CHEMISTRY 13th EDITION LABORATORY MANUAL

MINISTRY OF EDUCATION
MATRICULATION DIVISION

CHEMISTRY

LABORATORY MANUAL

SK015 &
SK025

13th EDITION

MATRICULATION DIVISION
MINISTRY OF EDUCATION MALAYSIA

CHEMISTRY

LABORATORY MANUAL
SEMESTER I & II
SK015 & SK025

MINISTRY OF EDUCATION MALAYSIA
MATRICULATION PROGRAMME

THIRTEENTH EDITION

First Printing, 2003
Second Printing, 2004
Third Printing, 2005 (Sixth Edition)
Fourth Printing, 2006 (Seventh Edition)
Fifth Printing, 2007 (Eighth Edition)
Sixth Printing, 2011 (Ninth Edition)
Seventh Printing, 2013 (Tenth Edition)
Eighth Printing, 2018 (Eleventh Edition)
Ninth Printing, 2020 (Twelfth Edition)
Tenth Printing, 2022 (Thirteenth Edition)
Copyright © 2022 Matriculation Division
Ministry of Education Malaysia

ALL RIGHTS RESERVED. No part of this publication may be reproduced or
transmitted in any form or by any means, electronic or mechanical, including
photocopying, recording or any information storage and retrieval system, without
the prior written permission from the Director of Matriculation Division, Ministry
of Education Malaysia.

Published in Malaysia by

Matriculation Division
Ministry of Education Malaysia,
Level 6 – 7, Block E15,
Government Complex Parcel E,
Federal Government Administrative Centre,
62604 Putrajaya,
MALAYSIA.
Tel: 603-8884 4083
Fax: 603-8884 4028
Website: www.moe.gov.my

Printed in Malaysia by

Malaysia National Library
Chemistry Laboratory Manual
Semester I & II
SK015 & SK025
Thirteenth Edition

e ISBN 978-983-2604-63-1

i

NATIONAL EDUCATION PHILOSOPHY

Education in Malaysia is an on-going effort towards
further developing the potential of individuals in a
holistic and integrated manner, so as to produce
individuals who are intellectually, spiritually and
physically balanced and harmonious based on a firm
belief in and devotion to God. Such an effort is designed
to produce Malaysian citizens who are knowledgeable
and competent, who possess high moral standards and
who are responsible and capable of achieving a high
level of personal well-being as well as being able to
contribute to the betterment of the family, society and
the nation at large.

NATIONAL SCIENCE EDUCATION
PHILOSOPHY

In consonance with the National Education Philosophy,
science education in Malaysia nurtures a science and
technology culture by focusing on the development of
individuals who are competitive, dynamic, robust and
resilient and able to master scientific knowledge and
technological competency.

ii

FOREWORD

I am delighted to write the foreword for the Laboratory Manual,
which aimed to equip students with knowledge, skills, and the
ability to be competitive undergraduates.

This Laboratory Manual is written in such a way to emphasise
students’ practical skills and their ability to read and understand
instructions, making assumptions, apply learnt skills and react
effectively in a safe environment. Science process skills such as
making accurate observations, taking measurement in correct
manner, using appropriate measuring apparatus, inferring,
hypothesizing, predicting, interpreting data, and controlling
variables are further developed during practical session. The
processes are incorporated to help students to enhance their
Higher Order Thinking Skills such as analytical, critical and
creative thinking skills. These 21st century skills are crucial to
prepare students to succeed in Industrial Revolution (I.R.) 4.0.

The manipulative skills such as handling the instruments, setting
up the apparatus correctly and drawing the diagrams can be
advanced through practical session. The laboratory experiments
are designed to encourage students to have enquiry mind. It
requires students to participate actively in the science process
skills before, during and after the experiment by preparing the pre-
report, making observations, analysing the results and in the
science process skills before, during, after the experiment by
preparing the pre-report, making observations, analysing the
results and drawing conclusions.

It is my hope and expectation that this manual will provide an
effective learning experience and referenced resource for all
students to equip themselves with the skills needed to fulfil the
prerequisite requirements in the first-year undergraduate studies.

DR HAJAH ROSNARIZAH BINTI ABDUL HALIM
Director
Matriculation Division

iii

CONTENTS

Students Learning Time (SLT) Page
Learning Outcomes v
Laboratory Safety v
Ethics In The Laboratory viii
Preparation For Experiment xi
xi
Experiment Semester I
1 Title 1
2 Determination of the formula unit of a compound 5
3 Acid Base Titration – Determination of the 10
4 concentration of hydrochloric acid solution 15
5 Determination of the molar mass of a metal 22
6 Charles’ Law and the ideal gas Law 29

Chemical Equilibrium

pH measurement and its applications

Experiment Semester II 36
1 Title 40
2 Rate of reaction 44
3 Determining the heat of reaction 49
4 Electrochemical cells 54
5 Reactions of aliphatic and aromatic hydrocarbons 58
6 Reactions of hydroxy compounds
Aldehydes and ketones

References 62
Acknowledgements 63

iv

Chemistry Lab Manual

1.0 Student Learning Time (SLT)

Students will be performing the experiment within the time allocated
for each practical work.

Face-to-face Non face-to-face
2 hour 0

2.0 Learning Outcomes

2.1 Matriculation Science Programme Educational Objectives

Upon a year of graduation from the programme, graduates
are:

1. Knowledgeable and technically competent in science
disciplines study in-line with higher educational
institution requirement.

2. Able to apply information and use data to solve
problems in science disciplines.

3. Able to communicate competently and collaborate
effectively in group work to compete in higher
education environment.

4. Able to use basic information technologies and engage
in life-long learning to continue the acquisition of new
knowledge and skills.

5. Able to demonstrate leadership skills and practice
good values and ethics in managing organisations.

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Chemistry Lab Manual

2.2 Matriculation Science Programme Learning Outcomes

At the end of the programme, students should be able to:

1. Acquire knowledge of science and mathematics as a
fundamental of higher level education.
(MQF LOC i – Knowledge and understanding)

2. Apply logical, analytical and critical thinking in
scientific studies and problem solving.
(MQF LOC ii – Cognitive skills)

3. Demonstrate manipulative skills in laboratory works.
(MQF LOC iii a – Practical skills)

4. Collaborate in group work with skills required for
higher education.
(MQF LOC iii b – Interpersonal skills)

5. Deliver ideas, information, problems and solution in
verbal and written communication.
(MQF LOC iii c – Communication skills)

6. Use basic digital technology to seek and analyse data
for management of information.
(MQF LOC iii d – Digital skills)

7. Interpret familiar and uncomplicated numerical data to
solve problems.
(MQF LOC iii e – Numeracy skills)

8. Demonstrate leadership, autonomy and responsibility
in managing organization.
(MQF LOC iii f – Leadership, autonomy and
responsibility)

9. Initiate self-improvement through independent
learning.
(MQF LOC iv – Personal and entrepreneurial skills)

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Chemistry Lab Manual

10. Practice good values attitude, ethics and accountability
in STEM and professionalism.
(MQF LOC v – Ethics and professionalism)

2.3 Course Learning Outcome

2.3.1 Chemistry 1

At the end of the course, student should be able to:

1. Explain the fundamental concepts in chemistry.
(C2, PLO 1, MQF LOC i)

2. Solve problems with analytical and critical thinking
using chemistry facts and principles.
(C4, PLO 2, MQF LOC ii)

3. Apply the appropriate scientific laboratory skills in
chemistry experiments.
(P3, PLO 3, MQF LOC iii a)

4. Analyse numerical problems in chemistry.
(C4, PLO 7, MQF LOC iii e)

2.3.2 Chemistry 2

At the end of the course, student should be able to:

1. Explain the fundamental concepts in chemistry.
(C2, PLO 1, MQF LOC i)

2. Solve problems with analytical and critical thinking
using chemistry facts and principles.
(C4, PLO 2, MQF LOC ii)

3. Apply the appropriate scientific laboratory skills in
chemistry experiments.
(P3, PLO 3, MQF LOC iii a)

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Chemistry Lab Manual

2.4 Objectives of Practical Sessions

The main purpose of the experiment is to give the student a
better insight of the concepts of Chemistry discussed in the
lectures by carrying out experiments. The aims of the
experiments are to enable students to:

1. Learn and practise the necessary safety precautions in
the laboratory.

2. Plan, understand and carry out the experiment.

3. Use the correct techniques in handling the apparatus.

4. Acquire scientific skills in measuring, recording and
analysing data.

5. Observe, measure and record data by giving
consideration to the consistency, accuracy and units
of the physical quantities.

6. Determine the errors in various physical quantities
obtained in the experiments.

7. Deduce logically and critically the conclusion based
on observation and data analysis.

3.0 Laboratory Safety

The Science Matriculation Programme requires the students to
attend practical classes two hours a week to complete six
experiments each semester.

In order for the laboratory to be a safe place to work in, students
should learn laboratory rules and regulations, including the correct
way of using laboratory apparatus and handling of chemicals before
starting any experiments.

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Chemistry Lab Manual

Laboratory rules and regulations.

1. Attendance is COMPULSORY. If you are unable to attend any
practical class, you should produce a medical certificate or a
letter of exemption.

2. Read, understand and plan your experiment before pre-lab
sessions and practical classes.

3. Wear shoes, lab coats and safety goggles at all times in the
laboratory.

4. Tie long hair or tuck head scarf under your lab coat

5. Do not wear contact lenses during experiments.

6. Foods and drinks are not allowed in the laboratory.

7. Do not perform any unauthorised experiments! Understand
and follow the specified procedures for each experiment.

8. Do not waste chemicals. Take only sufficient amount of
chemicals needed for your experiments.

9. Replace the lids or stoppers on the reagent bottles or
containers immediately after use.

10. Do not remove chemicals from the laboratory.

11. Handle volatile and hazardous compounds in the fume
cupboard. Avoid skin contact with all chemicals, wash off any
spillages.

12. Clean up spillages immediately. In case of a mercury spillage,
do not touch the mercury. Notify your instructor immediately.

13. Ensure there are no flames in the vicinity before working with
flammable chemicals

14. NEVER leave an ongoing experiment unattended.

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Chemistry Lab Manual

15. Be aware or familiar with the location and proper way of
handling safety equipment, including eyewash, safety shower,
fire blanket, fire alarm and fire extinguisher.

16. Turn off bunsen flames when not in use. Notify your
instructor immediately of any injury, fire or explosion

17. Do not throw any solid wastes into the sink. Dispose any
organic substances in the waste bottles provided.

18. Wash all glassware after use and return the apparatus to its
appropriate places.

19. Keep your work area clean and tidy.

20. Notify your instructor immediately of any injury, fire or
explosion

I have read and understood the laboratory rules and regulations
as stated above. I agree to abide by all these rules, follow the
instructions and act responsibly at all times.

Signature: Date:
Name: Practicum:
Matric number:
Signature Instructor: Date:

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Chemistry Lab Manual

4.0 Ethics in the laboratory

1. Follow the laboratory rules.

2. Students must be punctual for the practical session. Students
are not allowed to leave the laboratory before the practical
session ends without permission.

3. Co-operation between members of the group must be
encouraged so that each member can gain experience in
handling the apparatus and take part in the discussions about
the results of the experiments.

4. Record the data based on the observations and not based on
any assumptions. If the results obtained are different from
the theoretical value, state the possible reasons.

5. Get help from the instructor or the laboratory assistant
should any problems arise during the practical session.

5.0 Preparation for experiment

5.1 Pre-lab Sessions.

1. Read and understand the objectives and the theory of
the experiment.

2. Think and plan the working procedures properly for
the whole experiment. Make sure you have
appropriate table for the data.

3. Complete and submit the pre-lab questions provided.

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Chemistry Lab Manual

5.2 Practical Sessions
1. Check the apparatus provided.
2. Conduct the experiment carefully.
3. Record all measurements and observations made
during the experiment.
4. Keep the work area clean and tidy.

5.3 Post-lab Sessions

1. Explain what has been carried out and discuss the
findings of the experiment.

2. Introduce the format of report writing as below:

Objective ▪ state clearly
Theory ▪ write concisely in your own words
Procedure ▪ draw and label diagram if necessary
Results/ ▪ write in passive sentences about all the
Observation
steps taken during the experiment
Discussion ▪ data tabulation with units and

Conclusion uncertainties
▪ data processing (plotting graph,

calculation to obtain the results of the

experiments and its uncertainties)

▪ give comments about the experimental

results by comparing it with the standard

value.
▪ state the source of mistake(s) or error(s)

if any as well as any precaution(s) taken

to overcome them.
▪ answer all the questions given

▪ state briefly the results with reference to

the objectives of the experiment

Reminder: NO PLAGIARISM IS ALLOWED.

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CHEMISTRY 1
SK015

Chemistry Lab Manual – SK015

EXPERIMENT 1: DETERMINATION OF THE FORMULA UNIT
OF A COMPOUND

Objectives

At the end of this lesson, students should be able to:
i. synthesise a zinc chloride compound; and
ii. determine the formula unit of zinc chloride.

Introduction

One of the main properties of a compound is its chemical composition which
can be identified by determining the elements present. A quantitative
analysis can be used to determine the composition of an unknown
compound. Once the composition of the compound is known, it’s formula
unit can be determined. For example, a compound containing 0.1 mol of
silver and 0.1 mol of bromine will have a formula unit, AgBr.

In this experiment, a simple compound composed of zinc and chlorine will
be prepared. Once the mass of zinc and the mass of the compound are
known, the mass of chlorine can be determined. Using these masses, the
percentage composition of the compound can be calculated and the formula
unit can be deduced.

Apparatus Chemical Reagents

Hot plate 6 M HCl
Glass rod Zinc powder
White tile
Crucible tongs
50 mL crucible
Analytical balance
Measuring cylinder (10 mL)

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Chemistry Lab Manual – SK015

Procedure

1. Weigh the crucible and record the exact mass.

2. Place approximately 0.25 g of zinc powder into the crucible and
determine the exact mass of zinc powder.

3. Carefully add in 10 mL of 6 M HCl solution into the crucible
containing the zinc powder and stir gently with a glass rod. A vigorous
chemical reaction will occur and hydrogen gas will be released.

Carry out this step in a fume cupboard. Do not
CAUTION ! work near a fire source. Wet hydrogen gas can

cause explosions.

4. If the zinc powder does not dissolve completely, continue adding the
acid, 5 mL at a time until all zinc is dissolved. The amount of acid to be
used must not exceed 20 mL.

5. Place the crucible on a hot plate in the fume cupboard and heat the
content slowly so that the compound does not splatter during the
heating process.

6. Heat the compound gently until it is completely dry. Remove the
crucible from the hot plate immediately to avoid the compound from
melting.
(Precaution: Avoid from overheating)

7. Cover the crucible and allow it to cool to room temperature. Then
weigh the crucible and the compound. Record the mass.

8. Reheat the crucible to dry the compound. Let it cool to room
temperature and then weigh it again. Repeat the procedure until the
difference in mass does not exceed 0.02 g.

9. Determine the mass of zinc chloride from the final weight of the
sample (the smallest value). Calculate the mass of chlorine in the zinc
chloride.

10. Determine the formula unit of zinc chloride.

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Chemistry Lab Manual – SK015

POINT TO DISCUSS
1. Explain why the content is not weighed while it is still hot.
2. Explain why the crucible needs to be covered during cooling.
3. Write a balanced equation for the reaction between zinc and

hydrochloric acid.

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Chemistry Lab Manual – SK015

DATA SHEET

EXPERIMENT 1: DETERMINATION OF THE FORMULA UNIT
OF A COMPOUND

RESULTS Mass (g)

No. Item
1. Mass of empty crucible
2. Mass of crucible + zinc powder
3. Mass of zinc powder
4. Mass of crucible + zinc chloride:

i) First heating
ii) Second heating
iii) Third heating
5. Mass of zinc chloride
6. Mass of chlorine

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Chemistry Lab Manual – SK015

EXPERIMENT 2: ACID-BASE TITRATION − DETERMINATION
OF THE CONCENTRATION OF
HYDROCHLORIC ACID SOLUTION

Objectives

At the end of this lesson, students should be able to:
i. prepare a standard solution of oxalic acid.
ii. standardise 0.2 M NaOH solution.
iii. determine the concentration of HCl solution.
iv. acquire the correct techniques of titration

Introduction

Titration is a laboratory technique used to determine the concentration of a
solution using another solution with a known concentration.

Standards in acid-base titrations

One of the solutions involved in a titration is used as a standard solution.
The standard solution can be classified as either primary or secondary. A
primary standard solution is prepared by dissolving an accurately weighed
pure solid of a known molar mass in a known volume of distilled water.

A primary standard is used to determine the molarity of the other standard
solution, known as a secondary standard. For example, oxalic acid,
H2C2O4, and potassium hydrogen phthalate, KHC8H4O4, are two common
primary standards used to determine the concentration of bases (secondary
standard).

The NaOH solution used in titrations need to be standardized because they
contain impurities. Solid NaOH is hygroscopic (it absorbs moisture). Thus, it
is difficult to obtain its accurate mass. The standardized NaOH becomes the
secondary standard and can then be used to determine the concentration of
other acids such as HCl acid.

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Chemistry Lab Manual – SK015

Equivalence point and end point

An equivalence point is the point in a titration at which the added titrant
reacts completely with the electrolyte according to stoichiometry.To detect
this equivalence point, an indicator which produces a change in colour is
often used. The point at which the indicator changes colour is called the end
point. The end point and equivalence point should ideally be the same.

Chemical equations

In this acid-base titration, the neutralisation reactions involved are:

H2C2O4(aq) + 2NaOH(aq) → Na2C2O4(aq) + 2H2O(l) . . .(1)

HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l) . . .(2)

Apparatus Chemical Reagents

Burette x M HCl
Dropper 0.2 M NaOH
Glass rod Distilled water
White tile Phenolphthalein
Retort stand Hydrated oxalic acid, H2C2O4.2H2O
Filter funnel
50 mL beaker
25 mL pipette
Analytical balance
250 mL conical flask
250 mL volumetric flask
50 mL measuring cylinder

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Chemistry Lab Manual – SK015

Procedure

(A) Preparation of standard solution

1. Weigh to the nearest 0.0001 g about 3.00 g of hydrated oxalic
acid, H2C2O4.2H2O in a 50 mL beaker.

2. Add approximately 30 mL of distilled water to dissolve the
oxalic acid.

3. Transfer the solution into a 250 mL volumetric flask. Rinse the
beaker and pour the content into the flask. Add distilled water
up to the calibrated mark of the volumetric flask.

4. Stopper and shake the flask to obtain a homogeneous solution.

5. Calculate the concentration of the standard oxalic acid solution.

NOTE: Use this solution to standardize the NaOH solution in
Part (B).

(B) Standardisation of 0.2 M NaOH solution

1. Rinse a burette with a given NaOH solution to be standardized.

2. Fill the burette with the NaOH solution. Ensure there are no air
bubbles trapped at the tip.

3. Record the initial burette reading to two decimal places.

4. Pipette 25 mL of oxalic acid solution from Part (A) into a 250
mL conical flask. Add 2 drops of phenolphthalein to the oxalic
acid solution.

5. Place a white tile underneath the flask so that any colour
change can be clearly observed.

6. Titrate the acid with the NaOH solution from the burette.
During the titration, swirl the flask continuously.

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Chemistry Lab Manual – SK015

7. Rinse the unreacted solutions at the inner wall of the conical
flask with distilled water.

8. Upon reaching the end point, a temporary pink solution appears
but fades when the solution is swirled. Continue titrating until a
pale pink colour persists for more than 30 seconds. This is the
end point.

9. Record the final burette reading to two decimal places.
10. Repeat the titration three times.
11. Calculate the molarity of the NaOH solution.

(C) Determination of the molar concentration of HCl solution.
1. Pipette 25 mL of a given HCl solution into a 250 mL conical
flask.
2. Add two drops of phenolphthalein.
3. Repeat steps 5-9 as in Part (B).
4. Calculate the concentration of HCl.

POINT TO DISCUSS

Does the addition of water in step 7 (Part B) affect the result of the
titration? Explain.

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Chemistry Lab Manual – SK015

DATA SHEET

EXPERIMENT 2: ACID-BASE TITRATION − DETERMINATION
OF THE CONCENTRATION OF
HYDROCHLORIC ACID SOLUTION

RESULTS

(A) Preparation of standard oxalic acid solution

i. Exact mass of hydrated oxalic acid =
ii. Moles of hydrated oxalic acid =
iii. Molarity of oxalic acid =

(B) Standardisation of 0.2 M NaOH solution

Burette reading / mL Gross I II III

Final reading

Initial reading
Volume of NaOH used /
mL

Average volume of NaOH used =

Calculate the molarity of the NaOH solution.

(C) Determination of the molar concentration of HCl solution

Burette reading / mL Gross I II III

Final reading

Initial reading
Volume of NaOH used /
mL

Average volume of NaOH used =

Calculate the molarity of the HCl solution

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Chemistry Lab Manual – SK015

EXPERIMENT 3: DETERMINATION OF THE MOLAR MASS OF
A METAL

Objectives

At the end of this lesson, students should be able to:
i. standardize the hydrochloric acid solution.
ii. determine the molar mass of an alkaline earth metal by back- titration

method.

Introduction

A reactive metal, for example an alkaline earth metal, would readily react
with a strong acid such as hydrochloric acid. The general reaction between a
metal, M and an aqueous hydrochloric acid, HCl is as follows:

M(s) + 2HCI(aq) → MCl2(aq) + H2(g)

The molar mass of M can be determined by a back-titration. A back
titration is a two-stage analytical technique. The first stage involves the
reaction of a metal with an excess amount of acid of a known concentration.
In the second stage, the unreacted acid is titrated with a standardized base
solution to determine the amount of the remaining excess reactant.

In this experiment, the concentration of the acid is initially determined by the
normal titration before the reaction with metal M is carried out. M reacts
completely according to stoichiometric equation and if the amount of acid
used exceeds the amount of metal in terms of equivalence, then the resulting
solution would be acidic.

The excess acid can be determined by performing back-titration with sodium
hydroxide solution. The amount in moles of the reacted metal is determined
by comparing the moles of acid before and after the reaction.

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Chemistry Lab Manual – SK015

Apparatus Chemical Reagents

Scissors Distilled water
Dropper Phenolphthalein
White tile x M Hydrochloric acid, HCl
Pipette filler 1.0 M Sodium hydroxide, NaOH,
Filter funnel An unknown alkaline earth metal, M
Retort stand
50 mL beaker
50 mL burette
25 mL pipette
Analytical balance
250 mL conical flask
Abrasive cloth no.3 (36) / sandpaper
Aluminium oxide

Procedure

(A) Standardization of HCl solution

1. Rinse a clean burette with 1.0 M NaOH.

2. Fill the burette with 1.0 M NaOH solution.

3. Record the initial burette reading to two decimal places.

4. Pipette 25 mL HCl solution into a 250 mL conical flask. Add 2
drops of phenolphthalein to the acid.

5. Place a piece of white tile underneath the flask.

6. Titrate the acid with the NaOH solution. Swirl the flask
continuously.

7. Upon reaching the end point, a temporary pink solution will
appear but the colour will fade when it is swirled. Continue
titrating until the pale pink colour remains for more than 30
seconds. This is the end point.

8. Record the final reading of the burette.

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Chemistry Lab Manual – SK015

9. Repeat the titration three times.
10. Calculate the concentration of the HCI solution.

(B) Determination of the molar mass of a metal
1. Pipette 25 mL of HCl solution into 2 separate conical flasks.
2. Clean two pieces of metal M, each of approximately 4 cm long,
with a piece of abrasive cloth.
3. Weigh accurately the mass of each sample.
4. Cut each sample into smaller pieces. Place the samples
separately into the HCl solution. Swirl occasionally until the
metal is completely dissolved.
5. Add 2 drops of phenolphthalein.
6. Record the initial burette reading.
7. Titrate the unreacted HCl with the NaOH solution.
8. Record the final burette reading.
9. Repeat titration with the other sample.

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Chemistry Lab Manual – SK015

DATA SHEET

EXPERIMENT 3: DETERMINATION OF THE MOLAR MASS OF
A METAL

RESULTS
1. Titration of standard HCl solution

Concentration of NaOH = ___________ M

Volume of HCl = ___________ mL

Burette reading / mL Gross I II III

Final reading

Initial reading
Volume of NaOH /
mL

Average volume of NaOH =

2. Reaction of metal and HCl

Mass of metal (sample I) (g)
Mass of metal (sample II) (g)

3. Titration of unreacted HCl

Burette reading / mL Sample I Sample II
Final reading
Initial reading
Volume of NaOH (mL)

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Chemistry Lab Manual – SK015

CALCULATION
1. Calculate the molarity of the standard HCl solution.
2. Determine the number of moles of HCl in 25 mL of the standard

solution.
3. Calculate the number of moles of the unreacted HCl solution.

Sample I:

Sample II:

4. Calculate the number of moles of the reacted metal.
Sample I:

Sample II:

5. Determine the molar mass of metal in each sample.
Sample I:

Sample II:

Average molar mass of metal = _______

6. By comparing the results with elements in the periodic table,
determine the metal M.

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Chemistry Lab Manual – SK015

EXPERIMENT 4: CHARLES’ LAW AND THE IDEAL GAS LAW

Objectives

At the end of this lesson, students should be able to:
i. verify Charles’ Law.
ii. determine the molar mass of a volatile liquid.

Introduction

Charles’ Law states that the volume of a fixed mass of a given gas is
directly proportional to its absolute temperature at constant pressure. The law
is written as

V  T (n, P constant)

In this experiment, a quantity of air is trapped between the sealed end of a
thick-walled glass tube (with a small cross-sectional area) and a movable
plug of mercury. If the glass tube is held upright, the plug of mercury will
move to a position where the pressure of the air in the tube is equal to the
atmospheric pressure and a small pressure exerted by the plug. Thus, the
pressure of the trapped air is constant.

The volume, V, of the trapped air is obtained by multiplying the cross-
sectional area of the tube, A, with the height of the air column, h.

V= A x h

Assuming that the cross-sectional area is constant, the volume is directly
proportional to the height, i.e., V  h. Therefore, the height of the air column
can be used as a measure of the volume in this experiment. By measuring
this height at different temperatures we can determine the relationship
between the volume of the trapped air and its temperature at constant
pressure.

Ideal Gas Equation:

By combining the relationships govern by the gas laws, a general equation
known as the ideal gas equation can be obtained.

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Chemistry Lab Manual – SK015

Boyle’s Law
Volume of a fixed mass of a given gas is inversely proportional to its
pressure at constant temperature.

V  1 (n, T constant)
p

Avogadro’s Principle
All gases of equal volume will contain the same number of molecules at the
constant temperature and pressure.

V  n (T, P constant)

Charles’ Law

Volume of a fixed mass of a given gas is directly proportional to its absolute

temperature at constant pressure.

V  T (n, P constant)

Thus, combining the three laws, we get

V  nT
p

The above expression can be written as

V = RnT or PV = nRT ...........(1)
P

This is the ideal gas equation and R is called the gas constant. The number of
moles, n,

mass
n=

Molar mass, Mr

Therefore, the ideal gas equation can also be written as

RT ..........(2)
PV = m ( )

Mr

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Apparatus Chemical Reagents

Needle Ice
Wire gauze Methanol
Tripod stand Unknown liquid
Rubber band
Thermometer
Bunsen burner
Aluminum foil
Analytical balance
Barometer
Retort stand and clamp
Charles’ law apparatus
600 mL beaker
100 mL conical flask
100 mL measuring cylinder
Stopwatch

Procedure

(A) Charles’ Law

1. Tie a thermometer to a glass tube containing a plug of
mercury with a rubber band. The bulb of the thermometer is
placed approximately half-way up the column of the trapped
air as shown in Figure 4.1.

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Charles’ law apparatus

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2. Fill a 100 mL measuring cylinder with tap water. Place the
tube and the thermometer into the water until the air column
in the tube is immersed.

3. Leave for 5 minutes to ensure that the temperature of the
trapped air is equivalent to the temperature of the tap water.

4. Record the temperature and measure the height of the air
column.

5. Repeat Steps 2 until 4 using:

i. warm water (40 – 50°C)
ii. a mixture of ice and water
iii. a mixture of ice and 5 mL methanol

NOTE: Ensure that the mercury plug does not split into small
droplets.

(B) Determination of the molar mass of a gas

1. Cover a 100 mL conical flask with a piece of aluminium foil
and tie it loosely around the neck with a rubber band as shown
in Figure 4.2.

Figure 4.2 Figure 4.3

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2. Prick a tiny hole in the middle of the foil with a needle.

3. Weigh the apparatus accurately.

4. Remove the foil and place 5.0 mL of the unknown liquid into
the flask.

5. Replace the foil and tie it with a rubber band.

6. Clamp the neck of the flask and immerse it into a 600 mL
beaker containing water as shown in Figure 4.3.

7. Heat the water until all of the unknown liquid in the flask has
vaporised.

8. Record the temperature of the water bath when all the
unknown liquid has evaporated.

9. Take out the flask immediately by using the clamp.

10. Wipe the outer wall of the flask and the aluminium foil when
the flask is cooled.

11. Weigh the flask with the aluminium foil, rubber band and the
condensed unknown liquid.

12. Discard both the foil and the condensed liquid. Fill the flask
up to the brim with water and pour it into a measuring
cylinder. Record the volume of water.

13. Record the barometric pressure.

14. Calculate the molar mass of the unknown liquid using the
ideal gas equation.

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DATA SHEET

EXPERIMENT 4: CHARLES’ LAW AND THE IDEAL GAS LAW

RESULT
(A) Charles’ law

TABLE 1

Condition Temperature Volume
(0C) (Height of gas column)
Warm water
Tap water (mm)
Ice-water
Ice-methanol

1. Complete TABLE 1.

2. Plot the height of the column, h, against temperature, T, in celsius on a
graph paper. Based on the graph, state the relationship between volume
and temperature.

3. Extrapolate the line until h = 0, to obtain the absolute zero temperature.

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(B) Determination of the molar mass of the gas
TABLE 2

No Item Reading

1. Mass of flask + rubber band + cover (g)

2. Mass of flask + rubber band + cover + condensed
liquid (g)

3. Mass of condensed liquid (g)

4. Temperature of water bath (oC)

5. Barometric pressure (mm Hg)

6. Volume of flask (mL)

4. Complete TABLE 2.
5. Calculate the molar mass of the unknown liquid.

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EXPERIMENT 5: CHEMICAL EQUILIBRIUM

Objectives

At the end of this lesson, students should be able to:
i. study the effect of concentration and temperature on chemical

equilibrium.
ii. determine the equilibrium constant, Kc, of a reaction.

Introduction

There are two types of chemical reactions, namely irreversible and
reversible. A reversible reaction will reach a dynamic equilibrium when the
rate of the forward reaction is equal to the rate of the reverse reaction. At this
stage, one cannot observe any changes in the system as the concentration of
reactants are constant. This does not mean that the reactions have stopped,
instead, the reactions are still occurring but at the same rate.

The factors that influence chemical equilibrium are:
i. concentration
ii. temperature
iii. pressure (for reactions that involve gases)

A change in one of the factors on a system that is already at equilibrium, will
cause the reaction to move to the direction that minimizes the effect of
change. The direction of the change can be determined by applying Le
Chatelier’s Principle.

Le Chatelier’s Principle states that if a system at equilibrium is disturbed by
a change in temperature, pressure or concentration of one or more
components, the system will shift its equilibrium position in such a way so as
to counteract the effect of the disturbance.

The effect of concentration

According to the Le Chatelier’s principle, the change in concentration of any
substance in a mixture at equilibrium will cause the equilibrium position to
shift in the forward direction or reverse direction to re-attain the equilibrium.

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Consider a general reaction as follows:

A+B C+D

If substance A or B is added to a mixture at equilibrium, the reaction will
shift forward to reduce the concentration of A or B until equilibrium is re-
established.

On the other hand, if substance C or D is added, the equilibrium will shift in
the direction that will reduce the concentration of C or D, i.e. from right to
left until equilibrium is re-established.

The effect of temperature

The effect of temperature on an equilibrium system depends on whether the
reaction is exothermic or endothermic. Consider the following system:

E+F G + Heat

If the forward reaction is exothermic, then the heat released is considered as
one of the products. Heating the system will cause the equilibrium to shift in
the reverse direction so as to reduce the excess heat. Thus, the concentrations
of E and F increase while the concentration of G decreases. However, when
the system is cooled, the equilibrium will move forward to increase the heat
in the system. The same principle can be applied to explain an endothermic
system.

In this experiment, you will study the effect of changes in concentration and
temperature on two equilibrium systems. You can notice the shift in
equilibrium through changes in colour or phases such as precipitation or
dissolution.

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Apparatus Chemical Reagents

Burette 6 M HCl
Ice bath 0.2 M CoCl2
Test tube 2.5 M NaOH
Water bath 0.1 M KSCN
10 mL pipette 0.1 M Fe(NO3)3
100 mL beaker 0.5 M SbCl3 in 6 M HCl
100 mL conical flask
10 mL measuring cylinder
100 mL measuring cylinder

Procedure

(A) The effect of concentration in the formation of thiocyanoiron(III)
complex ion

The thiocyanoiron(III) complex ion is formed when iron(III) ion, Fe3+,
is added to the thiocyanate ion, SCN-. The equation for the reaction is

Fe3+ (aq) + 2SCN- (aq) [Fe(SCN)2]+(aq)
(Yellowish brown) (blood-red)

1. Place 2 mL of 0.1 M Fe(NO3)3 solution and 3 mL of 0.1 M
KSCN solution in a 100 mL beaker.

2. Add 50 mL of distilled water to reduce the intensity of the blood
red solution.

3. Place approximately 5 mL each of this solution into four test
tubes.

(a) To the first test tube, add 1 mL of 0.1 M Fe(NO3)3.
(b) To the second test tube, add 1 mL of 0.1 M KSCN.
(c) To the third test tube, add 6-8 drops of 2.5 M NaOH.
(d) The fourth test tube serves as a control.

4. Tabulate the observations.

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(B) The Effect of Temperature

The reaction between hexaaquocobalt(II) complex ion with chloride
ion produces tetrachlorocobalt(II) ion. The equation for the reaction is
given below:

[Co(H2O)6]2+(aq) + 4Cl-(aq) [CoCl4]2-(aq) + 6H2O(l)
(pink) (blue)

1. Place 2 mL of 0.2 M CoCl2 solution into a conical flask.

2. Add 20 mL of 6 M HCl and swirl the flask.

3. A purple solution should form, indicating a mixture of pink and
blue. If the solution appears pink, add more HCl; if it is blue, add
more distilled water.

4. Divide the purple solution into 3 separate test tubes.

(a) Leave one test tube at room temperature.
(b) Place the second test tube in an ice bath.
(c) Place the third test tube in a water bath at 80 – 90 °C.

5. Record the colour of the solution in each test tube. Remove the
second and the third test tubes and leave them at room
temperature. Observe the change in colour.

POINT TO DISCUSS

Determine whether the forward reaction is exothermic or endothermic.
Discuss.

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(C) Determination of the equilibrium constant.
The following reaction is an example of a heterogenous system:

SbCl3(aq) + H2O(l) SbOCl(s) + 2HCl(aq)

The expression for the equilibrium constant is:

Kc = [HCl]2
[SbCl 3 ]

1. Pipette 5.0 mL of 0.5 M SbCl3 in 6 M HCl into a conical flask.

2. Carefully add distilled water from a burette into the conical flask
while swirling until a faint white precipitate is obtained.

3. Record the volume of water added.

4. Calculate the value of the equilibrium constant, Kc.

POINT TO DISCUSS

Explain why the concentration of pure liquid and solid are excluded from the
equilibrium constant expression for a heterogeneous system.

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DATA SHEET
EXPERIMENT 5: CHEMICAL EQUILIBRIUM

RESULT

(A) The effect of concentration in the formation of thiocyanoiron(III)
complex ion

Test Experiment Observation
tube

I Addition of 1 mL of 0.1M
Fe(NO3)3

II Addition of 1 mL of 0.1M
KSCN

III Addition of 6-8 drops of
10% NaOH

(B) The Effect of Temperature

Test Experiment Observation
tube

I At room temperature

II 0 ⁰C (in ice bath)

III 80- 90 ⁰C

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When left at room temperature: Observation
Test
tube

II

III

(C) Determination of the equilibrium constant

Burette reading / mL Observation
Initial reading
Final reading
Volume of water

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EXPERIMENT 6: pH MEASUREMENT AND ITS APPLICATIONS

Objectives

At the end of this lesson, students should be able to:
i. use various methods to measure the pH of acids, bases and salts.
ii. determine the dissociation constant, Ka, of acetic acid.

Introduction

pH is a measure of acidity or basicity of a solution. pH is defined as the
negative logarithm of hydrogen ion concentration, [H+].

pH = − log [H+] ……….(1)

The pH scale ranges from 0 to 14. At 25°C, a neutral solution has a pH of 7.
An acidic solution has a pH of less than 7 while a basic solution has a pH
greater than 7.

There are two methods to determine pH in the laboratory. The first method
involves the use of indicators such as pH paper and the universal indicator.
The second method is using the pH meter.

Acids or bases which ionise completely are called strong acids or strong
bases. An example of a strong acid is HCl and a strong base is NaOH. Weak
acids and weak bases do not ionise completely. An example of a weak acid is
acetic acid, CH3COOH, and that of a weak base is ammonia, NH3.

Consider the ionisation of a weak acid, HA.

HA(aq) H+(aq) + A-(aq) ……….(2)

The equilibrium constant expression for the above reaction is written as:

Ka = [H+ ][A− ] ……….(3)
[HA]

where [H+], [A−] and [HA] represent the molar concentrations of species that

exist at equilibrium. Kais the dissociation constant for acid HA. A similar

expression of Kb can be written for weak bases.

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One of the methods to determine Ka is by adding a weak acid solution to its
conjugated base solution. The product of this process is an acidic buffer
solution. The conjugated base is obtained from the salt produced using the
titration method.

In this method, a known weak acid, HA is divided into two equal portions, X
and Y. The first portion, X is titrated with NaOH solution using
phenolphthalein as an indicator to detect the formation of a salt solution. A
change in colour, from colourless to light pink, indicates the end point. The
equation for the reaction is:-

OH−(aq) + HA(aq) A−(aq) + H2O(l) ……….(4)

In this reaction, HA reacts with NaOH to form NaA and H2O. NaA ionises
completely to form A− and Na+. The number of moles of A− formed is the
same as the number of moles of HA in the second portion, Y, which has not
been titrated.

The second portion of the weak acid HA is added to the conical flask
containing the salt NaA. In this mixture, the concentration of HA is equal to
the concentration of A− from the salt.

Since [A−] = [HA], and from Equation 3,
Ka = [H+]

The value of [H+] is obtained by measuring the pH; hence the value of Ka can
be calculated.

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Apparatus Chemical Reagents

Burette Methyl red
Test tube Methyl orange
pH meter Alizarin yellow
25 mL pipette Phenolphthalein
250 mL conical flask Universal indicator
0.1 M NaCl
0.1 M NH4NO3
0.1 M CH3COONa
0.1 M and 1.0 M NH3
0.01 M and 1.0 M HCl
0.1 M and 1.0 M CH3COOH
0.1 M, 0.2 M and 1.0 M NaOH

Procedure

(A) Determination of pH of acidic and basic solutions

1. (a) Place 2 mL of the following solutions into separate
test tubes.

i. 0.01 M HCl
ii. 1.0 M HCl
iii. 0.1 M CH3COOH
iv. 1.0 M CH3COOH
v. 0.1 M NaOH
vi. 0.1 M NH3

Use pH paper to determine the pH of the solutions.

(b) Use a pH meter to determine the pH of the following
solutions:

i. 0.01 M HCl
ii. 1.0 M HCl
iii. 0.1 M CH3COOH
iv. 1.0 M CH3COOH

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2. Fill the test tubes with 2 mL of each of the following solution:

i. 0.01 M HCl
ii. 0.1 M CH3COOH
iii. 0.1 M NH3

Add two drops of methyl red to each test tube. Record the
observation. Determine the pH range by comparing the colour
of the solutions with the chart provided.

Repeat step 2 with methyl orange.

3. Fill the test tubes with 2 mL of each of the following solution:

i. 0.1 M NaOH
ii. 1.0 M NaOH
iii. 0.1 M NH3
iv. 1.0 M NH3

Add two drops of alizarin yellow to each test tube. Record the
observation. Determine the pH range by comparing the colour
of the solutions with the chart provided.

(B) Determination of pH of salt solutions

1. Fill the test tube with 2 mL of each of the following solution:

i. 0.1 M NaCl
ii. 0.1 M CH3COONa
iii. 0.1 M NH4NO3

Using pH paper and universal indicator, determine the pH and
state whether the salt solutions are acidic, basic or neutral.

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(C) Determination of the dissociation constant of a weak acid, Ka

1. Pipette 25 mL of 0.1 M CH3COOH into two conical flasks, X
and Y.

2. Add 2 - 3 drops of phenolphthalein into the conical flask X
and titrate it with 0.2 M NaOH. When the volume of base
reaches 10 mL, add the titrant drop by drop. The end point is
reached when the solution becomes pink. Record the initial
and the final readings of the burette.

3. Mix the solution in step 2 with 25 mL of 0.1 M CH3COOH in
the conical flask Y. Determine the pH of this mixture using a
pH meter.

4. Calculate Ka from the value of pH obtained in step 3.

POINT TO DISCUSS

1. Calculate the percentage of ionisation of 0.1 M and 1.0 M acetic acid.
How does the percentage of ionisation change with its concentration?

2. Refer to the pH value of acetic acid in Part (A). Calculate its Ka and
compare this value to that obtained from Part (C).

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DATA SHEET
EXPERIMENT 6: pH MEASUREMENT AND ITS APPLICATIONS
RESULT
(A) pH measurements of acid and base solutions

Solution pH Indicator
paper
0.01 M pH Methyl Alizarin Methyl
HCl meter red Yellow orange

1.0 M -

---

0.1 M -
CH3COOH
---
1.0 M

0.1 M -- -
1.0 M
NaOH

- -- -

0.1 M -
NH3
- -- -
1.0 M

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(B) pH measurements of salt solutions

Salt solution pH paper Universal Acidic/ Basic/
NaCl 0.1 M Indicator Neutral

CH3COONa 0.1 M

NH4NO3 0.1 M

(C) Dissociation constant of a weak acid, Ka Data
Burette reading / mL

Final reading
Initial reading
Volume of NaOH used / mL

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