LAB MANUAL SK015 & SK025

MATRICULATION DIVISION

CHEMISTRY
LABORATORY MANUAL
SEMESTER I & II
SK015 & SK025

TWELFTH EDITION

MATRICULATION DIVISION
MINISTRY OF EDUCATION MALAYSIA

CHEMISTRY

LABORATORY MANUAL
SEMESTER I & II
SK015 & SK025

TWELFTH 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)
Copyright © 2020 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 : 603-88844083
Fax : 603-88844028
Website : http://www.moe.gov.my/bmkpm

Printed in Malaysia by Cataloguing-in-Publication Data

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

eISBN: 978-983-2604-50-1

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,
emotionally 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.



CONTENTS Page
i
Learning Outcomes
Introduction iii - iv
v
 Laboratory Safety vi
 Preparation For Experiment
 Report Writing

Semester I

Experiment Title

1 Determination of the formula unit of a compound 1
3
2 Acid Base Titration – Determination of the concentration of 7
hydrochloric acid solution 11
16
3 Determination of the molar mass of a metal 20

4 Charles’ Law and the ideal gas Law

5 Chemical Equilibrium

6 pH measurement and its applications

Semester II 24
Experiment Title 28
31
1 Rate of reaction 34
2 Determining the heat of reaction 38
3 Electrochemical cells 41
4 Reactions of aliphatic and aromatic hydrocarbons
5 Reactions of hydroxy compounds 44
6 Aldehydes and ketones 45

References
Acknowledgements

SK015 & SK025 Lab Manual

1.0 Learning Outcomes

1.1 Matriculation Science Programme Educational Objectives

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

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

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

3. Able to solve scientific and mathematical problems innovatively and
creatively.

4. Able to engage in life-long learning with strong commitment to
continue the acquisition of new knowledge and skills.

1.2 Matriculation Science Programme Learning Outcomes

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

1. Acquire knowledge of science and mathematics fundamental in higher
level education.
(PEO 1, MQF LOD 1)

2. Demonstrate manipulative skills in laboratory work.
(PEO 1, MQF LOD 2)

3. Communicate competently and collaborate effectively in group work
with skills needed for admission in higher education institutions.
(PEO 2, MQF LOD 5)

4. Apply logical, analytical and critical thinking in scientific studies and
problem solving.
(PEO 3, MQF LOD 6)

5. Independently seek and share information related to science and
mathematics.
(PEO 4, MQF LOD 7)

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1.3 Course Learning Outcome

13.1 Chemistry 1

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

1. Explain basic concepts and principles of physical chemistry in
novel and real life situations.
(C2, PLO1, MQF LOD1)

2. Demonstrate the correct techniques in handling laboratory
apparatus and chemicals when carrying out experiments.
(P3, PLO2, MQF LOD2)

3. Solve chemistry related problems by applying the basic concepts
and Principles in physical chemistry.
(C4, PLO4, CTPS3, MQF LOD 6)

13.2 Chemistry 2

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

1. Explain basic concepts and principles of organic chemistry in
novel and real life situations.
(C2, PLO1, MQF LOD1)

2. Demonstrate the correct techniques in handling laboratory
apparatus and chemicals when carrying out experiments.
(P3, PLO2, MQF LOD2)

3. Solve chemistry related problems by applying the basic concepts and
principles in organic chemistry.
(C4, PLO4, CTPS3, MQF LOD 6)

1.4 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.

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2.0 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.

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.

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.

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18. Wash all glasswares 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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3.0 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.

4.0 Preparation for experiment
4.1 Pre-lab Sessions.
i. Read and understand the objectives and the theory of the experiment.
ii. Think and plan the working procedures properly for the whole
experiment. Make sure you have appropriate table for the data.
iii. Complete and submit the pre-lab questions provided.

4.2 Practical Sessions
i. Check the apparatus provided.

ii. Conduct the experiment carefully.

iii. Record all measurements and observations made during the
experiment.

iv. Keep the work area clean and tidy.

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4.3 Post-lab Sessions
i. Explain what has been carried out and discuss the findings of the
experiment.
ii. Introduce the format of report writing as below:

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

Conclusion  data tabulation with units and 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

SK015 & SK025 Lab Manual

EXPERIMENT 1 DETERMINATION OF THE FORMULA UNIT OF A COMPOUND

Course Learning Objective

Demonstrate the correct techniques in handling laboratory apparatus and chemicals when carrying
out experiments. (P3, PLO 2, MQF LOD 2)

Learning Outcomes

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

Student Learning Time (SLT)

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

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 mole of silver and 0.1 mole 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)

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.

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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.

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

4. If the zinc powder does not dissolved 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.

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.

EXERCISE

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

Course Learning Objective

Demonstrate the correct techniques in handling laboratory apparatus and chemicals when carrying
out experiments. (P3, PLO 2, MQF LOD 2)

Learning Outcomes

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

Student Learning Time (SLT)

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

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.

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.

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

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.

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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.

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.

EXERCISE

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

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DATA SHEET

EXPERIMENT 2 ACID-BASE TITRATION

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) Standardisationof 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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EXPERIMENT 3 DETERMINATION OF THE MOLAR MASS OF A METAL

Course Learning Objective

Demonstrate the correct techniques in handling laboratory apparatus and chemicals when carrying
out experiments. (P3, PLO 2, MQF LOD 2)

Learning Outcomes

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.

Student Learning Time (SLT)

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

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.

Apparatus Chemical Reagents

Scissors Distilled water
White tile Phenolphthalein
Pipette filler
Filter funnel Dilute hydrochloric acid, HCl
Retort stand 0.1 M Sodium hydroxide, NaOH,
50 mL beaker An unknown alkaline earth metal, M
50 mL burette
25 mL pipette 25 mL
Analytical balance
250 mL conical flask
Abrasive cloth no.3 (36)
Aluminium oxide

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Procedure
(A) Standardization of HCl solution

1. Rinse a clean burette with 0.1 M NaOH.
2. Fill the burette with 0.1 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 persists for more than 30 seconds. This is the end point.
8. Record the final reading of the burette.
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.
5. Place the samples separately into the HCl solution. Swirl occasionally
until the metal is completely dissolved.
6. Add 2 drops of phenolphthalein.
7. Record the initial burette reading.
8. Titrate the unreacted HCl with the NaOH solution.
9. Record the final burette reading.
10. Repeat titration with the other sample.

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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 Sample I Sample II

Burette reading / mL
Final reading
Initial reading
Volume of NaOH (mL)

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:

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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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EXPERIMENT 4 CHARLES’ LAW AND THE IDEAL GAS LAW

Course Learning Objective

Demonstrate the correct techniques in handling laboratory apparatus and chemicals when carrying
out experiments. (P3, PLO 2, MQF LOD 2)

Learning Outcomes

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

Student Learning Time (SLT)

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

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.

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

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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

nT
Vp

The above expression can be written as

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

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

Apparatus Chemical reagents

Needle Ice
Wire gauze Methanol
Tripod stand Unknown liquid
Rubber band
Thermometer
Bunsen burner
Aluminum foil
Beaker (600 mL)
Analytical balance
Open tube manometer
Retort stand and clamp
Charles’ law apparatus
Conical flask (100 mL)
Measuring cylinder (100 mL)

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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.

Figure 4.1
Charles’ law apparatus

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 – 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.

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(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

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. 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

(A) Charles’ law TABLE 1 Volume
Temperature (Height of gas column)
Condition
Warm water
Tap water
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.

(B) Determination of the molar mass of the gas Reading

TABLE 2

No Item
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)

1. Complete TABLE 2.
2. Calculate the molar mass of the unknown liquid.

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EXPERIMENT 5 CHEMICAL EQUILIBRIUM

Course Learning Objective

Demonstrate the correct techniques in handling laboratory apparatus and chemicals when carrying
out experiments. (P3, PLO 2, MQF LOD 2)

Learning Outcomes

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.

Student Learning Time (SLT)

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

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.

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.

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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 o r
phases such as precipitation or dissolution.

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
Pipette (10 mL) 0.1 M Fe(NO3)3
Beaker (100 mL) 0.5 M SbCl3 in 6 M HCl
Conical flask (100 mL)
Measuring cylinder (10 mL and 100 mL)

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.

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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.

(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 – 90oC.

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.

EXERCISE
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 ]

Procedure

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.

EXERCISE

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

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EXPERIMENT 6 pH MEASUREMENT AND ITS APPLICATIONS

Course Learning Objective

Demonstrate the correct techniques in handling laboratory apparatus and chemicals when carrying
out experiments. (P3, PLO 2, MQF LOD 2)

Learning Outcomes

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.

Student Learning Time (SLT)

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

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.

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.

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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 H O. NaA ionises completely to form
2
+
A- and The number of moles of A- formed is the same as the number of moles of HA in
Na .

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.

Apparatus Chemical reagents

Burette pH paper
pH Meter Methyl red
Test tube Methyl orange
25 mL pipette Alizarin yellow
250 mL conical flask Phenolphthalein
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.

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(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

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.

EXERCISE

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 valueto
that obtained from Part (C).

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CHEMISTRY 2
SK025

SK015 & SK025 Lab Manual

EXPERIMENT 1 RATE OF REACTION

Course Learning Objective
Demonstrate the correct techniques in handling laboratory apparatus and chemicals when carrying
out experiments. (P3, PLO 2, MQF LOD 2)

Learning Outcomes

At the end of this lesson, students should be able to study the effect of concentration,
temperature and catalyst on the reaction rate

Student Learning Time (SLT)

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

Introduction

The reaction rate is the change in concentration of the reactants or products per unit time.
The factors that influence the rate of reaction are temperature, pressure, catalyst, size of
particles and concentration of reactants.

The rate of a reaction can be studied by observing the change in the chemical or physical
properties of species involved in the reaction. The reaction rate is inversely proportional to
the time of the reaction, i.e. the faster the reaction occurs, the shorter is the time for the
reaction to complete.

Apparatus Chemical reagents

Glass rod 0.1 M HCl
Water bath 10% MnSO
Stopwatch
Boiling tube 4
Thermometer
Pipette (10 mL) 2.0 M H SO
Burette (50 mL) 24
Conical flask (100 mL)
Laminated white paper with ‘X’ mark 0.02 M KMnO
Measuring cylinder (10mL) 4

0.2 M Na S O
22 3

0.25 M H C O
224

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Procedure

(A) The effect of concentration on the reaction rate

1. Place 50 mL of 0.2 M sodium thiosulphate, Na S O using a burette into a 100 mL
22 3
conical flask. Put the conical flask on the white paper with ‘X’ mark.

2. Pipette 10 mL of 0.1 M HCl into the conical flask and immediately start the
stopwatch. Stir continuously with a glass rod until the mark is no longer visible
and record the time.

Note: The ‘X’ mark should be observed from the top of the conical flask.

3. Repeat steps 1-2 with the addition of distilled water to the sodium thiosulphate as
instructed in Table 1.1.

Table 1.1
Concentration of reactant

Volume of Volume of Concentration Volume of Time 1
0.2 M Na S O distilled of 0.1 M HCI (s)
water t
22 3 (mL) Na S O (M) solution
22 3 (mL) (s-1)
solution 0.00
(mL) 10.00
50.00
10.00 10.00
40.00
20.00 10.00
30.00
30.00 10.00
20.00
40.00 10.00
10.00

4. Calculate the concentration of the sodium thiosulphate solution after the dilution

and the value of 1 .
t

1

5. Plot a graph of against the concentration of sodium thiosulphate solution.

t

6. Based on the graph, state the relationship between the concentration of the
sodium thiosulphate solution with time and the rate of reaction.

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(B) The effect of temperature and catalyst on the reaction rate

1. Label 4 boiling tubes as A1, A2, B1 and B2.

2. Place 10 mL of 0.25 M oxalic acid, H C O solution into boiling tubes A1 and A2.
224

3. Fill boiling tubes B1 and B2 with 5 mL of 0.02 M KMnO4 solution. Then add 10 mL
of 2.0 M H2SO4 solution to both tubes.

4. Add 5 drops of 10% MnSO solution to B . Stir the mixture.
42

5. Place tubes A1 and B1in a water bath at temperature of 30C for about 3 minutes.

6. While tube A1 is still in the water bath, pour the solutions from tube B1 into tube
A1. Start the stopwatch immediately.

7. Record the time taken for the mixture to turn colourless.

8. Repeat Steps 5 - 7 for tubes A2 and B2.

9. Follow Steps 2-7 for the temperatures of 35C, 40C and 50C. Record your
results in Table 1.2.

Table 1.2
Effect of temperature and catalyst on reaction rate

Temperature Without catalyst MnSO4 With catalyst MnSO4
(C ) (A1 + B1) (A2 + B2)

30 t (s) 1 (s-1) t (s) 1 (s-1)
35 t t
40
50

1

10. Plot against the temperature for the mixtures of A1+ B1 and A2 + B2solutions on

t

the same graph.

11. Based on the graph, deduce the relationship between

i. temperature and rate of reaction.

ii. catalyst and rate of reaction.

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EXERCISE
1. What is the function of the catalyst in the above reactions?

1

2. What does represent?

t

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EXPERIMENT 2 DETERMINING THE HEAT OF REACTION

Course Learning Objective

Demonstrate the correct techniques in handling laboratory apparatus and chemicals when carrying
out experiments. (P3, PLO 2, MQF LOD 2)

Learning Outcomes

At the end of this lesson, students should be able to:
i. determine the heat capacity of a calorimeter.
ii. determine the heat of neutralisation of HCl and NaOH

Student Learning Time (SLT)

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

Introduction

Heat released or absorbed during chemical reactions can be measured by using a
calorimeter. A calorimeter is a container that is thermally isolated from the environment. Heat

released by the chemical reaction, -q is absorbed by the solution and the calorimeter.
rx
-q = q + q (1)
rxn s c

where
q = heat absorbed by solution

s

q = heat absorbed by calorimeter
c

The heat absorbedby a calorimeter is proportional to the change in temperature. The
proportionality constant, C, is known as the heat capacity of a calorimeter. Heat capacity is
defined as the amount of heat required to increase the temperature of the calorimeter by
1C.

qc = C∆T ……….(2)

For a solution, the heat absorbed is proportional to the mass of the solution and the change

in temperature. The constant, c, is known as the specific heat capacity of solution per unit

mass. The specific heat capacity of a very dilute solution is equivalent to the specif ic heat

-1 -1

capacity of pure water ,4.18J g C The mass of the solution can be calculated by assuming

the density of the solution is the same as the densityof water.

qs = mscs ∆ T …………(3)

Heat released can be determined by measuring the temperature before and after the
reaction.

-q = C ∆T + m c ∆T ……………(4)
rxn c ss

where final temperature of system – initial temperature of system
∆T = mass of solution
heat capacity of calorimeter
m=
s specific heat capacity of solution

C=
c

c=
s

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Apparatus Chemical reagents

Pipette (25 mL) 1.0 M HCl
Beaker (100 mL) 1.0 M NaOH
Thermometer
Calorimeter or styrofoam cup

Procedure
(A) Determination of the heat capacity of a calorimeter

1. Set up a simple calorimeter as shown in Figure 2.1.

Figure 2.1
A simple calorimeter (Chang, 2005)

2. Measure the temperature, T1, of an empty calorimeter.

3. Pipette 50 mL of distilled water into a 100 mL beaker.

4. Heat the beaker to a temperature between 50 - 60°C.

5. Pour the hot water into the calorimeter. Close the lid immediately and measure
the initial temperature of the hot water, T2.

6. Observe the decrease in temperature every 10 seconds for 2 minutes. Record
the temperature that remains constant, T3.

7. Determine the heat capacity of the calorimeter.

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(B) Determination of the heat of neutralisation of 1.0 M HCl and 1.0 M NaOH
1. Pipette 25 mL of 1.0 M NaOH solution into the calorimeter and 25 mL of 1.0 M
HCl solution into a beaker. Record the initial temperature of each solution.
2. Without removing the thermometer, lift the lid slightly and quickly pour the HCl
solution into the calorimeter.
3. Quickly replace the lid of the calorimeter.
4. Stir the solution and record the maximum temperature reached.
5. Calculate the heat of neutralisation.

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EXPERIMENT 3 ELECTROCHEMICAL CELLS

Course Learning Objective

Demonstrate the correct techniques in handling laboratory apparatus and chemicals when carrying
out experiments. (P3, PLO 2, MQF LOD 2)

Learning Outcomes

At the end of this lesson, students should be able to:
i. arrange Al, Zn, Mg, Fe and Cu in an electrochemical series.
ii. determine the Faraday’s constant by electrolysis of CuSO4 solution.

Student Learning Time (SLT)

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

Introduction

Electrochemistry is a study of the relationship between electricity and chemistry. Generally
there are two types of electrochemical cells, namely galvanic and electrolytic cells. A
galvanic cell is an electrochemical cell in which redox reaction occurs spontaneously to
generate electricity.For a galvanic cell, oxidation occurs at the anode and electrons flow to
the cathode where reduction occurs.

A standard reduction potential is defined as a reduction potential obtained at a standard
condition where the concentration of solution is 1.0 M, the gas partial pressure is 1 atm and
temperature is 25 °C. The standard reduction potential values are arranged in a certain
order and the list is known as the Standard Reduction Potential Table or the emf Series.

The potential difference between the two half cells in an electrochemical cell is called cell
potential. The cell potential or the cell voltage at the standard condition can be written as:

Eocell = Eocathode - Eoanode

The cell potential at non-standard condition can be calculated by using the Nernst equation.

Ecell  Eo  0.0592 logQ
cell n

In this experiment, the cell potential is obtained from the voltmeter reading. By inserting the

value and the concentration of the electrolyte in the Nernst equation, the standard c ell
potential, Eocell can be determined.

An electrolytic cell uses electricity to produce chemical changes in an electrolyte. The cell is

made up of two electrodes connected to a battery which functions as a source of direct
current. During electrolysis, cations are reduced at the cathode while anions are oxidised at
the anode. The amount of substance formed at each electrode can be predicted based on
Faraday’s first law.

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Apparatus Chemical reagents

Tongs 0.1 M CuSO4
Ammeter 0.1 M ZnSO4
Hair dryer 0.1 M FeSO4
Voltmeter 0.1 M MgSO4
Stopwatch 0.1 M Al(NO3)3
Transformer Zinc strip
Sand paper / abrasive cloth
Crocodile clips Copper strip
Beaker (50 mL) Magnesium strip
Analytical balance
Salt bridge Iron strip (nail)
Measuring cylinder (50 mL) Aluminium strip
Carbon rod

Saturated KNO3/KCl

Note: 1. Clean the electrodes with sand paper / abrasive cloth before use.

2. Ensure that the filter paper to be used as salt bridge is completely soaked in
saturated KNO3/KCl solution. Avoid handling the salt bridge with bare hands.

Procedure

(A) Galvanic cell

1. Clean the metal strips with sand paper / abrasive cloth.

2. Fill a 50 mL beaker with 35 mL of 0.1 M CuSO4 and the other beaker with 35 mL
of 0.1 M ZnSO4.

3. Set up the apparatus as shown in Figure 3.1.

Zn Cu
Salt bridge
CuSO4
ZnSO4

Figure 3.1
Galvanic cell

4. Record the cell potential.

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5. Repeat Steps 1 – 4 by replacing Zn2+/Zn half cell with a

(a) magnesium strip in 0.1 M MgSO4
(b) aluminium strip in 0.1 M Al(NO3)3
(c) iron strip in 0.1 M FeSO4

6. Arrange the metals in ascending order of strength as reducing agents.

7. Verify the above order by calculating the standard reduction potential, Eored, of
each electrode.

(B) Determination of Faraday’s constant

1. Clean a copper electrode with a piece of sand paper / abrasive cloth.

2. Weigh the copper electrode accurately.

3. Set up apparatus as show in Figure 3.2. Fill a 50 mL beaker with 35 mL 0.1 M
CuSO4.
+ DC −

Carbon Copper
(anode) (cathode)

0.1 M CuSO4

Figure 3.2
An electrolytic cell

4. Complete the circuit by connecting the wires from each electrode to the ammeter
and transformer. Set the transformer to supply the direct current with a voltage of
3 V.

5. Run the electrolysis for 15 minutes.
6. Record the ammeter reading and your observation of each electrode.
7. Disconnect the circuit and record the exact time of electrolysis.
8. Dry the copper strip using a hair dryer.
9. Weigh again the copper strip.
10. Calculate the mass of copper deposited. Determine the Faraday’s constant.

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EXPERIMENT 4 REACTIONS OF ALIPHATIC AND AROMATIC HYDROCARBONS

Course Learning Objective

Demonstrate the correct techniques in handling laboratory apparatus and chemicals when carrying
out experiments. (P3, PLO 2, MQF LOD 2)

Learning Outcomes

At the end of this lesson, students should be able to:
i. study the chemical properties of an alkane, alkene and arene.
ii. differentiate an alkane from an alkene and arene.

Student Learning Time (SLT)

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

Introduction

Hydrocarbons are organic compounds that contain only carbon and hydrogen. Alkanes
which are also known as paraffins are saturated hydrocarbons. They do not contain double
or triple bonds. Hence, alkanes are relatively inert to chemical reactions.

Example of alkanes:

H H HH
HC HCCH

H HH

Methane Ethane Cyclohexane

Alkanes undergo free radical substitution reaction.

CH4 + CH2Cl2 CH3Br + HBr
Br2 uv

Alkenes are unsaturated hydrocarbons with at least one double bond between two carbon
atoms.

Example of alkenes:

H2C CH2

Ethene Cyclohexene

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Alkenes can easily undergo addition reactions at the C=C bond. For example, alkenes
undergo hydrogenation and halogenation to form alkanes and dihalides, respectively.

+CH3CH=CH2 CH2Cl2

Br2 CH3CHBrCH2Br

Alkenes also react with potassium permanganate solution in two different conditions:

a. In basic medium to form a diol.

H2C CH2 KMnO4/OH- HH + MnO2
room temp HC CH (brown precipitate)

OH OH

b. In hot acidic medium to form a carboxylic acid.

H3C CH3 KMnO4/H+ 2 H3C O
C C C

H H OH

Arenes are aromatic hydrocarbons with stable molecular structures.

Example of aromatic hydrocarbons:

CH3

Toluene Naphthalene Anthracene

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Although arenes have a very high degree of unsaturation, they are relatively inert towards all
addition reactions except at a very high pressure and temperature.

Ni

+ H2
high pressure, 200 °C

Arenes undergo electrophilic aromatic substitution reactions in the presence of a Lewis acid
catalyst.

+ Br2 FeBr3 Br

+ HBr

Apparatus Chemical reagents

Dropper Toluene
Test tube Cyclohexane
Rubber band Cyclohexene
Labeling paper Dichloromethane
Test tube rack 0.01 M KMnO4
Black sugar paper (6 x 12 cm) 4% bromine in dichloromethane

Procedure

(A) Reaction with bromine in dichloromethane

1. Label 6 dry, clean test tubes, A to F.

2. Place 1 mL of cyclohexane in test tubes A and B, 1 mL of cyclohexene in test
tubes C and D, and 1 mL of toluene in test tubes E and F.

3. Wrap test tubes A, C and E with black sugar papers.

4. Add 4 to 5 drops of 4% bromine in dichloromethane into each test tube.

5. Keep test tubes A, C and E in a dark place, and test tubes B, D and F in the
sunlight. Leave them for 15 minutes.

6. Record the observations.

Updated: 12/03/2020 36