Chemistry Laboratory Manual (SK015)

SK015 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.
(PEO2, 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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SK015 Lab Manual

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, PLO 1, MQF LOD 1)

2. Demonstrate the correct techniques in handling laboratory
apparatus and chemicals when carrying out experiments.
{P4, PLO2, MQF LOD2)

3. Solve chemistry related problems by applying basic concepts and
principles in physical chemistry.
(C4, PLO4, CTPS3, MQF LOD6)

13.2 Chemistry 2
At the end of the course, student should be able to:

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

2. Demonstrate the correct techniques in handling laboratory
apparatus and chemicals when carrying out experiments.
(P4, PLO2, MQF LOD2)

3. Solve chemistry related problems by applying basic concepts
and principles in physical and organic chemistry.
(C4, PLO4, CTPS3, MQF LOD6)

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 analyzing 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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SK015 Lab Manual

2.0 Laboratory Safety

The Science Matriculation Programme requires the students to attend practical classes
2 hours a week to complete 6 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 unauthorized 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 regent 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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SK015 Lab Manual

18. Wash all glass wares 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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SK015 Lab Manual
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SK015 Lab Manual

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 observation made during

the experiment. s

iv. Keep the work area clean and tidy.

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SK015 Lab Manual
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SK015 Lab Manual

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 • write concisely in your own words
• draw and label diagram if necessary

Procedure • write in passive sentences about all the
steps taken during the experiment

Results/ • data tabulation with units and uncertainties
Observation
• data processing (plotting graph, calculation to
obtain the results of the experiments and its
uncertainties)

Discussion • 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

Conclusion • state briefly the results with reference to the
objectives of the experiment

Reminder : NO PLAGIARISM IS ALLOWED.

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CHEMISTRY 1

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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. (P4, 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. determine 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 6MHCI
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 0.25g 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 HCI solution in to 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 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 20mL.

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.

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. (P4, 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 HCI 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 HCI 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(/) …(1)
HCI(aq) + NaOH(aq)
NaCl(aq) + H2O(g) ...(2)

Apparatus Chemical reagent

Burette x M HCI
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
250mL 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 250mL 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 acids 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 endpoint.

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 HCI solution.
1. Pipette 25 mL of a given HCI 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 HCI.

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

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EXPERIMENT 2 ACID-BASE TITRATION DATA SHEET
III
RESULTS III

(A) Preparation of standard oxalic acid solution

Exact mass of hydrated oxalic acid =
Moles of hydrated oxalic acid =
Molarity of oxalic acid =

(B) Standardisation of 0.2 M NaOH solution

Burette reading / mL Gross I II
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 HCI solution

Burette reading / mL Gross I II
Final reading
Initial reading
Volume of NaOH used / mL

Average volume of NaOH used =
Calculate the molarity of the HCI 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.(P4,PLO2,MQFLOD2)

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

Are active 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, HCI 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
Distilled water
Scissors Phenolphthalei
White tile Dilute hydrochloric acid, HCI
Pipette filler 0.1 M Sodium hydroxide, NaOH,
Filter funnel 0.2 An unknown alkaline earth metal, M
Retort stand
50 mL beaker 7
50 mL burette
25 mL pipette
Analytical balance
250 mLconical flask
Abrasive cloth no.3 (36)
Aluminium oxide

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Procedure
(A) Standardization of HCI 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 HCI 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 HCI 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 HCI 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 HCI 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 AMETAL
RESULTS

1. Titration of standard HCI solution

Concentration of NaOH = M
Volume of HCI = mL

Burette reading / mL Gross I II III
Final reading
Initial reading
Volume of NaOH used / mL

Average volume of NaOH =

2. Reaction of metal and HCI

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

3. Titration of unreacted HCI

Burette reading / mL Gross I II III
Final reading
Initial reading
Volume of NaOH used / mL

CALCULATION
1. Calculate the molarity of the standard HCI solution.
2. Determine the number of moles of HCI in 25 mL of the standard solution.
3. Calculate the number of moles of the unreacted HCI 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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EXPERIMENT4 CHARLES’ LAW AND THE IDEAL GAS LAW

Course Learning Objective

Demonstrate the correct techniques in handling laboratory apparatus and chemicals when
carrying out experiments. (P4, 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 α 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 ai is obtained by multiplying the cross-sectional area of
the tube, A, with the height of the air column, h.

V=Axh

Assuming that the cross-sectional area is constant, the volume is directly proportional to the
height ,i.e. , V α h. Therefore, the height of the air column can be used as ameasur eof 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 α P ( n, T constant )

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

V α 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 α T (n, P constant)

Thus, combining the three laws, we get

Vα nT
P

The above expression can be written as PV= nRT .(1)
V = RnT or
P

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

Therefore, the ideal gas equation can also be written as

PV = m RT .(2)
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 equivaIent 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.

Figure4.2 Figure4.3

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.

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
iquid 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 TABLE1.

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
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 (°C)

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. (P4,PLO2,MQFLOD2)

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, K„ 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, once 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 invoIve 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 away so as to counter act 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 or
phases such as precipitation or dissolution.

Apparatus Chemical reagents

Burette 6MHCI
Ice bath 0.3 M CoCI
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 SbCI3 in 6 M HCI
Conical flask (100mL)
Measuring cylinder( 10mL and 100mL)

Procedure

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

The thiocyano iron (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 5mL each of this solution into four test tubes.
(a) To the first test tube, add 1mL 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 hexaaquocobaIt(II) complex ion with chloride ion
produces tetrachIorocobaIt(II) ion. The equation for the reaction is given below:

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

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

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

3. A purple solution should form, indicating a mixture of pink and blue. If the
solution appears pink, add more HCI; 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.

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(/) SbOCl(s) + 2HCl(aq)

The expression for the equilibrium constant is

[HCI]2
[SbCl3]

Procedure
1. Pipette 5.0 mL of 0.5 M SbCI3 in 6 M HCI 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. (P4, 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, K„ 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 invoIves the use of
indicators such as pH paper and the universal indicator. The second method is using the pH
meter.

Acids or bases which ionize completely are called strong acids or strong bases. An example of
a strong acid is HCI and a strong base is NaOH. Weak acids and weak bases do not ionize

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:

= [H+][A−] ……………………….(3)
Ka[HA]

where [H+], [A-] and [HA] represent the molar concentrations of species that exist at for
equilibrium. Ka is the dissociation constant 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

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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(g) ................................. (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 vaIue 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 NaCI
0.1 M NH4NO3
0.1 M CH3COONa
0.1 M and 1.0 M NH3
0.01 M and 1.0 M HCI
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 HCI
ii. 1.0 M HCI
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 HCI
ii. 1.0 M HCI
iii. 0.1M CH3COOH
iv. 1.0M CH3COOH

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

i. 0.01 M HCI
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.

0.2 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.

2. 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.

3. 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 value
to that obtained from Part (C).

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REFERENCES

Ali,R.(1995) Panduan Amali Kimi aAsas, Kursus Pengajian Tingg iFajar Bakti,
Selangor.

Baum, S.J., Sandwick, R.K. (1994) Laboratory Exercises in Organic and
Biological Chemistry. Prentice Hall. New Jersey. United States of America.

Beran, J.A. (1996) A Study of Chemical and Physical Changes, 2’ d

Edition. John Wiley & Sons Inc. United States of America.

Chemistry Department of University Malaya. (2001) Laboratory Manual
Organic Chemistry (SCES1220). Universiti Malaya. Malaysia.

Ritchie,R.(2000)Revise AS Chemistry.Letts Educational Ltd..United
States of America.

Ryan, L. (1996) Chemistry for You, Stanley, horpes (Publishers) Ltd.
England.

Seager,S.L.,Slabaugh,M.R.(2000)Introductory Chemistsr or Today,
4th Edition.ThomsonLearning.California.UnitedStatesofAmerica.

Stanley, A.J. et.aI (2000) Discovering Chemistry . A Year-12 Chemistry
Text Book. Open book Publishers. South Australia, Australia.

Universiti Teknologi Malaysia (2001) Amafi Kimia Am, Jawatankuasa
Penerbitan dan Penulisan Fakulti Sains UTM. Penerbit UTM. Malaysia.

Ware, G., Deretic, G. (1995) Senior Chemistry : Practical Manual,
Heinemann. Victoria.

Chang,R. (2005) Chemistry, 8" Edition. MacGraw Hill. United States of
America.

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ACKNOWLEDGEMENTS

The Matriculation Di\/ision, Ministry of Education wishes to extend heartfelt thanks and outmost
gratitude to the following individuals who have contributed valuable input, support and
suggestions in completing this manual.

A very special thanks and appreciation are due to these individuals:

Reviewers Director of Matriculation
Division, Ministry OfEducation
Dr. Baiduriah bintiYaakub
Deputy Director,
En.AzmanbinAbd.Karim MatriculationDivision, Ministry
OfEducation
Dr. Shah Jahan binAssanarkutty
Senior Principal
Dr. Saharawati binti Shahar AssistantDirector,
MatriculationDivision
Prof.Dr.Zanariah binti Abdullah Principal
Prof.Dr. Rosiyah binti Yahya AssistantDirector,
Noor Hayati binti Abu Bakar MatriculationDivision
Rusiati binti Md.Som Gan Fie Universiti Malaya
Chuen
Nor Hayati binti Abu Bakar Universiti Malaya
Dani Asmadi bin Ibrahim
Wan NorIzana binti Wan Mohammad Kolej Matrikulasi Melaka
Noor Fatihah binti Zulkeply
Khatijah binti Ali Kolej Matrikulasi Negeri Sembilan
Fauziah binti lsmail Nor Kolej Matrikulasi Negeri Sembilan
Aziah binti Lazim Kolej Matrikulasi Pulau Pinang
Maizuyah binti Omar Kolej Matrikulasi Negeri Sembilan
Siti Warda binti Selamat Kolej Matrikulasi Negeri Sembilan
lsharae bin Abdul Mosamad Kolej Matrikulasi Neqeri Sembilan
Kolej Matrikulasi Johor
Kolej Matrikulasi Perak
Kolej Matrikulasi Selangor
Kolej Matrikulasi Selangor
Kolej Matrikulasi Selangor
Bahagian Matrikulasi
KPM