PRACTICE MODULE CHEMISTRY SK015

PRACTICE MODULE CHAPTER 2 SK015

ii) Calculate the wavelength and frequency of second line in Balmer series.
[ Ans : 4.862 x 10-7 m @ 486.2 nm ; 6.171 x 1014 s-1 ]

iii) Calculate the wavelength and frequency of first line in Pfund series.
[ Ans : 1.875 x 10-6 m @ 1875 nm ; 1.60 x 1014 s-1]

46

PRACTICE MODULE CHAPTER 2 SK015

7. a) i) Define ionization energy.

ii) Calculate the ionization energy of hydrogen atom in kJmol-1.
[ Ans : 1312 kJmol-1]

8. Calculate the minimum energy required to completely remove an electron from
Lyman series. [Ans : 1312 kJmol-1]

47

PRACTICE MODULE CHAPTER 2 SK015

9. A photon of light with wavelength 434 nm falls in the visible light region and forms a line
in the emission spectrum of hydrogen atom.
a) State the name of this emission spectrum series.
__________________________
b) Determine the transition of electrons that corresponding to the above wavelength.

c) Calculate the energy emitted for the transition [Ans : 4.58 x 10-19 J]

48

PRACTICE MODULE CHAPTER 2 SK015

10. The hydrogen emission spectrum in visible region is shown in the figure below :

a b cd

Given that line a is the first line in visible region,

a) Draw the energy level diagram to show the transition of electron for line b,c and d.

Energy

n=∞
n=8
n=7
n=6

n=5

n=4

n=3

n=2

n=1

b) Name the emission series that produce this series of lines.
_________________________

c) Calculate the wavelength corresponding to line a. [Ans : 656nm]

11. Give weaknesses of Bohr’s atomic model.

i)
_________________________________________________________________

ii)
________________________________________________________________

49

PRACTICE MODULE CHAPTER 2 SK015

iii) _______________________________________________________________
iv) _______________________________________________________________
2.2 Quantum number of electron
1. State and describe all the for quantum numbers on and electron?

2. What values of the angular momentum (ℓ) and magnetic quantum number (m) are
allowed for a principle quantum number (n) of 4?

3. Give all possible m values for orbitals that have given value of each of the following:
(a) ℓ = 3
(b) n=2
(c) n=6, ℓ=1

50

PRACTICE MODULE CHAPTER 2 SK015

4. Fill in the missing quantum numbers and subshell name:

n ℓm name
(a) 0 4p

(b) 2 10 2s
(c) 3 2 –2
(d)

5. Write an acceptable value for each of the missing quantum numbers
(a) n = 3, ℓ = ?, m = 2

(b) n = ?, ℓ = 2, m = –1
(c) n = 4, ℓ = 2, m = ?
(d) n = 1, ℓ = 0, m = ?

6. Are the following quantum number combinations allowed? If not, explain the reason.
(a) n = 1; ℓ = 0; m = 0
(b) n = 2; ℓ = 2; m = +1
(c) n = 7; ℓ = 1; m = +2
(d) n = 3; ℓ = 1; m = –1

7. For the following subshells give the values of the quantum numbers (n, ℓ, m) and the
number of orbitals in each subshell:
(a) 4p
(b) 3d
(c) 3s
(d) 5f

51

PRACTICE MODULE CHAPTER 2 SK015

8. How many orbitals in an atom can have each of the following designation:
(a) 5f
(b) 4p
(c) 5d
(d) n = 2

9. Write the four quantum number for an electron in a 6s orbital?
_____________________________________________________________________

10. Write the set of four quantum number for all electrons in a 5p orbital?

11. Which of the following are possible sets of quantum numbers for an electron?
For those sets that are NOT possible, give a reason.
(a) ( 1, 0, 1, ½ )
(b) ( 9, 7, –6, +½ )
(c) ( 2, 1, 0, 0 )
(d) ( 1, 1, 1, +½ )
(e) ( 3, 2, –3, +½ )
(f) (4, 0, 0, –½)

12. Sketch all the shape(s) of:
(a) s
(b) p
(c) d

Ans: (a)

52

PRACTICE MODULE CHAPTER 2 SK015

(b)

(c)

13. What is the similarity(s) and difference(s) between 1s and 2s orbital?
___________________________________________________________________
___________________________________________________________________

14. What is the similarity(s) and difference(s) between 3dxy and 3dxz orbital?
___________________________________________________________________
___________________________________________________________________

2.3 Electronic configuration of atom
1. State Aufbau principle, Hund’s Rule and Pauli’s exclusion principle.

2. For each of the following pairs of hydrogen orbitals, indicate which has higher energy:
(a) 1s , 2s
(b) 2p , 3p
(c) 3dxy , 3dyz
(d) 3s , 3d
53

PRACTICE MODULE CHAPTER 2 SK015

3. Which of the following pairs is lower of energy in electron atoms:
(a) 2s , 2p
(b) 3p , 3d
(c) 3s , 4s
(d) 4d , 5f
(e) 3d , 4s

4. Write the electronic configuration for each of the following atoms:
(a) Co (Z = 27) = ___________________________________________
(b) Mn (Z = 25) = ___________________________________________
(c) Ni (Z = 28) =___________________________________________
(d) Sc (Z = 21) = ___________________________________________
(e) Zn (Z = 30) = ___________________________________________
(f) V (Z = 23) = ___________________________________________

5. Write the electron configuration for each of the following atoms:
(a) Ge (Z = 32)
(b) Ar (Z = 18)
(c) Br ( Z = 35)
(d) Se (Z = 34)
(e) Kr (Z = 36)

6. What are the possible quantum numbers for the last (outermost) electron in Cl?

7. The electron configuration of a neutral atom is 1s2 2s2 2p6 3s2.Write a complete set of
quantum numbers for each of the electrons. Name the element.

54

PRACTICE MODULE CHAPTER 2 SK015

8. Draw the orbital diagram and write the electronic configuration for N, Cl and Ne.

9. Which of the following is the correct orbital diagram for the valence electronic
configuration of phosphosrus (Z = 15)?. If incorrect, give a reason.

10. Explain why each of the following ground state configuration is incorrect, and correct it:
55

PRACTICE MODULE CHAPTER 2 SK015

11. Write a set of quantum number for the third electron and a set for the eighth electron of
the F atom. (Z=9)

12. Write the electronic configurations of the following ions:
(a) Ca2+ (Z=20)
(b) N3– (Z=7)
(c) Br– (Z=35)
(d) Al3+ (Z=13)
(e) P3– (Z =15)
(f) Se2– (Z=34)
56

PRACTICE MODULE CHAPTER 2 SK015

13. Which of the following species has most unpaired electrons? S+, S, or S–. Explain your
reason for suggested answer.

14 Indicate the number of unpaired electrons present in each of the following atoms:
(a) B (Z = 5)
(b) Ne (Z = 10)
(c) P (Z = 15)
(d) Sc (Z = 21)
(e) Mn (Z = 25)
(f) Se (Z = 34)
(g) Kr (Z = 36)
(h) Fe (Z = 26)

57

PRACTICE MODULE CHAPTER 2 SK015

15. How many electrons are there in the third shell (n = 3)?
__________________________________________________________________

16. The element P (mass number = 39) and Q (mass number = 80) contain 20 and 45
neutrons respectively in their nucleus. Give their electronic configuration separately.

17. Chromium, Cr and Copper, Cu are in transition element in d-block of the Periodic
Table. The proton number of Chromium and Copper are 24 and 29 respectively.

a) Write the expected and actual electronic configuration of Chromium and Copper.

b) State the reason for the anomaly accured.

58

PRACTICE MODULE CHAPTER 2 SK015

OBJECTIVE QUESTION

1. When an electron excites or drops from its orbital to another orbital, energy is

A. Emitted only
B. Absorbed only
C. No effect
D. Emitted and absorbed

2. The orbits in which electrons move according to Bohr are

A. Elliptical
B. Cylindrical
C. Circular
D. Oval

3. When electron remains between orbits its momentum is

A. Quantized
B. Dequantized
C. Emitted
D. Changed always

4. The energy of each orbit is
A. Changed
B. Fixed
C. Not same
D. Effected

5. In Bohr’s model of the atom, the energy level closest to the nucleus would be the
A. Lowest energy level
B. Average energy level
C. Highest energy level
D. Valence energy level

6. What happens to the energy of an electron if it moves from one energy level to another
energy level farther away from the nucleus?

A. Loses energy
B. Becomes an ion

59

PRACTICE MODULE CHAPTER 2 SK015

C. Gains energy
D. Stays constant

7. When an an electron jumps from the fourth orbit to the second orbit, one gets the

A. Second line of Balmer series
B. First line of Pfund series
C. Second line of Lyman series
D. Second line of Paschen series

8. Which of the following series in the spectrum of hydrogen atom lies in the visible of the
electromagnetic spectrum?

A. Balmer series
B. Brackett series
C. Lyman series
D. Paschen series

9. Which quantum numbers gives the shell to which the electron belongs?

A. n
B. l
C. m
D. s

10. What is the range of Azimuthal Quantum Number, l?

A. 0 to n
B. 0 to s
C. 0 to n-1
D. 0 to s-1

11. Which of the following can be the quantum numbers for an orbital?

A. n = 4, l = 4, m = 3
B. n = 2, l = 3, m = 1
C. n = 3, l = 2, m = -1
D. n = 3, l = 0, m = -3

12. Which of the following quantum number gives the shape of atomic orbital of sub-shell?

A. n
B. l
C. m
D. s

60

PRACTICE MODULE CHAPTER 2 SK015

13. Which of the following sets of quantum numbers is not allowed?
A. n = 3, l = 1, m = -1
B. n = 2, l = 0, m = 0
C. n = 3, l = 2, m = -3
D. n = 2, l = l, m = 0

14. The maximum number of electrons that can be accommodated in a subshell l = 3 is
A. 2
B. 10
C. 6
D. 14

PAST YEAR QUESTION
Session 2018/2019 (PSPM)
1. a) The wavelength that produces a line, B in Brackett series is 2165.6 nm.

i) Determine the transition that forms the B line
ii) Another line, C was formed with wavelength of 1817.5 nm. Explain qualitatively, whether

line B or line C, has the higher energy emitted.

b) The electronic configuration of element D is 1s2 2s2 2p6 3s2 3p3.
Draw and label the 3D shape of orbitals occupied by the valence electrons.

Session 2019/2020 (PSPM)
a) The line spectrum of hydrogen in the visible region in shown in the following diagram

Line
X

Frequency increases
i) Name the series of the line spectrum.
ii) Sketch an energy level diagram of a hydrogen atom for the formation of line X. Explain

how line X is formed.

61

PRACTICE MODULE CHAPTER 2 SK015

iii) Calculate the wavelength corresponding to line X.
b) Describe the anomalous electronic configuration of chromium atom.

Prepared by : Mdm Noor Najihah & Miss Noorazreen

62

PRACTICE MODULE CHAPTER 3 SK015

CHAPTER 3.0 : PERIODIC TABLE

MIND MAP
GENERAL OVERVIEW

Period

Classification of group
Elements
Periodic block across
Table Periodicity period
atomic
radii down the
group

1st row transition
element

isoelectronics 1st IE
ionization energy 2nd IE

ionic radii period 2 &3

acid base group 3

electronegativity

WORKSHEET 3.1

3.1 Classification of element

1. How many inner, outer, and valence electrons are present in atom of the following

elements?

(a) O (Z = 8) (f) Br (Z = 35)

(b) Ga (Z = 31) (g) Al (Z = 13)

(c) Ca (Z = 20) (h) Cr (Z = 24)

(d) Fe (Z = 26) (i) P (Z = 15)

(e) Se (Z = 34) (j) F (Z = 9)

Ans: Answer
Elements
a. O Inner e Outer e Valence e
b. Ga
c. Ca

63

PRACTICE MODULE CHAPTER 3 SK015

d. Fe
e. Se
f. Br
g. Al
h. Cr
i. P
j. F

2. Identify the group and period of each element with valence electronic configuration:
(a) ns2 np4
(b) ns2
(c) ns2 np6
(d) ns1
(e) ns2 (n − 1)d1−10

3. Without referring to periodic table, classify the following elements according to their

group and block in the periodic table.

(a) K (Z = 19)

(b) Cl (Z = 17)
€ Ti (Z = 22)
(d) Kr (Z = 36)
€ Co (Z = 27)
(f) Si (Z = 14)

(g) Ca (Z = 20)

(h) Ga (Z = 31)

64

PRACTICE MODULE CHAPTER 3 SK015
4. Choose a pair of the following electronic configuration that would represent similar

chemical properties of their atoms
(a) 1s2 2s2 2p5
(b) 1s2 2s1
(c) 1s2 2s2 2p6
(d) 1s2 2s2 2p6 3s2 3p5
(e) 1s2 2s2 2p6 3s2 3p6 4s1
(f) 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6

5. From each partial orbital diagram given below, write the electron configuration and state
its group number:

3.2 Periodicity
WORKSHEET 3.2.1
1. By referring to the periodic table, arrange the elements in each set in order of increasing

atomic radius :
(a) Rb , K , Cs
(b) C , O , Be
(c) Mg , K , Ca
(d) Ge , Pb , Sn
(e) F , Ne , Na
(f) Be , Mg , Na

65

PRACTICE MODULE CHAPTER 3 SK015

2. By referring to the Periodic Table, rank the elements in order of decreasing atomic radii:
(a) Ca , Mg , Sr
(b) K , Ga , Ca
(c) Br , Rb , Kr
(d) Sr , Ca , Rb
(e) Se , Br , Cl
(f) I , Xe , Ba

3. Determine which species is smaller size and explain in details.
(a) Na or Na+
(b) Cl or Cl–

66

PRACTICE MODULE CHAPTER 3 SK015
4. Determine which species is larger size and explain in details.

(a) Al or Al3+
(b) S or S2–

WORKSHEET 3.2.2
1. Define the electronegativity.

_____________________________________________________________________
2. Compare the electronegativity of each pair of atoms given, choose which element has

larger electronegativity
(a) S and O
(b) Mg and P.
(c) I and Br
(d) P and Si
(e) Cl and P
(f) I and F

67

PRACTICE MODULE CHAPTER 3 SK015
3. By refer to the periodic table, rank the elements in each of the following sets in

order of decreasing ionization energy:
(a) Kr , He , Ar
(b) K , Na , Rb
(c) Sr , Ca , Ba
(d) N , B , Ne
(e) As , Sb , P
(f) Na , Li , K
(g) Be , F , C
(h) Cl , Ar , Na
(i) Cl , Br , F

4. Based on the electronic configurations shown, answer all the questions.
X = [Ar] 3d8 4s2 Z = [Ar] 3d10 4s2 4p5

(a) An atom of which element is expected to have the larger ionization energy?
(b) Which atom is smaller ?

5. Given the electronic configuration for Element X and Z, answer these questions below.
X = [Kr] 4d10 5s1 Z = [Ar] 3d10 4s2 4p4
(a) Is the element X a metal or nonmetal?
(b) Which element has the larger atomic radius?
(c) Which element would have the greater first ionization energy?

68

PRACTICE MODULE CHAPTER 3 SK015
WORKSHEET 3.2.3
1. In general, ionization energy increases from left to right across a given period.
Boron (B), however, has lower ionization energy than beryllium (Be). Explain.

2. Why F has a larger first ionization than atom O?

3. Write the symbol and the electronic configuration of the element decribed.
(a) The metal in Group 1 whose atom is the smallest.
_____________
(b) The alkaline earth metal with the heaviest atom.
___________________________
(c) The transition metal in Period 4 whose atom contains the fewest protons.
___________________________
(d) The largest metalloid atom with the maximum number of unpaired p electrons.
___________________________
(e) An element in Period 4 whose +2 ion is isoelectronics with argon.
____________________________

69

PRACTICE MODULE CHAPTER 3 SK015

WORKSHEET 3.3

OBJECTIVE QUESTION

1. Which of the following is an inert gas?
A. Oxygen
B. Nitrogen
C. Argon
D. Hydrogen

2. The arrangement of elements in the modern periodic table is based on their
A. electronic configuration.
B. atomic size.
C. atomic number.
D. atomic mass.

3. Which element is more electronegative among halogens?
A. Br
B. Cl
C. F
D. I

4. Why does atomic size increase down a group?
A. Due to the addition of more electrons
B. Due to decrease in number of protons
C. Due to the addition of an extra shell
D. All of the above

5. Element 'X' forms a chloride with the formula XCl2, which is a solid with high melting point.
Identify X.
A. Na
B. Al
C. Mg
D. Si

6. Which of the following pairs of elements have the same number of valence electrons?
A. Zn and O
B. Cl and S
C. Na and Ca
D. C and Ge

7. How many electrons are there in the outermost orbit of each element in group 2?
A. 1
B. 2
C. 3
D. 7

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PRACTICE MODULE CHAPTER 3 SK015

8. Identify the period which contains s, p, d block elements.
A. 2nd period
B. 3rd period
C. 4th period
D. 1st period

9. Which element contains half-filled 'd' orbital?
A. Vanadium
B. Zinc
C. Chromium
D. Titanium

10. An element having atomic number 34 can be placed in which of the given group and period
of the modern periodic table?
A. Group 15 and 4th period
B. Group 16 and 5th period
C. Group 14 and 7th period
D. Group 16 and 4th period

11. Elements in the same vertical group of the periodic table have same
A. Number of valence electrons
B. Atomic number
C. Atomic mass
D. Atomic volume

12. Which set of elements is listed in order of increasing ionization energy?
A. Sb < As < S < P < Cl
B. Cl < Sb < P < As < S
C. As < Cl < P < S < Sb
D. Sb < As < Cl < S < P

13. In which of the following pairs are elements belonging to the same group?
A. Boron & Beryllium
B. Nitrogen & Phosphorous
C. Magnesium & Aluminium
D. Gallium & Helium

14. An element has electronic configuration 1s2 2s2 2p2. It belongs to
A. Group 2
B. Group 4
C. Group 14
D. Group 12

15. Which comparison of ionization energy is correct
A. Mg < Al
B. Si > P
C. Mg > Al
D. both b & c

71

PRACTICE MODULE CHAPTER 3 SK015

16. Shielding effect across the period
A. Increases
B. Decreases
C. Cannot be predicted
D. Remains constant

17. Oxidation state of an atom represents
A. Number of electrons gained
B. Number of electrons lost
C. Apparent charge in compound
D. Its vacancies

18. Keeping in view the size of atom which order is correct one?
A. Mg > Sr
B. Ba > Mg
C. Li > Cs
D. Cl > I

19. On moving across a period from left to right the ionisation energy increases because
A. value of principal quantum number increases
B. effective nuclear charge increases
C. atomic size increases
D. nuclear charge increases

20. The nuclear charge in periodic table….
A. Across a period.
B. Down a group.
C. d-block.
D. s-block.

21. Which of the following describes the physical properties of a transition element ?

Melting points /0C Boiling points/0C Density/ gcm-3 Ability to conduct

electricity

A. 323 1750 11.8 Good
B. 660 1150 1.7 Good
C. 1543 2850 7.7 Good
D. 1411 2500 2.3 weak

22. Mendelev arranged elements in the Periodic Table according to

A. Increasing order of proton numbers
B. Increasing order of relative atomic mass
C. Increasing order of neutron numbers
D. Increasing order of atomic radius

72

PRACTICE MODULE CHAPTER 3 SK015

23. Across the second period ( Lithium to fluorine) in the Periodic Table, it was discovered that
the

A. Electronegativity of elements decreases.
B. Ionization energy of elements throughout the series.
C. Atomic radius of elements decreases.
D. Strength of elements as reducing agents increases.

24. The outermost electronic configuration of elements P and Q are ns2 np3 and ns2 np5
respectively. Which of the following pairs represents the formula and the type of bonds formed
by the elements?

A. PQ3 , covalent
B. PQ3 , ionic
C. PQ5 , ionic
D. P3Q, covalent

25. Which of the following set of diagrams best represents the relative sizes of the atom for
sodium and chlorine ?

Sodium Chlorine

Atom Ion Atom Ion

A.

B.

C.
D.

26. Which of the following could be the strongest reducing agent ?
A. Br-
B. F-
C. Li+
D. Al3+

73

PRACTICE MODULE CHAPTER 3 SK015

27. An element M forms positive ion and has one chloride only. Which of the following shows
the electronic configurations of M ?

A. 1s2 2s2 2p6 3s2
B. 1s2 2s2 2p4
C. 1s2 2s2 2p6
D. 1s2 2s2 2p6 3s2 3p6 4s1 3d10

28. X and Y are two elements in the Periodic Table. The proton number of X is 26, and Y2+ ion
is isoelectronic with X, which of the following is correct about X and Y ?

Element /Ion Electronic configuration
1s2 2s2 2p6 3s2 3p6 4s2 3d4
A. X 1s2 2s2 2p6 3s2 3p6 4s2 3d10
1s2 2s2 2p6 3s2 3p6 3d8
B. Y 1s2 2s2 2p6 3s2 3p6 4s2 3d6
C. Y2+
D. Y2+

29. Which of the following elements is expected to form the largest ion with a noble gas
configuration ?

A. Nitrogen
B. Phosphorus
C. Chlorine
D. Calcium

30. Where would you find the most reactive metals on the periodic table?

A. Group 1
B. Group 13
C. Group 2
D. Group 18

Prepared by : Mdm Noor Asmahan Abdullah

74

PRACTICE MODULE CHAPTER 4 SK015
CHAPTER 4: CHEMICAL BONDING
Mind map /I-Think Map

4.1 LEWIS STRUCTURE

LEWIS Stable FORMAL EXCEPTION RESONANCE
DOT Configuration CHARGE OF OCTET
SYMBOL 2 @ more
OCTET of ion lewis structure
FORMATION RULE
OF for a single
LEWIS FORMULA: MOST molecule that
- IONIC STRUCTURE no of PLAUSIBLE
- COVALENT valence STRUCTURE cannot be
- DATIVE @ WITH electron – represented
COORDINATE [no of lone - INCOMPLETE accurately by
- SINGLE pair + ½ of OCTET only one lewis
- DOUBLE bonding
- TRIPLE electron] - EXPANDED structure.
OCTET
BOND
- ODD NO.
ELECTRONS

- Noble gas
configuration

- Pseudo noble gas
configuration

- Half filled orbital

75

4.2 MOLECULAR SHAPE & PRACTICE MODULE CHAPTER 4 SK015
POLARITY
4.3 ORBITAL OVERLAPPING &
Definition HYBRIDISATION

5 basic Sigma & pi
shape bond

Bond polarity VSEPR THEORY Hybrid of
& dipole central atom
moment (Valence Shell Electron Pair
Repulsion Theory)

Polarity of molecule based Illustrate hybridisation &
on shape & resultant overlapping of orbital
dipole moment

76

PRACTICE MODULE CHAPTER 4 SK015

Dipole-dipole Factors that influence
forces (DD) vdW & HB

van der Waals • Molecular
Forces polarity

London force • Molecular
(LF) @ shape

Dispersion • Molecular size
Forces (polarisability)

4.4
INTERMOLECULAR

FORCES

Effect on physical
properties

Hydrogen • Boiling point • Number of HB
bonding (HB) • Solubility • Electronegativity
• Density of
of element
water compare
to ice

77

PRACTICE MODULE CHAPTER 4 SK015

Formation

Relate metallic • Malleability
bond to • Ductility
• Electrical
properties of
metal conductivity using
band theory.
• Thermal conductivity

4.5
METALLIC

BOND

Factors affects • Numbers of valence
strength of metallic electron

bond. • Size of atoms

Relate • Molecular structure Element of:
boiling/melting • Type of bonding
• Intermolecular force • Period 3
point to • Group 1
• Group 17

78

PRACTICE MODULE CHAPTER 4 SK015

4.1 LEWIS STRUCTURE

WORKSHEET 4.1

1. State the type of bonding (ionic or covalent) in the following compound:

(a) CsF(s) (f) O3(g)
(b) N2(g) (g) MgCl2(s)
(c) ICl3(g) (h) BrO2(g)
(d) N2O(g) (i) H2S(g)
(e) LiCl(s) (j) CaO(s)

2. a) State the octet rule.

__________________________________________________________________

b) write the electronic configuration for the following ions and state the type of
stability.
i. 30Zn2+
ii. 35Br-
iii. 25Mn2+
iv. 16S2-
v. 7N3-
vi. 31Ga3+
vii. 26Fe3+

79

PRACTICE MODULE CHAPTER 4 SK015

3. In which of the following bonding patterns does X obey the octet rule?

4. By using Lewis symbols show the formation of the following ions from the atoms, and
determine the formula of the compound:
(a) Na+ and O2–
(b) Na+ and Br–
(c) Mg2+ and O2–

5. (a) Describe the formation of dative bond by using NH4+, H3O+, Al2Cl6 and BF3NH3 as
examples.

80

PRACTICE MODULE CHAPTER 4 SK015

6. Write a Lewis structure for:

(a) ICl
(b) Br2
(c) H2O
(d) SF2
(e) H2S
(f) OF2
(g) PCl3
(h) CO2
(i) CS2
(j) C2H4

7. Draw the Lewis structure of the:
(a) CO32–
(b) ClO2–
(c) ClO4–
(d) SO32–
(e) SO42–

81

PRACTICE MODULE CHAPTER 4 SK015

8. Write Lewis structures for the following structures:
(a) NO-
(b) CN-
(c) H3O+
(d) NH4+
(e) NO3-

9. Assign formal charges to each of the atoms in the following structures.

82

PRACTICE MODULE CHAPTER 4 SK015

10. Draw a Lewis structure and calculate the formal charge of each atom in:
(a) AlH4–
(b) SCO
(c) CN–
(d) ClO–
(e) BF4–
(f) CO
(g) BrCl3

11. Write the Lewis structure of BF3 that follow the octet rule and not obey octet rule. Calculate
the formal charge of each atom. Which one is most plausible? Explain.

83

PRACTICE MODULE CHAPTER 4 SK015

12 Based on the species below, Write the Lewis structure that follow the octet rule and not
obey octet rule. Calculate the formal charge of each atom. Which one is most plausible?
Explain.
(a) BeF2
(b) CO2

13. The following species do not obey the octet rule. Draw a Lewis structure for each, and
state the type of octet rule exception:
(a) NO2
(b) ICl3
(c) AlCl3
(d) BrF5
(e) SO3

14. (a) What is meant by resonance structures?
_________________________________________________________________________
(b) Draw all the resonance structures of
i. N3-.
ii. NCO-
iii. NO3-
iv. SO3
v. N2O

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PRACTICE MODULE CHAPTER 4 SK015

4.2 MOLECULAR SHAPE AND POLARITY

WORKSHEET 4.2
1. Explain Valence Shell Electron Pair Repulsion theory (VSEPR).
2. Give all the basic molecular shapes and their bond angle.

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PRACTICE MODULE CHAPTER 4 SK015

3. Determine the number of e- groups around the central atom of the molecular shape below:

No of electron group Arrangement

Linear

Trigonal Planar

Tetrahedral

Trigonal bipyramidal

octahedral

4. Predict and explain the shape of molecular shape and bond angle in the given species.
(a) COCl2
(b) CS2
(c) CBr4
(d) SbF5
(e) SeF6

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PRACTICE MODULE CHAPTER 4 SK015

5. Draw and determine the shape of molecules for following species.

Compund Molecular structure

(a) CO2

(a) NH4+

(b) NO3-

(c) SiH4

(d) BF3

(e) CH2O

(f) CO32-

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PRACTICE MODULE CHAPTER 4 SK015

6. Show the polarity of each bond with polar arrows:
(a) N–B
(b) N–O
(c) S–O
(d) N–H
(e) Cl–O

7. Chloroform (CHCl3) is a polar compound but carbon tetrachloride (CCl4) is nonpolar.
Why? Draw the molecular shape and indicate the bond polarity.

8. Explain the polarity of PCl2F3.

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PRACTICE MODULE CHAPTER 4 SK015

9. Which of the following characteristics apply to BH3?
1.nonpolar molecule
2.polar bonds
3.trigonal-pyramidal molecular geometry
_________________________________________________________________

10. Predict whether each of the following molecules is polar and show the direction of bond
polarity and net dipole moment.
(a) Boron trichloride, BCl3
(b) Carbonyl sulfide, SCO
(c) Dichloromethane, CH2Cl2
(d) Chlorine, Cl2
(e) Hydrogen Chloride, HCl

12. Determine whether the following molecules are polar or nonpolar:
(a) CO2
(b) H2O
(c) SO3
(d) CCl4
(e) CHCl3
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PRACTICE MODULE CHAPTER 4 SK015

4.3 ORBITAL OVERLAP AND HYBRIDISATION
WORKSHEET 4.3

1. Describe the formation of sigma and pi bond.

2. Determine the hybridization of the central atom in each of the following compound.

(a) H2O (h) ClO2–
(b) H2S (i) ClO4–
(c) NO3- (j) SO32–
(d) CO2 (k) SO42–

(e) SeF6 (l) CS2

(f) SbF6 (m) COCl2

(g) BrF5 (n) BrCl3

3 Indicate the type of hybridisation of carbon and nitrogen atoms in the molecule below.

HH N
HN CN C C

HN H H

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PRACTICE MODULE CHAPTER 4 SK015

4. Used the hybridisation theory to explain the formation of following molecules and draw the
orbital overlapping.
i. O2
ii. ClO3-

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PRACTICE MODULE CHAPTER 4 SK015

5. Xenon difluoride is a powerful fluorinating agent with the chemical formula XeF2, and one of
the most stable Xenon compounds.
i. Draw the Lewis structures of xenon difluoride

ii. State the type of octet rule exception.
_______________________________________

iii. Determine the electron pair geometry.
__________________________________________________________

iv. Determine the molecule shape.
___________________________

v. Determine the polarity of the molecule and explain
vi. State the type of the hybridization of the central atom in xenon difluoride

molecule
vii. Describe the hybridization process in xenon difluoride.

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PRACTICE MODULE CHAPTER 4 SK015

6 Consider the reaction
BF3 + NH3 → F3B-NH3 .

Describe the changes in hybridization (if any) of the B and N atoms as a result of this
reaction.

4.4: INTERMOLECULAR FORCES
WORKSHEET 4.4

1. Describe types of intermolecular forces and give example for each.

2. Determine the type of intermolecular forces exist in the following molecules:

ANSWER Molecules intermolecular forces
(a) O2

(b) HCl

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PRACTICE MODULE CHAPTER 4 SK015

(c) H2O
(d) HF

3. 1-Propanol C3H7OH and methoxyethane CH3 –O-C2H5 have the same molecular weight.
Which has the higher boiling point?

4. Why do the lightest compounds such as NH3, H2O, and HF have much higher boiling points
than expected?

5. Explain why ice is less dense than water although they have same molecular formula.

4.5: METALLIC BOND
WORKSHEET 4.5

1. a. Define metallic bond.

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PRACTICE MODULE CHAPTER 4 SK015

b. Explain the formation of ionic bond in Mg and Ca with using electron sea model.

2. State 4 properties of metal and explain. Explaination
ANSWER Properties

State 2 factors that’s affecting the strength of metallic bond and explain.
3.

Factor Explaination.

Which element, aluminium or magnesium has a higher melting point? Explain your answer.
4.

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