PRACTICE MODULE CHEMISTRY SK015

PRACTICE MODULE CHAPTER 4 SK015

5. State and explain the trend in boiling point on descending Group 1 of the periodic table.

WORKSHEET 4.6
MULTIPLE CHOICE QUESTIONS
1. How many electrons should carbon have around its Lewis dot model?

A. 1
B. 3
C. 4
D. 5
2. According to the octet rule most elements need _______ valence electrons.
A. 2
B. 8
C. 6
D. 18
3. When electrons are shared unequally, chemists characterize these types of bonds as_______.
A. polar covalent
B. ionic
C. pure covalent
D. nonpolar covalent

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PRACTICE MODULE CHAPTER 4 SK015

4. Which of the following terms describes electrons or pairs of electrons not involved in chemical
bonding?
A. Bonding electrons
B. Lone-pair electrons
C. Valence electrons
D. Core electrons

5. A lone pair is defined as.
A. A pair of bonding electrons
B. One non-bonding electron
C. A pair of non-bonding electrons
D. A pair of electrons on the central atom

6. Atoms with greatly differing electronegativity values are expected to form
A. polar covalent bonds
B. nonpolar covalent bonds
C. triple bonds
D. ionic bonds

7. The attraction of two ions due to opposite charge is known as..
A. ionic bonding
B. covalent bonding
C. metallic bonding
D. dative bonding

8. Which of the following pairs of elements would form an ionic bond?
A. K and Ca
B. N and O
C. He and Al
D. P and Li

9. Which one of the following pairs atoms is most likely to form an ionic bond?
A. Na and F
B. C and F
C. N and F
D. O and F

10. Which is a correct Lewis structure for carbon dioxide, CO2?
A.
B.
C.
D.

11. Formal charge is _______________________________
A. the absolute value of the charge on a polyatomic anion or cation.
B. the difference between the number of lone pairs of electrons and shared pairs of electrons on
any atom in a Lewis structure.

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PRACTICE MODULE CHAPTER 4 SK015

C. equal to the number of valence electrons in a free atom minus the number of shared in
covalent bonds.

D. the difference between the number of valence electrons in a free atom and the number of
electrons assigned to the atom in a Lewis structure.

12. In which structure(s) below does the oxygen have a formal charge of +1?

A. I only
B. II only
C. I and IV
D. I, III, and IV
13. VSEPR stands for..
A. Valid Shell Electricity Program Remedy
B. Valence Shell Electron Pair Repulsion
C. Velocity Sustaining Electron Proton Review
D. Valence Shell Electricity Proximal Repulsion

14. The concept that electron pairs located in the valence shell of an atom bonded to other atoms tend to
stay as far apart as possible so as to minimize repulsions between them is incorporated in the..
A. Pauli principle
B. valence shell electron pair repulsion theory
C. electronegativity and polar bonds theory
D. Aufbau principle

15. Which of the following is the correct order for the electron pair repulsions?
A. lone pair-lone pair < bonding pair-bonding pair < bonding pair-lone pair
B. lone pair-lone pair < bonding pair-lone pair < bonding pair-bonding pair
C. bonding pair-bonding pair < bonding pair-lone pair < lone pair-lone pair
D. bonding pair-bonding pair < lone pair-lone pair < bonding pair-lone pair

16. The shape of a molecule with six bonding pairs and no lone pairs is...
A. hexahedral
B. octahedral
C. tetrahedral
D. trigonal bipyramidal

17. Choose the correct name of the molecular geometry for the given Lewis Structure?

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PRACTICE MODULE CHAPTER 4 SK015

A. trigonal planar
B. trigonal pyramidal
C. tetrahedral
D. bent

18. The shape of carbon dioxide is described as...
A. linear
B. octahedral
C. tetrahedral
D. trigonal planar

19. Bonding pairs of electrons in SF6 are….
A. 2
B. 4
C. 6
D. 8

20. The bond angles in PF5 are...
A. all 72°
B. 90° and 120°
C. 109.5° and 120°
D. 109.5° and 90°

21. Which substance would be polar?
A. O2
B. HCl
C. CO2
D. CH4

22. Which covalent bond is the most polar?
A. F-F
B. N-F
C. Cl-F
D. C-F

23. Which has a dipole moment?
A. CO32-
B. SO42-
C. SO2
D. CO2

24. Which substance would contain the strongest intermolecular force?
A. O2
B. H3N
C. CO2
D. CH4

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PRACTICE MODULE CHAPTER 4 SK015

25. Which one of the following represents the weakest interaction between two species?
A. Hydrogen bond
B. Dipole-dipole force
C. Ionic bond
D. Dispersion force

26. Which species does not contain sp3-hybridized atom?
A. BH3
B. BH4-
C. NH3
D. NH4+

27. Which set of hybridization states of C1, C2, and C3 of the following molecule is correct?

A. sp2, sp2, sp2
B. sp2, sp2, sp
C. sp3, sp2, sp
D. sp3, sp2, sp2

28. A  (pi) bond is present in molecule. This bond is formed when
A. 2s orbitals overlap
B. Side overlapping of two parallel p orbitals
C. Overlapping of two sp orbitals
D. Overlapping of two sp2 orbitals

29. A triple bond contains ___ sigma bond(s) and ___ pi bond(s).
A. 0, 3
B. 3, 0
C. 2, 1
D. 1, 2

30. The perchloric acid molecule, HClO4 contains:

A. 13 lone pairs, 1 bond, and 4 bonds.
B. 8 lone pairs, 2 bonds, and 7 bonds.
C. 2 lone pairs, 3 bonds, and 4 bonds.
D. 11 lone pairs, no bonds, and 5 bonds.

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PRACTICE MODULE CHAPTER 4 SK015

31. A neutral molecule having the general formula AB3 has two unshared pair of electrons on A. What
is the hybridization of A?

A. sp
B. sp2
C. sp3
D. sp3d

32. Which of the following pairs of molecules and their molecular geometries is WRONG?

A. NF3 - trigonal planar
B. H2O – bent
C. BF3 - trigonal planar
D. AsF5 - trigonal bipyramidal

33. Which substance would form a double bond when bonding covalently?
A. water
B. oxygen
C. nitrogen
D. fluorine

34. The boiling point of water is exceptionally high because___________
A. there is a covalent bond between H and O
B. water molecule is linear
C. water molecules associate due to hydrogen bonding
D. water molecules are not linear

34. London forces exist_______
A. for all molecules.
B. only for molecules with nonpolar bonds.
C. only for molecules with polar bonds.
D. only for molecules with metallic bonds.

35. The correct order of increasing attractive strength for weak intermolecular forces is...
A. dipole-dipole forces, hydrogen bonding, London forces
B. London forces, dipole-dipole forces, hydrogen bonding
C. hydrogen bonding, dipole-dipole forces, London forces
D. hydrogen bonding, London forces, dipole-dipole forces.

36. Which one of the following exhibits intermolecular hydrogen bonding?
A. HF
B. HCl
C. HBr
D. HI

37. Which of the following intermolecular forces results from the attraction between temporary dipoles
and their induced temporary dipoles?
A. Ionic bond
B. Hydrogen bond
C. London forces

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PRACTICE MODULE CHAPTER 4 SK015

D. Van der Waals forces.

38. Which of these is not an intermolecular force?
A. covalent bonding
B. hydrogen bonding
C. London forces
D. dipole-dipole forces

39. What does malleable mean?
A. able to be shaped
B. will break easily
C. can be used for wire
D. is shiny

40. Metallic bonding is...
A. a type of covalent bond.
B. a type of ionic bond.
C. an attraction between positive and negative ions.
D. an attraction between positive ions and electrons.

WORKSHEET 4.7

MODEL OF PSPM QUESTIONS
1. Antimony (Sb) is in the Period 5 and Group 15 of the periodic table.

a) Give the valence shell electronic configuration of antimony.
b) Based on its valence shell electronic configuration predict the possible formulae of two

fluorides of antimony.
c) Draw the Lewis structure and predict the shapes of the two fluorides of antimony.
d) State the polarity of the two compounds.
e) State the hybridization of antimony as the central atom and show the overlapping
orbitals.

[17 marks]
2. a) Two possible Lewis structures of phosgene, COCl2 are

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PRACTICE MODULE CHAPTER 4 SK015

O Cl
Cl C Cl O C Cl

structure A structure B

Determine which structure is more plausible and explain. [4 marks ]

b) Consider the species below NH2- ion and NH3 molecule

i. Determine the molecular geometry of both species

ii. Using VSEPR theory, explain the difference of H-N-H angle in NH2 and NH3
[10 marks]

c) Mg metal is easily deformed by an applied force, whereas magnesium fluoride is
shattered. Explain why do these things behave so differently. [3 marks]

Prepared by : Mdm Julianawati & Mdm Wan Rosilah

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PRACTICE MODULE CHAPTER 4 SK015

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PRACTICE MODULE CHAPTER 5 SK015

CHAPTER 5: STATES OF MATTER

5.1: GAS

Basic kinetic molecular theory of gases for an ideal gas

Boyle’s law
V α 1/P (constant T and n)

P1V1 = P2V2

Gas laws

GAS

Charles’s law
V α T (at constant P and n)

Avogadro’s law
V α n (at constant P and T)

Ideal gas equation: PV = nRT

Dalton’s law: PTotal = PA + PB + PC + …

Ideal and non-ideal gas behavior (in term of intermolecular forces &
molecular volume)

Condition at which real gas approach to ideal gas:
at very low pressure (constant temperature)
at high temperature (at constant volume)

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PRACTICE MODULE CHAPTER 5 SK015

WORKSHEET 1
Kinetic Molecular Theory of Gases
1. Write four assumptions of the kinetic molecular theory of gases.

Basic Assumptions of Kinetic Molecular Theory

12

________________________________ ___________________________________
________________________________ ___________________________________

3 4

___________________________________ _______________________________________
___________________________________ _______________________________________

2. Describe the effects on gas molecules in a closed container when
temperature is increased

volume is decreased

Gas Law

3. State the gas law by relating the variables P, V, T, and n.

Boyle’s law Charles’s law Avogadro’s law

4. What is absolute zero?
______________________________________________________________________________

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PRACTICE MODULE CHAPTER 5 SK015

5. Sketch and interpret the graph.

Boyle’s law V against P V against 1 / P PV against V @ P

Interpretatio V against T(K) V against T(°C)
n

Charles’s
law

Interpretatio
n

Conversion Pressure Unit

6. Convert the pressure value to atmospheric pressure.

680 mmHg 680 Pa

680 torr 68 kPa

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PRACTICE MODULE CHAPTER 5 SK015

Applying Gas Laws

7. 100 cm3 of hydrocarbon at 1 atm is compressed to 37.0 cm3 at constant temperature. What is the new

pressure? [ Ans : 2.70 atm]

Information Given Conversion

Target
Calculation

8. A sample of gas at 2.00 atm and 30°C is heated at constant pressure to 50.0°C. Its final volume is 5.0
L. What was its original volume? [Ans :4.7L]

Information Given Conversion

Target
Calculation

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PRACTICE MODULE CHAPTER 5 SK015

9. A gas is heated from −10.0°C to 50.0°C at 1 atm and simultaneously compressed to one half of its
original volume. What will be the final pressure in mmHg? [Ans : 1.87 x 10-3 mmHg]

Information Given Conversion

Target
Calculation

10.Calculate the volume (in liters) at STP occupied by 0.40 mol O2 gas.[Ans :8.96L]

Information Given Conversion

Target
Calculation

Ideal gas Equation
11.What is an ideal gas?

_________________________________________________________________________________
_________________________________________________________________________________

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PRACTICE MODULE CHAPTER 5 SK015

12.What is the formula for an ideal gas equation?
_________________________________________________________________________________

Relationship of Ideal Gas Equation with Mass, Molar Mass and Density of Gas

13. Fill in the blank molar mass density
Ideal gas equation

Applying Ideal Gas Equation.
14.What volume will 2.20 g of carbon dioxide occupy at 30.0°C and 755 torr. [Ans :1.25L]

Information Given Conversion

Target
Calculation

15. The pressure of a sample of gas collected in a 0.340 L gas bulb at temperature of 23.0°C is 0.750 atm.

What is the molecular mass of the gas if the sample weighed 0.478 g to begin with?[ Ans: 45.5 amu]

Information Given Conversion

Target

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PRACTICE MODULE CHAPTER 5 SK015

Calculation

16. What is the density (in g L-1) of oxygen at 742 mmHg and 25.0°C? [Ans : 1.28gL-1]

Information Given Conversion

Target
Calculation

Dalton’s Law of Partial Pressure

17.Define Dalton’s law of partial pressure.
_______________________________________________________________________________
_______________________________________________________________________________

18.Fill in the blank.

At constant 1 mol O2
volume and 1 mol He
temperature
P total
1 mol 1 mol
He O2

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PRACTICE MODULE CHAPTER 5 SK015

19. What is the volume of the gaseous before and after the valve of the container is opened?

Argon Neon

V container = 2.5 L V container = 3.5 L

Gas Volume before Volume after
Argon, Ar
Neon, Ne

Applying Dalton’s Law of Partial Pressure

Mole fraction, X A = Mole of gas A(nA ) (nT )
Total number of
moles

Sum of mole fraction of all gases in the mixtures always equal to1

XA + XB + XC +... =1

Partial pressure,PA = XA .PT
PT = PA + PB + PC +...
Where PT = total pressure of a mixture of non − reacting gases

20.A mixture of gases at 1 atm contains 50% nitrogen, 20% oxygen and 30% neon by volume. What is

the partial pressure of each gas? [Ans : PNe = 0.30 atm, PO2 = 0.20 atm, PN2= 0.50 atm]

Information Given Conversion

Target

Calculation

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PRACTICE MODULE CHAPTER 5 SK015

21. 100 mL of argon at 25°C and 700 torr is mixed with oxygen and transferred to a 600 mL container at

the same temperature. The total pressure of the mixture is 400 torr. What is the partial pressure of

argon and nitrogen? [PAr = 1.17X 102 torr, PO2= 283 torr]

Information Given Conversion

Target
Calculation

Collection of Gas by Displacement of Water

P total = P gas + P water vapor
22. The decomposition of potassium chlorate to potassium chloride and oxygen gas is shown as

2KClO3(s) → 2KCl(s) + 3O2(g)
A sample of 3.10 L of oxygen gas is collected over water with the total pressure of 745 torr at 25.0°C.
Calculate the mass of oxygen gas collected. [Vapor pressure of water at 25°C=23.8 torr] [Ans : 3.84g]

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PRACTICE MODULE CHAPTER 5 SK015

Information Given Conversion
Target

Calculation

Ideal and non-ideal behaviors of gases
23.State the conditions at which a gas deviate from ideal behavior. Explain your answer.

Conditions at which a gas deviate from ideal behavior

____________________________________ ________________________________________
____________________________________ ________________________________________
____________________________________ ________________________________________
____________________________________ ________________________________________
____________________________________ ________________________________________
____________________________________ ________________________________________

24.State the conditions at which a gas approach the ideal behavior.
_________________________________________________________________________________
_________________________________________________________________________________

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PRACTICE MODULE CHAPTER 5 SK015

van der Waals Equation

25. Write the van der Waals equation for a real gas. For HCl and N2 which molecule has a higher value
of a and b? Justify your answer.
________________________________________________________________________________
________________________________________________________________________________
________________________________________________________________________________
________________________________________________________________________________
________________________________________________________________________________

WORKSHEET 2

Objective Questions

1. How does Charles’s law explain kinetic molecular theory?
A. As temperature increases, molecules move faster.
B. The pressure of a gas increases at constant temperature.
C. The pressure in a container increase as more molecules are added.
D. The volume of a gas is negligible compared to the volume of the container.

2. If the temperature of a gas at constant pressure is plotted on the x-axis against its volume on the y-

axis, a straight line is obtained which cuts the x-axis at

A. 0°C
B. −273.15°C
C. −275.13°C

D. 25 K

3. Which of the following statements is not true about gas?
A. The attractive forces between gas molecules for an ideal gas are negligible.
B. Real gases will behave ideally at high pressure and low temperature.
C. According to Charles’s law, at constant pressure, the volume of a certain mass of gas is directly
proportional to the temperature.
D. Molecules of ideal gas always collide with each other and the container walls but there is no loss
of kinetic energy.

4. The table below show the van der Waals constants for water and Sulphur dioxide.

Molecule a b
(atm dm3 mol−2) (dm3 mol−1)
H2O
SO2 5.54 0.0305

6.87 0.0568

A. H2O molecules are smaller and more attracted to each other than SO2 molecules.
B. H2O molecules are larger and more attracted to each other than SO2 molecules.
C. H2O molecules are smaller and less attracted to each other than SO2 molecules.
D. H2O molecules are larger and less attracted to each other than SO2 molecules.

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PRACTICE MODULE CHAPTER 5 SK015

5. According to Dalton’s law of partial pressures, in a mixture of gases, each gas
A. exerts only a fraction of its normal pressure.
B. exerts the same pressure as each of the other gases.
C. exerts a pressure equal to the total pressure of the mixture.
D. exerts a pressure equal to the pressure it would exert if it alone occupied the total volume

5.2: LIQUID

Properties Shape & Volume
Surface Tension

Viscosity
Compressibility

Diffusion

Processes Vaporization
Condensation

LIQUID Vapor Pressure: The pressure exerted by vapor in equilibrium with its liquid in
a closed container.

Boiling Point: The temperature at which the vapor pressure of a liquid is equal
to atmospheric pressure.

Boiling process Intermolecular force , vapor pressure
Relationship Vapor pressure boiling point

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PRACTICE MODULE CHAPTER 5 SK015

WORKSHEET 1

Properties of Liquid

1. Give any three properties of liquid.
__________________________________________________________________________________
__________________________________________________________________________________
__________________________________________________________________________________
__________________________________________________________________________________

2. State the properties of the liquid based on the meaning given.
(a) A measure of the resistance of fluid to flow

(b) A measure of the ability of the substance to be
forced into a smaller volume.

(c) The amount of energy needed to stretch or
expand the surface area of a liquid.

3. Give the suitable explanation for all the questions below. Answer

Question
(a) Why are water droplets

round?

(b) Why does the viscosity
of a liquid decrease with
heating?

(c) Why liquids are slightly
compressible?

3. Explain the kinetic concept of a liquid.
__________________________________________________________________________________
__________________________________________________________________________________
__________________________________________________________________________________
__________________________________________________________________________________
__________________________________________________________________________________

Vaporization and Condensation

4. (a) Define the term of vaporization and condensation.
(a) Explain vaporization and condensation process based on kinetic energy and intermolecular forces.

Vaporization Definition
___________________________________________________________________

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PRACTICE MODULE CHAPTER 5 SK015

Explanation
_________________________________________________________________

________________________________________________________________
________________________________________________________________
________________________________________________________________
________________________________________________________________
________________________________________________________________
________________________________________________________________

Condensatio Definition
n ________________________________________________________________
________________________________________________________________

Explanation
________________________________________________________________
________________________________________________________________
________________________________________________________________
________________________________________________________________
________________________________________________________________
________________________________________________________________
________________________________________________________________

Factors That Influence the Rate of Vaporization
6. Explain how three factors below influence the rate of vaporization.

Factors

Surface area Temperature Intermolecular attractive
forces
_______________________ _________________________
_______________________ _________________________ _______________________
_______________________ _________________________ _______________________
_______________________ _________________________ _______________________
_______________________ _________________________ _______________________
_______________________ _________________________ _______________________
_______________________ _________________________ _______________________
_______________________ _________________________ _______________________
_______________________

Vapor Pressure

7. Define vapor pressure of a liquid.
__________________________________________________________________________________
__________________________________________________________________________________
__________________________________________________________________________________

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PRACTICE MODULE CHAPTER 5 SK015

8. Complete the answer scheme below.
Dynamic equilibrium

Description
escape

collide continually with the wall of container and condense into…

Boiling
9. Complete the answer scheme below.

occurs at…

BBooiliilinninggg
is defined as…

Boiling point is called as…

The temperature at which the atmospheric pressure
vapor pressure of the liquid is
equal to... is called as…

1 atm
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PRACTICE MODULE CHAPTER 5 SK015

10.What do you understand by the term boiling?
_______________________________________________________________________________
_______________________________________________________________________________
_______________________________________________________________________________
_______________________________________________________________________________
_______________________________________________________________________________

11.Give two differences between vaporization and boiling.
_______________________________________________________________________________
_______________________________________________________________________________
_______________________________________________________________________________
_______________________________________________________________________________

12.Explain in your own words, why does water boil at a lower temperature on the mountain than at sea
level.
_______________________________________________________________________________
_______________________________________________________________________________
_______________________________________________________________________________
_______________________________________________________________________________

Relationship Between Intermolecular Forces, Vapor Pressure and Boiling Point.
13.Complete the answer scheme below.

Liquid with… strength of
has… intermolecular

forces

has…

vapour pressure

Thus, has energy Thus, has

boiling point
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PRACTICE MODULE CHAPTER 5 SK015

14.Arrange the following substances in order of increasing strength of intermolecular attraction based on
their boiling points.
Boiling point of; butanol (117°C), ethane (−88.6°C), water (100°C)

WORKSHEET 2

Objective Questions

1. Which of the following is true when a liquid change to a gas at the boiling point?
A. The temperatures stay constant.
B. The energy change is exothermic.
C. The kinetic energy of the liquid particles decreases.
D. The arrangement of particles becomes more orderly.

2. What happens to the surface tension of water when methanol is added?
A. The surface tension increases because methanol dissolves in water.
B. The surface tension increases due to the increased liquid pressure.
C. The surface tension does not change as methanol can form hydrogen bond with water.
D. The surface tension decreases because methanol disrupts the hydrogen bonds between the
water molecules.

3. The boiling point of water is defined as
A. the state where liquid changes to vapour.
B. the pressure at which the temperature reaches 100°C.
C. the temperature at which its vapour pressure is equal to the external pressure.
D. the temperature and pressure when solid, liquid and vapor can exist in equilibrium.

4. Water evaporates faster at higher temperature because
A. the molecule moves farther apart.
B. the surface area of water increases.
C. the surface area of water increases.
D. water molecules have larger kinetic energies.

5. Which of the following is true about liquid?
A. The vapour pressure of a pure liquid in a closed container increases two-fold if the volume of the
container is halved.
B. At constant temperature, each molecule in the liquid moves with the same velocity.
C. The vapour pressure of a liquid depends on its molecular size.
D. The boiling point of a liquid is dependent on the volume of the liquid.

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5.3: SOLID Fixed shape of a solid

SOLID Freezing (solidification)

Process Melting (fusion)
Sublimation
Deposition

Amorphous and crystalline solids

Metallic

Types of crystalline Ionic
solids Molecular covalent

Giant covalent

WORKSHEET 1 Explanation

1. Complete the table below:

Properties
Shape
Volume
Compressibility
Ability to flow
Density
Motion between particle

2. Explain the fixed – shape of a solid.
_______________________________________________________________________________
_______________________________________________________________________________

3. Which drawing represents a solid?

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PRACTICE MODULE CHAPTER 5 SK015

__________________ Add New state Process

4. Complete the table.
State

5. (a) What is the reverse process to melting?
____________________________________________________________________________

(b) What is the reverse process of evaporation?
_____________________________________________________________________________

(c) How do you think the particles in a substance behave when we give them more energy?
_____________________________________________________________________________

(d) Explain the steps that a solid must go through to become a gas.
_____________________________________________________________________________

6. Complete the table below: Explanation
Change of state

When heat is added and a solid change to a liquid

When heat is removed and a gas changes state to a liquid

When heat is added and the particles at the surface of a liquid change to
the gas state

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When heat is removed and a liquid change to a solid

7. Solid can be divide into two which is amorphous solid and crystalline solid. Give the differences
between both solid.

Crystalline solid Amorphous solid

Arrangement
of particles

Formation

Physical
properties

Example

8. State the types of crystalline solid in the following molecules/ atoms.

Molecule/ atom Types of crystalline solid

NaCl

Mg

SiO2

S8

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PRACTICE MODULE CHAPTER 5 SK015

WORKSHEET 2

Objective Questions

1. Which one of the following is non-crystalline or amorphous?
A. Diamond
B. Graphite
C. Glass
D. Common Salt

2. Which type of solid crystals will conduct heat and electricity?
A. Metallic
B. Ionic
C. Molecular
D. Covalent

3. Solids have...
A. a definite shape and volume
B. a definite shape but no definite volume
C. no definite shape but a definite volume
D. no definite shape or volume

4. Solids melt when solid particles _____ energy.
A. release
B. absorb
C. maintain
D. share

5. What is the change from a solid to a gas due to the gaining of energy?
A. Condensation
B. Vaporization
C. Sublimation
D. Deposition

6. When a solid changes into a liquid, the process is_______
A. melting
B. freezing
C. vaporization
D. condensation

7. Solid state is denser than the liquid and gaseous states of the same substance. Which of the following
is an exception to this rule?
A. Mercury
B. Carbon dioxide (dry ice)
C. Ice
D. NaCl

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8. Which of the following is an amorphous solid?
A. Quartz
B. Quartz glass
C. Graphite
D. Salt (NaCl)

9. Which of the following describes a general solid?
A. Compressible
B. Incompressible
C. Fluid
D. Semi-compressible

10.Which of the following is true of solids?
A. Solids maintain a defined shape and size under all conditions.
B. All solids have a crystalline structure.
C. All solids maintain a defined shape and size if conditions remain constant.
D. All solids have a lattice structure at the atomic level.

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5.4 PHASE DIAGRAM

• Phase: A homogeneous part of a system in
contact with other parts of the system but
separated from them by well-defined
boundary.

• Triple point: The temperature at which solid,
liquid and gas phases coexist in equilibrium.

• Critical point: The temperature and
pressure beyond which the liquid and gas
phases are indistinguishable from each other

PHASE
DIAGRAM

Phase diagram of CO2 Phase diagram of H2O
(+ve slope) (-ve slope)

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PRACTICE MODULE CHAPTER 5 SK015

WORKSHEET 1

1. Describe the following systems according to the component number, the phase number, and the
phase type.

Statement Component Phase Phase type

Ice and powdered glass number number
Water and oil
Water and alcohol
Water (10 g) and salt (1 g)
Water (10 g) and salt (20 g)

2.

(a) What phase is present in region A? Region B? Region C?
A: ___________________________________________
B: ___________________________________________
C: ___________________________________________

(b) What phases are in equilibrium at point 1? Point 2? Point 3? Point 5?
Point 1: _______________________________________
Point 2: _______________________________________
Point 3: _______________________________________
Point 5: _______________________________________

(c) Determine the critical point and triple point.
Critical point: ___________________________________
Triple point: ____________________________________

3. Referring the phase diagram of CO2. What phase(s) are present under these conditions:
(a) T =−70oC and P = 1.0 atm: ______________________________________________
(b) T =−40oC and P = 15.5 atm: _____________________________________________
(c) T =−80oC and P = 4.7 atm: ______________________________________________

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4. What phase of CO2 exists at 1.25 atm pressure and a temperature of (refer phase diagram of CO2)?
(a) −90oC: ___________________________________________
(b) −60oC: ___________________________________________
(c) 0oC: ___________________________________________

5. What phases of CO2 are present (refer phase diagram of CO2)?
(a) at a temperature of −78oC and a pressure of 1.0 atm: ____________________________________
(b) at −57oC and a pressure of 5.2 atm:
__________________________________________________

6. List phase that would be observed if a sample of CO2 at 10 atm pressure were heated from − 80oC
to 40oC. (Refer phase diagram of CO2)
_______________________________________________________________________________
__

7.

Use the phase diagram above to answer these questions.
(a) What is the temperature and pressure at the triple point?

____________________________________________
(b) Starting at the triple point, what phase exists when the pressure is held constant, and the

temperature is increased to 65oC?
____________________________________________
(c) Starting T = −70oC and P = 4 atm, what phase change occurs when the pressure is held
constant, and the temperature increased to −30oC?
____________________________________________
(d) Starting T = −30oC and P = 5 atm, what phase change occurs when the temperature is held
constant and the pressure increased to 10 atm?
____________________________________________

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PRACTICE MODULE CHAPTER 5 SK015

8. On the phase diagram for CO2, the point that corresponds to a temperature of – 13oC and a pressure
of 7.5 atm is on the vapour pressure curve where liquid phase and gas phase are in equilibrium. What
phase of CO2 exists when the pressure remains the same and the temperature increased by several
degrees?
________________________________________________

9.

Predict what would happen because of the following changes: (state the phase changes)
(a) Starting at A, we raise the temperature at constant pressure.

__________________________________________________
(b) Stating at C, we lower, the temperature at constant pressure.

__________________________________________________
(c) Starting at B, we lower the pressure at constant temperature.

__________________________________________________

10. Describe the change in phase when H2O at 100oC and 0.5 atm is slowly cooled to – 5oC, followed by
pressure increase to 1 atm. (refer phase diagram of H2O)
_______________________________________________________________________________
_______________________________________________________________________________
_______________________________________________________________________________
_______________________________________________________________________________

11.

Using the phase diagram above, determine the state of water at the following temperatures and
pressures:
(a) −10°C and 50 kPa: ________________________________

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PRACTICE MODULE CHAPTER 5 SK015

(b) 25°C and 90 kPa: _________________________________
(c) 50°C and 40 kPa: _________________________________
(d) 80°C and 5 kPa: __________________________________
(e) −10°C and 0.3 kPa: ________________________________
(f) 50°C and 0.3 kPa: _________________________________
12.Sketch and labeled the phase diagram for CO2.

WORKSHEET 2

Objective Questions

1. What is happening at the triple point on a phase diagram?
A. Solid, liquid and gas exist at equilibrium
B. Only gas and liquid exist at equilibrium
C. Only solid and liquid exist at equilibrium
D. Only solid and gas exist at equilibrium

2. Which of the following cannot be obtained using a phase diagram?
A. Melting temperatures of various phases
B. Temperature range for solidification
C. Equilibrium solid solubility
D. Purity of materials

3. On a phase diagram an isotherm indicates which of the following?
A. A region where the composition of the system is constant
B. A region where the pressure is constant
C. An area below which only the solid phase exists
D. A region where the temperature is constant

4. What is the point at which all phases can exist in equilibrium?
A. Critical point
B. Equilibrium point
C. Triple point
D. Super Critical Point

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PRACTICE MODULE CHAPTER 5 SK015

5. What are the three areas on the phase diagram for?
A. Rock, paper, and scissor
B. Solid, liquid and gas
C. Metal, nonmetal, and metalloid
D. Solid, water and gas.

6. What two phases does melting concern?
A. Plasma and fluid
B. Solid and gas
C. Liquid and gas
D. Solid and liquid

7. What is a phase diagram?
A. A graph of the physical state of a substance (solid, liquid, or gas) and the temperature and
pressure of the substance.
B. A diagram showing the phases of a gas.
C. A graph of the physical state of a substance (solid, liquid, or gas) and the temperature and state
of the substance.
D. A graph of the physical state of a substance (solid, liquid, or gas) and the reactants and products
of the substance.

8. What is NOT something that a phase equilibrium line show?
A. The equilibrium point between solid and liquid.
B. The equilibrium point between liquid and gas.
C. The equilibrium point between solid and gas.
D. The equilibrium point between gas and vapor.

9. Which best describes the significance of a "critical point" on a phase diagram?
A. The lowest temperature and pressure at which a substance may exist in a liquid phase.
B. The highest temperature and pressure at which a substance may exist in distinct solid and liquid
phases.
C. The highest temperature and pressure at which a substance may exist in distinct liquid and gas
phases
D. The temperature and pressure at which a substance may exist in equilibrium between the solid,
liquid, and gas phases.

10. Phase diagrams are used to depict changes in the properties of a solution at different temperatures
and pressures. Below is a phase diagram of a polar solution. What is the line A representing?

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PRACTICE MODULE CHAPTER 5 SK015

A. Deposition/sublimation
B. Freezing/melting
C. Vaporization/condensation
D. None of these

MODEL OF PSPM QUESTIONS

1. (a) State Boyle’s law. Explain Boyle’s law in terms of kinetic molecular teory.
(b) A cylinder with a movable piston containing 7.4 g of nitrogen dioxide, NO2 gas has a pressure of
115 kPa at 50°C. Calculate the gas pressure (in atm) if the volume of the cylinder is reduced to
1.50 L under the same temperature.
[PSPM 2007]

2. (a) For exactly 1 mole of an ideal gas, write the ideal gas equation. Give the two assumptions used in

ideal gas equation.

(b)The van der Waals equation is used to explain the non-ideal behavior of gases,

 P + a  (V − b)= RT
 V2 

Where a and b are the van der Waals constants. Explain how the assumptions gave in (a) are

corrected in the van der Waals equation.

[PSPM 2011]

3. The boiling point of several liquids is given as below:

Liquids Boiling point (°C)

Methanol 65

Ethanol 78
Propanol 97
Butanol 117

Arrange these liquids in order of decreasing vapour pressure. Explain your answer.
[PSPM 2001]

4. (a) Limestone, CaCO3, decomposed to solid calcium oxide and carbon dioxide gas when heated at
high temperature. At 30oC, a volume of 107.3 mL of the gas was collected by displacement of water
with a total pressure of 1 atm. Calculate:
i. the number of moles of carbon dioxide produced
ii. the mass of limestone decomposed
[Vapour pressure of water at 30oC is 31.8 mmHg]
(b) Briefly describe three types of crystalline solids in terms of their interparticle forces.
[PSPM 2014]

Prepared by : Mdm Siti Sarah & Mdm Farhana

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PRACTICE MODULE CHAPTER 6 SK015

CHAPTER 6 : CHEMICAL EQUILIBRIUM
6.1 DYNAMIC EQUILIBRIUM
Mind map / i-Think Map

REVERSIBLE Reactions that occur in
REACTION both directions (forward
and backward directions)

Equilibrium that exists in a
closed system when the rate

forward=rate reverse

6.1 DYNAMIC DYNAMIC rate forward=rate
EQUILIBRIUM EQUILIBRIUM reverse

LAW OF MASS Characteristic of [ ] reactant &
ACTION system in equilibirium product are
constants over the
Also known as
equilibrium law time

The reaction
quotient (Q) = The

equilibrium
constant (K)

definition: the ratio of the concentration of
the products, each raised to the power of its
coefficient, to the concentration of reactants,
each raised to the power of its coefficient, is a

constant

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PRACTICE MODULE CHAPTER 6 SK015

WORKSHEET 6.1
1. Explain the following term;

i. Reversible reaction:_______________________________________________________________

_______________________________________________________________________________

ii. Dynamic equilibrium:______________________________________________________________

_______________________________________________________________________________

_______________________________________________________________________________

iii. Law of mass action: ___________________________________________________________

____________________________________________________________________________

2. State 3 characteristics of a system in dynamic equilibrium.

i. ___________________________________________________________________________

ii. ___________________________________________________________________________

iii. ___________________________________________________________________________

3. Complete the graph of dynamic equilibrium for the following reaction.

N2O4 (g) 2NO2 (g)

Rate Concentration

Time Time
4. Write the law of mass action for the following reactions:

a. 2N2O5(g) ⇌ 4NO2(g) + O2(g)

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PRACTICE MODULE CHAPTER 6 SK015
b. Cd2+(aq) + 4CN-(aq) ⇌ Cd(CN)42-(aq)

c. H2(g) + I2(g) ⇌ 2HI(g)

d. 3H2(g) + N2(g) ⇌ 2NH3(g)

5. The equilibrium concentrations from two experiments of reactions (c) and (d) above are tabulated below.
Calculate the K for each experiment.

Experiment for [ H2(g) ]eq [ I2 (g) ]eq [HI(g) ]eq K
reaction (c)
0.0018 0.0031 0.0177
1 0.0011 0.0011 0.0084
2

Experiment for [H2(g)]eq [ N2 (g) ]eq [NH3(g)]eq K
reaction (d)
1.00 0.0866
1 0.500 1.15 0.412
2 1.35

6. At 745 K, K is 0.118 for the following reaction:

N2(g)+3H2(g)⇌2NH3(g)

What is the equilibrium constant for each related reaction at 745K?

a. 2NH3(g) ⇌ N2(g)+3H2(g)

b. 1/2N2(g)+3/2H2(g)⇌NH3(g)

c. 2N2(g)+6H2(g)⇌4NH3(g)

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PRACTICE MODULE CHAPTER 6 SK015

6.2 EQUILIBRIUM CONSTANT
MINDMAP / i-THINK MAP

is a measures the relative
amounts of products and
reactants present during a
reaction at a particular point

in time.

Reaction Q>K, Q=K, no
Quotient reaction net

(Q) will change
progress
Fraction of from right
molecule
dissociate to left
into smaller
molecules/ predict Q<K,
ions/ atoms direction of reaction

reaction will
progress
Degree of from left
Dissociation to right

(α)

6.2 DEFINE HOMOGENEOUS -
EQUILIBRIUM Reactants and
solve
equilibrium CONSTANT products are in the

problem relation same phase
between
if given if given Kp & Kc HETEROGENEOUS -
equilibrium initial
quantities, quantities, Reactants and
solve Kp & use ICE products are in
table
Kc different phases

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PRACTICE MODULE CHAPTER 6 SK015
WORKSHEET 6.2(a)
1. Define the following term,

i. Homogeneous equilibria: _________________________________________________________
ii. Heterogeneous equilibria: _________________________________________________________
2. Write the equilibrium constant expression Kc and Kp (if any) for each of the following reactions.
i. 2H2O2(g) ⇌ 2H2O(g) + O2(g)

ii. 6H2O2(g) ⇌ 6 H2O(g) + 3O2(g)

iii. The reverse of the reaction (i)

iv. 2PbS(s) + 3O2(g) ⇌ 2PbO(s) + 2SO2(g)

v. MgCl2(s) ⇌ Mg2+(aq) + 2Cl-(aq)

vi. The reverse of the reaction (v)

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PRACTICE MODULE CHAPTER 6 SK015

WORKSHEET 6.2(b)
1. What is the equilibrium constant, Kc expression for the following reaction;

3Fe(s) + 4H2O(g ) ⇌ Fe3O4(s) + 4H2(g)

2. The equilibrium constant, Kc, for the reaction of 2NOCl(g) ⇌ 2NO(g) + Cl2(g) is 2.4 x 10-7. What is the
equilibrium constant, Kc, for this reaction 1/3Cl2 (g) + 2/3NO (g) ⇌ 2/3NOCl(g)?

Ans: 1.61 x 102

3. Calculate Kp for the reaction at 527°C, if Kc = 7.9×104 at this temperature.
2SO2(g) + O2(g) ⇌ 2SO3(g)

Ans: 1.2 x 103

4. A mixture of 9.22 moles of A, 10.11 moles of B, and 27.83 moles of C is placed in a one-liter container at a
certain temperature. The reaction is allowed to reach equilibrium. At equilibrium the number of moles of B is
18.32. Calculate the equilibrium constant for the reaction:
A(g) + 2B(g) ⇌ 3C(g)

Ans: 0.835
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PRACTICE MODULE CHAPTER 6 SK015
5. a. At a certain temperature, Kc is 4.13 x 10 -2 for the equilibrium:

2IBr(g) ⇌ I2(g) + Br2(g)
Assume that equilibrium is established at the above temperature by adding only IBr(g) to the reaction flask.
What are the concentrations of I2(g) and Br2(g) in equilibrium with 0.0124 moles/liter of IBr(g)?

Ans: [ 2] = [ 2] = 2.52 x 10-3 M
b. What was the initial concentration of IBr before equilibrium was established?

Ans: 0.0174 M
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PRACTICE MODULE CHAPTER 6 SK015
6. 0.924 mole of A(g) is placed in a 1.00 liter container at 700°C, where it is 38.8 % dissociated when equilibrium

was established.
3A(g) ⇌ 5B(g) + 2C(g)

What is the value of the equilibrium constant, Kc, at the same temperature?

Ans: 0.0241
7. The equilibrium constant for the reaction is 2.60 x 10-3 at 1100°C. If 0.820 mole of NO(g) and 0.223 mole

each of N2(g) and O2(g) are mixed in a 1.00 liter container at 1100°C, what are the concentrations of NO(g),
N2(g), and O2(g) at equilibrium?

2NO(g) ⇌ N2(g) + O2(g)

Ans: [NO]= 1.148 M, [N2] = 0.059M, [Cl2]= 0.059 M
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PRACTICE MODULE CHAPTER 6 SK015

8. Pure phosphorous pentachloride is injected into an otherwise empty reaction vessel so that it has an initial

pressure of 1.64 atm at 500K. If the system remains at 500K, what are the equilibrium partial pressures of

the three gases?

PCl5(g) ⇌ PCl3(g) + Cl2(g) Kp = 0.49 at 500K

: 5 = 0.9557 , 3 = 0.6843 , 2 = 0.6843

WORKSHEET 6.2(c)

1. a. What is the numerical value of the equilibrium constant, Kc, for the reaction:
3N2(g) + 3O2(g) ⇌ 6NO(g)

if the equilibrium constant for the reaction; 2NO (g) ⇌ N2(g) + O2(g) is 3.5 x 10-6.

b. What is the equilibrium constant expression, Kc for the reaction: Ans: 2.33 x 1016
2Ni(s)+ 2CO2(g) ⇌ 2CO(g) + 2NiO(s) Ans: K= [CO]2/[CO2]2

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PRACTICE MODULE CHAPTER 6 SK015
2. A mixture of 1.16 mole of A, 1.35 mole of B and 0.641 mole of C is placed in a one-liter container at a certain

temperature. The reaction was allowed to reach equilibrium. At equilibrium. the number of moles of A is 1.95.
Calculate the equilibrium constant, Kc, for the reaction:

2A(g) ⇌ 2B(g) + C(g)

Ans: 0.0203
3. At 500 K the reaction PCl5(g) ⇌ PCl3(g) + Cl2(g) has Kp = 0.497. In an equilibrium mixture at 500 K, the partial

pressure of PCl5 is 0.860 atm and that of PCl3 is 0.350 atm. What is the partial pressure of Cl2 in the
equilibrium mixture?

Ans: 1.22 atm
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PRACTICE MODULE CHAPTER 6 SK015

4. Sulfur trioxide decomposes at high temperature in a sealed container:
2SO3(g) ⇌ 2SO2(g) + O2(g)

Initially, the vessel is charged at 1000 K with SO3(g) at a partial pressure of 0.500 atm. At equilibrium the
SO3(g) partial pressure is 0.200 atm. Calculate the value of Kp at 1000K.

Ans: 0.338

5. In the synthesis of ammonia from nitrogen and hydrogen,
N2(g) + 3H2 (g) ⇌ 2NH3(g)

Kc = 9.60 at 300 °C. Calculate Kp for this reaction at this temperature.

Ans: 4.34 x10-3
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PRACTICE MODULE CHAPTER 6 SK015
6. 0.822 mole of SO3(g) is placed in a 1.00 liter container at 600K. 36.7% of the SO3(g) are decomposed when

equilibrium is established.
2SO3(g) ⇌ 2SO2(g) + O2(g)

What is the value of the equilibrium constant, Kc and Kp, at the same temperature?

Ans: Kc= 5.08 x 10-2 , Kp = 2.501
7. 0.0487 mol of NOCl(g) was injected into an otherwise empty 1.00-L reaction vessel at 500 K. It decomposed

into nitrogen monoxide and chlorine:
2NOCl(g) ⇌ 2NO(g) + Cl2(g)

When the system reached equilibrium at 500K, the total pressure within the reaction vessel was found to be
2.63 atm. What is Kp for the above reaction at 500K?

Ans: 1.87
142