

F F0
0

iii. number of bonding pair : 4
number of lone pair:1

iv. molecular geometry: see-saw
F SF

FF
v. valence orbital diagram

S (ground state):

3s 3p 3d
S (excited state):

3s 3p 3d
S (hybrid):

sp3d 3d
F (ground state):

2s 2p
vi. orbital overlapping diagram

sp3d

F sp3d S F
2p 3 2p

sp d

 sp3d sp3d 

ICl3 F
i. & ii. 
F 2p
2p

Lewis structure & formal charge
00

0 Cl I Cl

Cl 0

iii. number of bonding pair : 3
number of lone pair:2

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iv. molecular geometry: T-shape
Cl I Cl

Cl

v. valence orbital diagram
I (ground state):

5s 5p 5d
I (excited state):

5s 5p 5d
I (hybrid):

sp3d 5d
Cl (ground state):

3s 3p
vi. orbital overlapping diagram

 3 sp3d 

sp d

3p sp3d sp3d 3p
Cl Cl
sp3d I



Cl

3p

XeF2 Lewis structure & formal charge
i. & ii.
0 00
F Xe F

iii. number of bonding pair : 2
number of lone pair:3

iv. molecular geometry: linear
F Xe F

v. valence orbital diagram
Xe (ground state):

5s 5p 5d
Xe (excited state):

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5s 5p 5d
Xe (hybrid):

sp3d 5d
F (ground state):
2s 2p

vi. orbital overlapping diagram

F Xesp3d F
2p 2p
sp3d 3

sp d

 sp3d sp3d 

SF6 Lewis structure & formal charge
i. & ii. 0
F

0F F 0

S
0F F 0

F
0

iii. number of bonding pair : 6
number of lone pair: 0

iv. molecular geometry: octahedral
F

FF

S

FF
F

v. valence orbital diagram
S (ground state):

3s 3p 3d
S (excited state):

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3s 3p 3d
S (hybrid):

sp3d2 3d
F (ground state):

2s 2p

vi. orbital overlapping diagram

2p

F

2p   2p
F 2p
sp3d2 F
 F
2p F sp3d2 sp3d2

sp3ds2pS3dP2sp3d2 



F

2p

BrF5
i. & ii. Lewis structure & formal charge

0F 0 F0
0F Br

F0

F
0

iii. number of bonding pair : 5
number of lone pair: 1

iv. molecular geometry: square pyramidal
FF

Br
FF

F

v. valence orbital diagram
Br (ground state):

5s 5p 5d

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Br (excited state):

5s 5p 5d
Br (hybrid):

sp3d2 5d
F (ground state):

2s 2p

vi. orbital overlapping diagram

2p  2p
F 2p
sp3d2 F
 F
2p F sp3d2 sp3d2

sp3ds2pB3rdP2sp3d2 



F

2p

XeF4 Lewis structure & formal charge
i. & ii.

0F 0 F0
0F Xe

F0

iii. number of bonding pair : 4
number of lone pair: 2

iv. molecular geometry: square planar
FF

Xe
FF

v. valence orbital diagram
Xe (ground state):

5s 5p 5d
Xe (excited state):

5s 5p 5d

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Xe (hybrid): CHEMISTRY SK015

sp3d2 5d
F (ground state):

2s 2p

vi. orbital overlapping diagram

2p  2p
F 2p
sp3d2 F
 F
2p F 32 32

sp d sp d

sp3ds2pX3edP2sp3d2 

7. a) -hydrogen bond
-dipole-dipole forces
-metallic bond
- London dispersion forces

b) i. CH3Br – van der Waals forces
CH3F – van der Waals forces

Boiling point of CH3Br > CH3F.
• CH3Br bigger than CH3F.
• Strength of van der Waals forces between molecules CH3Br > CH3F.
• More energy needed to overcome forces between CH3Br molecules than

CH3F molecules.

ii. NH3 – hydrogen bond
CH4 – van der Waals forces

Boiling point of NH3 > CH4.
• Hydrogen bond stronger than van der Waals forces.
• More energy needed to break hydrogen bonds between NH3 molecules.

iii. CCl4 – non- polar molecules - London dispersion forces
CH3Cl – polar molecules - dipole-dipole forces

Boiling point of CH3Cl > CCl4.
• Dipole-dipole forces stronger than London dispersion forces.
• More energy needed to overcome dipole-dipole forces between CH3Cl

molecules.

iv. H2O – hydrogen bond
NH3 – hydrogen bond

Boiling point of H2O > NH3.
• O is more electronegative than N.
• Strength of hydrogen bond between molecules H2O > NH3.
• More energy needed to break hydrogen bonds between H2O molecules.

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v. H2O – hydrogen bond
HF – hydrogen bond

Boiling point of H2O > HF.
• H2O can form more hydrogen bond than HF.
• Strength of hydrogen bond between molecules H2O > HF.
• More energy needed to break hydrogen bonds between H2O molecules.

c) i. Be

ee e ee
Be2+ Be2+
e Be2+ e Be2+ e ee

Be2+ e Be2+ Be2+ Be2+
e e
e e e e e
Be2+ e Be2+
Be2+ Be2+

ee ee e

• When Be atoms are arrange closely packed to each other, each of Be atom

will released its valence electrons and forming sea of delocalized electrons

and Be2+ ions.

• The attractive forces between Be2+ ions and sea of electrons forming

metallic bond.

Mg

ee e ee
Mg2+ Mg2+
e Mg2+ e Mg2+ e ee

Mg2+ e Mg2+ Mg2+ Mg2+
e e
e e e e e
Mg2+ eMg2+
Mg2+ Mg2+

ee ee e

• When Mg atoms are arrange closely packed to each other, each of Mg

atom will released its valence electrons and forming sea of delocalized

electrons and Mg2+ ions.

• The attractive forces between Mg2+ ions and sea of electrons forming

metallic bond.

ii. - size of ion Mg2+ > Be2+

iii. - melting point of Be > Mgl.
- size of ion Mg2+ > Be2+.

- Be metal has stronger metallic bond.

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KUMBE 4 possible structures and formal charge
1. a)
0 -1
b) O O

H CO H H C O H0
0 00 0 0 0 +1

A B

most plausible structure.
Structure A is the most plausible structure because the formal charge for each atom is
zero.

i. SF2
0F
S0
0 F stable, formal charge of all atoms are zero.

ii. SF3+ *correction

0F F 0
S+

0 F unstable, formal charge of S is +1.

iii. SF4

0F 0 F 0

S

0F F0
stable, formal charge of all atoms are zero.

iv. SF5+ *correction

0F F0
0F S+

F0

F

0 unstable, formal charge of S is +1.
v. SF6

0
F
0F F 0
0
S
0F F 0

F
0 stable, formal charge of all atoms are zero.

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2. a) SiF4

FF

Si
FF

• 4 electron pairs around central atom.
• Electron pair arrangement: tetrahedral
• 4 Bonding pair and no lone pair
• Based on VSEPR theory, the valence electron pairs around central atom are

oriented as far as possible to minimize the repulsion between them.
• The strength of repulsion of Bonding Pair-Bonding Pair are equals.
• Molecular geometry: tetrahedral

SF4

FF

S
FF

• 5 electron pairs around central atom.
• Electron pair arrangement: trigonal bipyramidal
• 4 Bonding pair and 1 lone pair
• Based on VSEPR theory, the valence electron pairs around central atom are

oriented as far as possible to minimize the repulsion between them.
• The strength of repulsion are Lone Pair-Bonding Pair > Bonding Pair-Bonding Pair.
• Molecular geometry: see-saw

XeF4

FF

Xe
FF

• 6 electron pairs around central atom.
• Electron pair arrangement: octahedral
• 4 Bonding pair and 2 lone pair
• Based on VSEPR theory, the valence electron pairs around central atom are

oriented as far as possible to minimize the repulsion between them.
• The strength of repulsion are Lone pair-bonding > Lone Pair-Bonding Pair >

Bonding Pair-Bonding Pair.
• Molecular geometry: square planar.

b) All species have 4 electron pairs around their central atom. Therefore, Electron pair
arrangement for all species are tetrahedral

NH2-
• 2 Bonding pair and 2 lone pair
• Based on VSEPR theory, the valence electron pairs around central atom are

oriented as far as possible to minimize the repulsion between them.
• The strength of repulsion are Lone pair-bonding > Lone Pair-Bonding Pair >

Bonding Pair-Bonding Pair.
• Molecular geometry: bent
• The H-N-H angle: 105o.

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NH3
• 3 Bonding pair and 1 lone pair
• Based on VSEPR theory, the valence electron pairs around central atom are

oriented as far as possible to minimize the repulsion between them.
• The strength of repulsion are Lone Pair-Bonding Pair > Bonding Pair-Bonding Pair.
• Molecular geometry: trigonal pyramidal
• The H-N-H angle: 107o.

NH4+
• 4 Bonding pair and no lone pair
• Based on VSEPR theory, the valence electron pairs around central atom are

oriented as far as possible to minimize the repulsion between them.
• The strength of repulsion of Bonding Pair-Bonding Pair are equals.
• Molecular geometry: tetrahedral
• The H-N-H angle: 109o.

3. PF3

FPF

F

• Molecular geometry trigonal pyramidal.

PF
FF
• F is more electronegative than P. P-F bond is polar.
• Bond dipole cannot cancel each other. μ ≠ 0.
• Polar.

BF3 F
FB

F

• Molecular geometry trigonal planar.
F

FB

F

• F is more electronegative than B. B-F bond is polar.
• Bond dipole can cancel each other. μ = 0.
• Non Polar.

4. a) type of hybrid

sp2

H aO sp3

H N Ca Cb O b H

sp3 HH

sp3 sp2

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CHEMISTRY SK015

b) orbital overlapping diagram
N (ground state):

2s 2p
N (excited state):

2s 2p
N (hybrid):

3

sp
Ca (ground state):

2s 2p
Ca (excited state):

2s 2p
Ca (hybrid):

sp3
Cb (ground state):

2s 2p
Cb (excited state):

2s 2p
Cb (hybrid):

2 2p

sp

Oa (ground state):

2s 2p
Oa (excited state):

2s 2p
Oa (hybrid state):

sp2 2p
Ob (ground state):

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2s 2p
Ob (excited state):

2s 2p
Ob (hybrid state):

3

sp
H (ground state):

1s

2 O sp2
a
sp

H 2p
sp2

H 1s  
sp2
1s 
sp3
 sp3 P

N 2p sp2 

sp3 sp3  sp2 C 1s
b H

C  3 O sp3
a b
sp
3 3 
sp
sp

3 sp3 sp3 sp3

 sp 

H 1s 1s H

5. a) Boiling point of ethanol > dimethyl ether.
• Ethanol can form Hydrogen Bond.
• dimethyl ether can only form van der Waals forces.
• Hydrogen Bond stronger than van der Waals forces.
• More energy needed to break the hydrogen bond between ethanol.

b) Boiling point of calcium is higher than that of potassium.
• Ca has 2 Valence electrons, while K has 1 Valence electrons.
• Strength of Metallic Bond between Ca atoms > between K atoms.
• More energy needed to break the metallic bond between Ca atoms.

c) ice floats on water.
• In ice, water is tetrahedrally bonded to other four water molecules to form open
hexagonal structure with many empty space between molecules.
• Thus, ice occupies larger volume than liquid.
• As a result, water is less dense than water.

d) Copper can easily be shaped into pipes and drawn into wires.
• In solid state, Cu atoms are arranged closely packed.
• When sufficient force is applied to the metal, one layer of atoms can slide over
another without disrupting the metallic bonding.
• As a result, metals are malleable and can be drawn into wires (ductile).

92