Physical chemistry experiments

UTAR

FHSC1114 Physical Chemistry

Trimester 1

Lab manual version 4.0

Foundation in Science

1

Topic 1: Introductory to laboratory safety and apparatus

________________________________________________________________________

1. Materials requirement There are various materials that students must bring along during laboratory session.

Students will be asked to leave the laboratory if they fail to bring the items listed.

Laboratory manual Appropriate PPE (Personal Protective Equipment) Protective laboratory coating (Lab Coat) to protect you and your attire. Covered shoes to protect you from chemical burn or other hazards. Open toes

shoes are strictly not allowed.

Safety glasses or goggles to protect your eyes. Some other experimental materials which are requested by the lecturer. Calculator or scientific calculator. Record book to record observation and experiment results during lab sessions.

2. Practical Exercises To get the most out of the practical exercises, you are required to obey the

instructions given. These instructions have been designed to provide you with

experience in the following skills:

Following instructions Handling apparatus Having due regards for safety Making accurate observations Recording results in an appropriate form Presenting quantitative results Gives valuable discussion Drawing conclusions

Following Instructions

Follow the sequence of instructions as they are provided in an appropriate order.

Experimental procedures must be read through before carrying out the experiment. Draw

out the procedure in a flow chart for better understanding of the steps for the experiment.

Apparatus Handling

Before carrying out an experiment, it is important for scientists to plan and gather for

their experiment apparatus. As such, you are advised to list down the apparatus to be used

for the entire experiment before the start of the experiment. You will be able to master

the basic technique of using different types of apparatus. These include measuring

cylinder, bulb pipette, graduated glass pipette, volumetric flasks and burette. For different

type of tasks, different set of the apparatus are to be used.

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Laboratory Safety

Always move slowly and carefully in the lab.

During and after practical session, never touch your mouth or eyes with fingers before thorough wash of hands with soap and water.

Make sure glass objects (e.g: thermometer or glass rod) are not placed unattended on the bench to prevent it from being rolled off from the working bench to the floor.

Always put on your PPE such as lab coat, safety glasses and wear covered shoes in the lab.

Remarks:

Proper heating using Bunsen burner (even with heating of water).

Proper handling of any liquids, particularly those identified as corrosive, irritant, toxic or harmful.

Careful handling of corrosive work.

Allow hot plate, Bunsen burner, tripods, gauzes and beakers to cool down before handling them.

Keep long hair tied and do not wear dangly earrings.

Do not allow electrical equipment to come into contact with water.

If you are not sure about how to carry out a scientific procedure, ask the lecturer or lab officer.

Make sure you understand the rationale and consequences of your actions before you act.

Follow all safety instructions given in the manual or provided by the lecturer/tutor for particular experiments (e.g. use of gloves or mask).

Making Accurate Observation

The experiment will make it clear about the needed observation, e.g. the color changes

when two solutions are added together or time taken for a chemical reaction. Ensure that

you know the proper handling of relevant equipment before the start of experiment.

Think carefully about the precision of your observations. You may need to find out

reference for color description from external sources (e.g. reference books or online

references) before the practical session.

Recording Results in an Appropriate Form

Results can be recorded in various ways. Often it is helpful to record raw data in a table.

Most data will be in the form of numbers, e.g. quantitative data (also known as numerical

data). However, some data, e.g. color of solution, are qualitative. Bear in mind that the

best way for data collection is to avoid missing out any observation you have made, and

keep your raw data in safe hand.

Presenting Quantitative Results

Presentation of data can be made using table, graph or other visual means to ease result

analysis. You will have to choose the best way to present the experimental results.

Drawing Conclusion

Conclusions should be drawn from and supported by experimental results.

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3. Writing a Scientific Report

Title

Refers to the subject investigated.

Introduction

State the hypotheses

Give well-defined reason(s) for making hypotheses.

Explain the chemical basis of the experiment.

Cite sources to substantiate background information.

Explain how the method used will produce information relevant to the hypotheses.

State a prediction based on the hypotheses.

Material and Methods

Use appropriate format.

Give enough details (so that the experiment can be duplicated).

State the control treatment, replication, and standardized variables that were used.

Results

Summarize the data (do not include raw data).

Present the data in an appropriate format (table or graph).

Present tables and figures neatly so they are easily read.

Label the axes of each graph correctly.

Give units of measurement where appropriate.

Write a descriptive caption for each table and figure.

Include a short paragraph pointing out the important results (do not interpret the data).

Do not create your own data which is not true.

Discussion

State whether the hypotheses was supported or proven false by the results, or else state that the results were inconclusive.

Cite specific results that support your conclusions.

Compare the results, with your predictions and explain any unexpected results.

Compare the results with other research or information available.

Discuss any weaknesses the experimental design or problems with the execution of the experiment.

Discuss how you might extend or improve the experiment.

Conclusion

State conclusion which is supported by results

Restate important results.

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Literature Cited

Use proper citation form in the text.

Use proper citation form in the Literature Cited section.

Refer in the text to any source listed this section.

Acknowledgment

State any appropriate and necessary acknowledgment.

4. Apparatus and Equipments

Apparatus and Equipments Description and Function

Beaker

Beaker is a simple container for liquids, very commonly used in laboratories.

Beakers are generally cylindrical in shape, with a flat bottom.

Beakers are available in a wide range of sizes, from 1 ml up to several liters.

Conical flask/

Erlenmeyer flask

An Erlenmeyer flask (conical flask) is a type of widely used laboratory flask which features a conical base with a

cylindrical neck.

They are usually marked (graduated) on the side to indicate the approximate volume of their contents.

The conical flask's counterpart is the beaker. However the main difference is its narrow neck.

The neck allows the flask to be stoppered using rubber bungs or cotton wool.

The conical shape allows the contents to be swirled or stirred during an experiment (as is required in titration);

the narrow neck keeps the contents from spilling

Filtering funnel

A filtering funnel is a pipe with a wide, often conical mouth and a narrow stem.

It is used to channel liquid or fine-grained substances into containers with a small opening. Without a funnel, much

spillage will occur

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Measuring cylinder

A graduated cylinder also referred to as a measuring cylinder, type of laboratory glassware comprised of a tall

cylinder with a range of calibrated markings that is used

for visually measuring the volumes of liquids in a

quantitative manner.

A graduated cylinder can be made of translucent plastic or borosilicate glass.

Volumetric flask

A volumetric flask refers to graduated glassware used for the measurement of volume of liquids when the

bottom of the meniscus is perfectly aligned with mark on

the neck of the flask.

Pipette

Pipettes are used to transfer a specific volume of liquid.

2 most common used pipette in a science laboratory are: a) Volumetric pipette (Bulb pipette) b) Measuring pipette (Serological pipette)

Burette

Burette is a vertical cylindrical piece of laboratory glassware with a volumetric graduation on its full length

and a precision tap, or stopcock, on the bottom.

It is used to dispense known amount of a liquid reagent in experiments for which such precision is necessary, such as

a titration experiment.

Burettes are extremely precise: class A burettes are accurate to 0.05 ml.

Volumetric

pipette

Measuring

pipette

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Reagent bottles

Reagent bottles are used to hold small stocks of reagents and samples for use.

Reagent bottles must be clean but not necessarily dry.

Desiccators

Desiccators are sealable enclosures containing desiccants used for preserving moisture-sensitive items.

A typical desiccator contains two compartments separated by a perforated plate; the desiccant is placed at

the lower compartment while the reagent is placed in the

upper compartment.

It should not be used to dry an object, but to maintain an already dried object indefinitely in a dry condition.

Separating funnel

Separating funnel also known as separatory funnel or separation funnel. It is used to partition the components

of a mixture of immiscible liquids with different

densities.

Typically, one of the liquids will be water, and the other an organic solvent such as ether or chloroform.

The funnel, which is in the shape of a cone surmounted by a hemisphere, has a stopper at the top

Electronic balance

An electronic balance uses electromagnet to balance the weight on the pan.

This unit widely used in science laboratory as weighing equipment.

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Topic 2: Precipitation of barium (II) sulphate

________________________________________________________________________

Introduction:

Barium (II) nitrate reacts with sulphuric acid to produce barium (II) sulphate as white

precipitate which is difficult to dissolve in water. If sulphuric acid is in excess compared

to barium (II) nitrate, the entire ion Ba2+

will precipitate as barium (II) sulphate.

Therefore, barium (II) nitrate is a limiting reagent and the quantity of precipitate formed

is determined by this limiting reagent. The excess acid will then be separated using

distilled water.

Apparatus and Equipments:

Pipette (10mL) Hot water bath

Beaker (100mL) Filter funnel

Glass rod Filter paper

Watch glass Electronic balance

Materials: 0.1 M Barium (II) nitrate solution, Ba(NO3)2

0.5 M Sulphuric acid, H2SO4

Procedures:

1. Pipette 10mL of Ba(NO3)2 into a 100mL beaker. 2. Add 10mL of diluted H2SO4 and stir with a glass rod. Allow the precipitate to settle. 3. Weigh a clean and dry watch glass with a filter paper. Then, put the filter paper into

the filter funnel.

4. Carefully filter all the liquid in the beaker. 5. Add 20mL of distilled water into the same beaker, stir for 1 minute, and rinse the

glass rod and the side of the beaker. Wait for a while and filter like the previous step.

6. Decant the filtrate and remove the filter paper with precipitate to the watch glass. 7. Dry the precipitate in the oven until there is no liquid left on the watch glass.

When the precipitate dries, cool and measure the weight. Repeat this step until the

mass is same or less than 0.02g different.

Treatment of Data:

Mass of dry watch glass with filter paper g

Mass of watch glass with filter paper + precipitate g

Mass of BaSO4 precipitate g

Calculation:

1. Calculate the molecular weights of Ba(NO3)2 and BaSO4.

2. Calculate the moles of BaSO4 are produced in this experiment.

3. Calculate the moles of Ba(NO3)2 are required to react with H2SO4.

4. Determine which is limiting reactant in this experiment.

5. Determine the percentage yield of BaSO4.

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Topic 3: Determination of the amount of dissolved oxygen in a water sample by

iodometry the Winklers method ________________________________________________________________________

Introduction:

Myriad forms of life exist in lakes, streams and oceans. These creatures depend on

dissolved oxygen (D.O.) in water for their life support. Occasionally something will

happen which depletes the oxygen content of a natural water system.

The dissolved oxygen content is an important index when considering the suitability of

water for town supply. Good potable water will give a D.O. value close to the theoretical

value for a saturated solution for oxygen in water. When there is pollution from organic

matter and other trade effluents, the D.O. is used up in various biochemical oxidation

processes and it is only slowly replaced through surface adsorption. Such water will give

a low D.O. content until oxidation is completed. Adequate D.O. is necessary for the life

of fish and other aquatic organisms.

Gases which are dissolved in water obey Henrys law to a first approximation:

Xi = k . Pi

Where Xi is the mole fraction of the gas in solution, k is the Henrys law constant and Pi is the partial pressure of the gas. For oxygen, at 20

oC, k = 2.5 x 10

-5 atm

-1 and P (O2) =

0.21 atm for air.

Also, it is approximately true that X (O2) = n (O2) / n (H2O) since n (O2)

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4Mn(OH)2 (s) + O2 + 2H2O 4Mn(OH)3 (s)

When acidified, the manganese (III) oxidizes iodide which is already present to iodine:

2Mn(OH)3 (s) + 2I-

+ 6H+ I2 + 6H2O + 2Mn

2+

The liberated iodine, which is equivalent to the D.O. content, may then be titrated with

standard sodium thiosulphate:

I2 + 2S2O32- S4O6

2- + 2I

-

Apparatus:

Volumetric pipette (100 mL) Reagent bottle (1 L)

Conical flask (500 mL) Conical flask stopper

Burette Retort stand

Burette clamp Measuring cylinder (5 mL)

Pasteur pipette

Materials:

Manganese sulfate solution

Alkaline-iodide solution

0.025 M Sodium thiosulphate solution

Concentrated sulphuric acid

Starch solution

Procedures:

1. When sampling water, care must be taken to ensure that a good representative sample

of the water to be analyzed is obtained. For most purposes, this includes attention to

dissolved gases. Therefore, the water sample should be taken in a clean bottle which

must be filled to overflowing and tightly sealed with stopper without introduction of

air. If the water is sampled from a tap, it must be allowed to run for at least 5 minutes

prior to sampling. For this purpose, you may collect the water into a 1 L reagent

bottle fitted with a stopper.

2. Carefully remove the stopper from the conical flask and add 4 mL of the manganese

sulfate solution, discharging the reagent from the tip of a pipette put well below the

water surface. Replace the stopper.

3. Similarly, introduce 4 mL of the alkaline-iodide solution.

4. Place the stopper in the bottle; be sure that no air becomes entrapped. Some overflow

may occur. Thoroughly mix the content by inversion and rotation. Manganese

hydroxide is precipitated and will settle on standing.

5. When the precipitate has settled, introduce 4 mL of concentrated sulfuric acid with

the tip of the pipette well below the surface of the solution.

6. Replace the stopper and mix until the precipitate dissolves completely. The dissolved

oxygen now liberates free iodine from the potassium iodide present.

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7. Pipette 200 mL of the acidified sample into a 500 mL conical flask and titrate with

0.025 M sodium thiosulphate until the solution becomes pale yellow.

8. Add a few drops of starch indicator and continue the titration to the disappearance of

blue color.

(If the blue color doesnt appear after adding starch, repeat the titration and adding starch before start of titration)

9. Repeat the titration twice.

Treatment of data:

Titration number 1 2 3

Final reading

Initial reading

Volume used (cm3)

Average volume of titrant required for titration = ____________ cm3

Calculation:

1. Determine the concentration (ppm) of dissolved oxygen (DO) in the water sample.

2. Chemical equation below shows the reaction between iodine and thiosulphate.

I2 + 2S2O3-2

S4O6-2

+ 2I-

Calculate the oxidation states for the underlined elements and determine the oxidizing

and reducing agents.

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Topic 4: Gases Boyles Law and Charless Law ________________________________________________________________________

Introduction:

Boyle's Law states that the volume of a given mass of gas sample is inversely

proportional to the pressure applied to the gas if the temperature is kept constant.

P x V = constant

11 VP = 22 VP

Charless Law states that the volume of a given mass of gas sample is directly proportional to its absolute temperature, if the pressure is kept constant.

T

V = constant

2

2

1

1

T

V

T

V

The data collected from the experiment can be plotted on graphs in order to see the

relationship between volume and pressure and the relationship between volume and

temperature.


Syringe (25 mL) Pressure Gauge

Beaker (500 mL) Thermometer

Plastic tubing Lubricating oil

Hot water bath Retort stand

Burette Clamp Hotplate

Procedures:

Part 1 Boyles law

1. Pull the syringe plunger to 20 mL mark. 2. Connect the end of the syringe to the plastic tubing. 3. Connect the other end of the plastic tubing to the pressure gauge. (Make sure all of

the tubing are sealed securely)

4. Record the pressure when the syringe plunger is at 20 mL. 5. Reduce the volume in the steps of 2.5 mL until the 5 mL mark and record the pressure

for each.

6. Plot a graph of pressure (P) versus volume (V) and pressure (P) versus inverse of volume (1/V).

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Part 2 Charless law

1. Prepare about 400 mL of warm water in a beaker that is placed on top of a hotplate. 2. Lubricate the syringe plunger with oil to ensure smooth sliding in the syringe tube. 3. Pull the plunger to 5 mL mark and seal the end of the syringe. 4. At 50 C, immerse the syringe tube in the warm water by adjusting the height of the

burette clamp holding it. Wait about 1 minute for the air in the syringe to come to the

same temperature as the water.

5. Record the water temperature and the volume on the syringe. 6. Increase the water temperature to 55 C (wait about 1 minute for equilibrium to

occur) and record the water temperature and volume on the syringe.

7. Repeat step 6, by increasing the temperature in the steps of 5 C until 80 C. 8. Plot a graph of volume (V) versus temperature (C).

Treatment of data:

Part 1 Boyles law

Volume (mL) Pressure (kPa)

20.0

17.5

15.0

12.5

10.0

7.5

5.0

Calculation:

1. Determine the relationship of volume and pressure and find the proportionality

constant, k.

2. Based on your data, what would be the pressure to be if the volume of the syringe was

increased to 45 mL? Show the calculation steps.

Part 2 Charless law

Temperature(C) Volume (mL)

50

55

60

65

70

75

80

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Calculation:

1. Determine the relationship of volume and temperature and find the proportionality

constant, k.

2. Based on your data, postulate the x-intercept and y-intercept.

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Topic 5: Thermochemistry Determination of heat of neutralization of hydrochloric acid by sodium hydroxide

________________________________________________________________________

Introduction:

In adiabatic calorimetry, we try to keep as much as possible of the heat change that is

caused by a reaction of interest confined to the reaction vessel and its content, i.e., we

attempt to keep heat from leaking into and out of the calorimeter, usually by insulating

the reaction vessel. When there is no heat leakage to the surroundings, all the heat

evolved or absorbed by the reaction of interest must remain in the reaction vessel and its

contents, so it raises or lowers the temperature. It is this temperature change that we

measure and relate to the heat evolution or absorption that has occurred.

The relationship between the amount of heat, qs that enters or leaves a system and the

attendant temperature change, T; may be expressed as

qs

C = C = total heat capacity of the system

T

If the system is composed of several parts, e.g., of a reaction vessel plus the solution it

contains, then the total heat capacity of the system is the sum of the heat capacities of the

individual parts, C = Ci . Since, in the ideal case, no heat is allowed to enter or escape, the sum of the heat from the reaction (qr) and the heat from the system (qs) must be equal

to zero,

qr + qs = 0

so that,

qr + C T = 0

Heat capacity is a positive quantity. Therefore, if the temperature of the system rises, i.e.,

T is positive, then qr will be a negative quantity.

In practice, it is impossible to make a perfect calorimeter, one which has no heat leakage,

so measuring the temperature change due to the heat from the reaction is complicated by

temperature changes from heat leaking into or out of the calorimeter. By taking these

factors into account and correcting them, a graphical procedure is usually used to obtain

the temperature change T, as shown in the following figure of temperature versus time for an exothermic reaction.

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The initial temperature of the reaction vessel and contents was a little below room

temperature, so heat was leaking into the system. This explains the slow rise in

temperature before the reaction was initiated. After the initial sharp rise in temperature

when the reaction has begun, the temperature began to decrease due to heat loss from the

calorimeter. At first, these readings were a bit erratic, as temperature equilibration was

not yet attained throughout the system.

After some time, the equilibration (and the reaction has completed, the heat loss became

steady and the temperature decreased linearly. Extrapolation of this linear portion of the

curve back to the time when the reaction began, gives a fairly accurate value for the

temperature which would have been attained if the reaction and temperature equilibration

had taken place instantaneously as if that no heat loss has occurred.

Generally, a chemical change is usually accompanied by an evolution or absorption of

energy (e.g. heat); certain physical changes also produce an energy change. The heat of

neutralization of an acid by a base is the energy evolved when 1 mol of the acid is

neutralized by 1 mol of a base, the reaction being carried out in dilute aqueous solution.

For a strong acid and a strong base, the heat of neutralization is effectively constant.

This is because strong acids and strong bases, and the salts they form are all completely

ionized in dilute solution. Thus the reaction between any strong acid and strong base is

simply the formation of unionized water from H+

and OH- ions.

e.g.: HCl + KOH KCl + H2O

H+

+ Cl- + K

+ + OH

- K

+ + Cl

- + H2O

H+

+ OH-

H2O H = Heat of neutralization.

H

Decreasing slope due to loss of heat to the calorimeter walls

Starting time

Room temperature

T

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The constancy of the heat of neutralization of any strong acid and any strong base

provides simple, but convincing evidence that strong acids and bases are, in fact,

completely ionized.

The temperature rises accompanying the mixing of a solution of a base with a solution of

acid at the same initial temperature in a calorimeter is caused by the heat released during

the neutralization process.


Joule calorimeter

Digital stop-watch

Thermometer (reading to 0.1 oC)

Measuring cylinder (50 mL)

Hot water bath

Beaker (100 mL)

Materials:

0.5 M Sodium hydroxide, NaOH

0.5 M Hydrochloric acid, HCl

Distilled water

Procedures:

Part 1 Calibration of calorimeter

1. Measure 50 cm3

of distilled water into a calorimeter.

2. Record the temperature of the water (1 minute interval) for 5 minutes.

3. Add another 50 cm3

of water which has initial temperature of 50 oC into the

calorimeter and record the temperature (1 minute interval) for another 10 minutes.

Part 2 Determination of heat of neutralization

1. Measure 50 cm3 of 0.5 M NaOH solution into the calibrated calorimeter.

2. Record the temperature of the base every 2 minutes.

3. Measure 50 cm3 of 0.5 M HCl into another container.

4. Insert a thermometer and cool or warm slightly until the temperature of the acid is the

same as that of the base.

5. Now add the acid to the base quickly and note the time of mixing.

6. Record the temperature (30 seconds interval) for first 5 minutes and (1 minute

interval) for the next 10 minutes.

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Treatment of data:

Part 1 With 50 cm3 distilled water

Time (minute) Temperature (oC)

1.00

2.00

3.00

4.00

5.00

Addition of 50 cm3 distilled water

Time (minute) Temperature (oC)

1.00

2.00

3.00

4.00

5.00

6.00

7.00

8.00

9.00

10.00

(Specific heat capacity of water = 4.184 J g-1

K-1

; density of water = 1.00 g cm-3

)

We can imagine the process that occurs on mixing the hot water with the water already in

the calorimeter as being broken into two parts: the release of enough heat by the high

temperature water reduces its temperature from 50 o C to Tfinal.. Hence,

qr = Mass of hot water (50 oC) x Specific heat of water x Fall in temperature (Tfinal 50

oC)

and

qs = [Mass of distilled water x Specific heat of water x Rise in temperature (Tfinal Tinitial room temp)] + [Heat capacity of calorimeter x Rise in temperature (Tfinal Tinitial room temp)]

qr + qs = 0

Plot a graph of temperature against time.

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Part 2 With 50 cm3 of 0.5 M NaOH

Time (minutes) Temperature (oC)

2.00

4.00

6.00

8.00

10.00

12.00

Addition of 50 cm3

of 0.5 M HCl

Time (seconds) Temperature (oC)

30

60

90

120

150

180

210

240

270

300

360

420

480

540

600

660

720

780

840

900

(Specific heat capacity of water = 4.184 J g-1

K-1

; density of solution = 1.00 g cm-3

)

Assuming the specific heat of the solutions to be the same as that of water, the heat

evolved by the neutralization process is given by

Hneutralization = MNaOH CNaOH T + CCalorimeter T + MHCl CHCl T

Plot a graph of temperature against time.

Calculation:

1. Calculate the heat of reaction (qr) for Part 1.

2. Determine the heat capacity of the calorimeter.

3. Determine the enthalpy of neutralization (Hneutralization) for Part 2. 4. Determine the molar enthalpy of neutralization.

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Topic 6: Reaction kinetics Determination of the activation energy of the reaction between oxalic acid and potassium permanganate

________________________________________________________________________

Introduction:

Reaction between oppositely charged icons are often very fast but the reaction between

similarly charged ions, e.g. permanganate and oxalate, may proceed at a rate which is

measurable:

2KMnO4 + 5H2C2O4 + 3H2SO4 K2SO4 + 2MnSO4 + 8H2O + 10CO2

or

2MnO4- + 16H +

+ 5C2O4 2-

K2SO4 + 2MnSO4 + 8H2O + 10CO2

The rate of this reaction is measured by the time taken for the disappearance of the purple

color of the potassium permanganate.


Measuring cylinder, test tube, water bath, stopwatch

Materials:

0.02M Potassium Permanganate

1 M Sulphuric Acid

0.5 M Oxalic Acid

Test tube

Water bath

Measuring cylinder (5mL)

Stopwatch

Procedure:

1. In a test tube, measure 2 cm3 of 0.02 M potassium permanganate and 4 cm

3 of 1 M

sulphuric acid.

2. In another test tube, place 2 cm3

of oxalic acid.

3. Place the test tubes in water bath at 35 oC.

4. When the solutions have attained this temperature, pour the oxalic acid into the

acidified permanganate solution and immediately start the time.

5. Record the time taken for the permanganate to decolorize.

6. Repeat the experiment at higher temperatures of 40, 45, 50, 55 and 60 oC.

Treatment of data:

Temperature, T (C) 35 40 45 50 55 60

Temperature, T (K)

1 / T (K-1

)

Reaction time, t (sec)

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ln 1/t

Plot a graph of ln 1/t against

1/T.

Calculation: 1. Calculate the activation energy, Ea in J mol

-1 according to the Arhenius equation

k = A e -Ea / RT

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Topic 7: Rate law of an iodine clock reaction

________________________________________________________________________

Introduction:

In this experiment you will study the reaction of hydrogen peroxide with iodide ion in the

presence of acid.

H2O2 (aq) + 2 I- (aq) + 2 H3O

+ (aq) I2 (aq) + 4 H2O (l)

The rate of the reaction is very much depending on the concentrations. The easiest way to

measure the rate of reaction is by determining the rate of formation of iodine by reaction

between iodine (I2) and thiosulfate ion (S2O32-

) to form the tetrathionate ion (S4O62-

)

according to:

I2 (aq) + 2 S2O32-

(aq) S4O62-

(aq) + 2 I- (aq)

The rate of a chemical reaction also depends on temperature, pressure and other physical

characteristics of the reaction surroundings. However, the first consideration a chemist

gives to a chemical reaction is the concentration of the reactants. Thus, chemists often

change the concentration of reactants so that they can study the effect of such changes on

the rate of the reaction.


Graduated pipette(1 mL and 25 mL) 250 mL conical flask

Digital stop watch Hot water bath

Measuring cylinder Thermometer

Materials: 0.36 M Sulphuric acid, H2SO4

0.025 M Potassium iodide, KI

0.8 M Hydrogen peroxide, H2O2

0.0025 M Thiosulfate, S2O32-

Starch solution

Distilled water

Procedures:

Part 1 Measurement of Rate For One Set Conditions 1. To a clean conical flask, add the following in order.

43 cm3 distilled water

35 cm3 0.36 M H2SO4

10 cm3 0.025 M KI

10 cm3 0.0025 M thiosulfate

1 cm3 starch solution

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2. Pipette 1 cm3 of 0.8 M hydrogen peroxide in the fume cupboard. With one eye on a

watch, add the peroxide and noting the time to the nearest second.

3. Note the time at which the thiosulfate is all used up and blue colour appears. Record

the time elapsed.

4. Measure the temperature of the reaction mixture.

5. Repeat the experiment.

Part 2 The Effect of Peroxide 1. Proceed as in Part 1 but

(i) use 2 cm3 peroxide and 42 cm

3 distilled water.

(ii) use 0.5 cm3 of peroxide and 43.5 cm

3 distilled water.

2. Plot a graph of rate against concentration of peroxide in the 3 mixtures and describe

it. Use axes starting from zero for rate and for [H2O2].

Part 3 The Effect of Iodide 1. Proceed as in Part 1 but

(i) use 20 cm3 KI and 33 cm

3 distilled water

(ii) use 5 cm3 KI and 48 cm

3 distilled water.

2. Plot a graph of rate against concentration of iodide in the 3 mixtures and describe it.

Use axes starting from zero for rate and for [I-].

Part 4 The Effect of Acid 1. Proceed as in Part 1 but

(i) use 8 cm3 distilled water and 70 cm

3 sulphuric acid.

(ii) use 60.5 cm3 distilled water and 17.5 cm

3 sulphuric acid.

2. Plot a graph of rate against [H3O+] in the 3 mixtures and describe it. Use axes starting

from zero for rate and for [H3O+].

Treatment of data:

Part 1 Measurement of rate for one set conditions:

Number of trial 1 2

Concentration of peroxide (M)

Concentration of iodide (M)

Concentration of sulphuric acid (M)

Concentration of thiosulphate (M)

Time elapsed (s)

Temperature of reaction mixture (oC)

Average time elapsed = _________________________ s

Average rate of reaction of peroxide = _______________________ M s-1

Average rate of reaction of iodide = _________________________ M s-1

Average rate of reaction of sulphuric acid = ___________________ M s-1

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Part 2 The effect of peroxide:

(i) 2 cm3 peroxide and 42 cm

3 distilled water

Number of trial 1 2


Time elapsed (s)


Average time elapsed = ______________________ s

Average rate of reaction = ____________________ M s-1

(ii) 0.5 cm3 of peroxide and 43.5 cm

3 distilled water

Number of trial 1 2


Time elapsed (s)


Average time elapsed = _____________________ s

Average rate of reaction = ___________________ M s-1

Part 3 The effect of iodide:

(i) 20 cm3 KI and 33 cm

3 distilled water

Number of trial 1 2

Concentration of KI (M)

Time elapsed (s)


Average time elapsed = _______________________ s

Average rate of reaction = _____________________ M s-1

(ii) 5 cm3 KI and 48 cm

3 distilled water

Number of trial 1 2

Concentration of KI (M)

Time elapsed (s)




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Part 4 The effect of acid:

(i) 8 cm3 distilled water and 70 cm

3 sulphuric acid

Number of trial 1 2

Concentration of sulphuric acid(M)

Time elapsed (s)




(ii) 60.5 cm3 distilled water and 17.5 cm

3 sulphuric acid.

Number of trial 1 2

Concentration of sulphuric acid(M)

Time elapsed (s)


Average time elapsed = ______________________ s

Average rate of reaction = ____________________ M s-1

Calculation:

1. Using Part 2 (i) as an example, show the steps in calculating

(a) concentration of peroxide

(b) average rate of reaction

2. For Part 1,

(a) calculate the number of moles of thiosulfate added to the reaction mixture for use

in reaction

(b) calculate the number of moles of iodine produced by reaction

(c) calculate the rate of reaction (concentration of iodine produced by reaction

divided by the time taken for the blue colour to appear)

3. Based on the results obtained, determine the rate law for this experiment.

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Topic 8: Determination of pKa of Bromocresol Green

Introduction:

Bromocresol green is a weak acid which dissociates in water according to the equation:

HB- (aq) + H2O B

2- (aq) + H3O

+ (aq)

yellow blue

The yellow, acidic form of the indicator is a singly charged anion which can be

represented by the structure:

C

SO3-

O

Br

Br

CH3

CH3

Br

OH

Br

Na+

The hydrogen of the phenol group is weakly acidic and dissociates to give the basic form,

a blue coloured, doubly charged anion.

At equilibrium, the pKa of the indicator is related to the pH of the solution by the

equation:

2

10log

B

HBa

C

CpHpK

Ka= the equilibrium constant for the dissociation of the acid HB-

CHB-= concentration of the undissociated form of indicator

CB2-

= concentration of the dissociated form of indicator.

The bromocresol green in the acidic solution is completely in its protonated form, HB-. In

the basic solution, it is completely in the unprotonated form, B2-

.


Erlenmeyer flask (100 mL) Stopper

Graduated pipette (2 mL) Pipette (10 mL)

Volumetric flask (100 mL) pH meter

Measuring cylinder (10 mL) UV-Visible Spectrophotometer

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Materials:

Bromocresol green indicator solution

2 M Hydrochloric acid, HCl

2 M Sodium hydroxide, NaOH

Buffer solution (pH 5)

Distilled water

Procedures:

1. Bromocresol green in basic solution: Put 10 cm3 of 2 M NaOH in a 100 cm3 volumetric flask. Then, pipette 2 cm

3 of bromocresol green indicator solution into the

flask. Fill to the mark with distilled water, stopper the flask and mix the contents

thoroughly.

2. Bromocresol green in acidic solution: Put 10 cm3 of 2 M HCl in a 100 cm3 volumetric flask. Then, pipette 2 cm

3 of bromocresol green indicator solution into the

flask. Fill to the mark with distilled water, stopper the flask and mix the contents

thoroughly.

3. Bromocresol green in buffer solution: Put 10 cm3 of pH 5 buffer solution in a 100 cm

3 volumetric flask. Then, pipette 2 cm

3 of bromocresol green indicator solution

into the flask. Fill to the mark with distilled water, stopper the flask and mix the

contents thoroughly. Measure the pH of this solution by using pH meter.

4. Use distilled water as the blank. Measure the spectrum of these solutions from wavelength 400 630 nm.

Treatment of data:

Bromocresol green in pH max Absorbance

acidic solution

basic solution

buffer solution

Cuvette length = ______________ cm

Wavelength (nm) Absorbance of bromocresol green in

acidic solution buffer solution basic solution

400

410

420

430

440

450

460

470

480

490

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500

510

520

530

540

550

560

570

580

590

600

610

620

630

Plot a graph of absorbance against wavelength for each of solution with different pH.

Calculation:

1. Determine the isosbestic point for bromocresol green.

2. Calculate the pKa value for indicator using the pH of the indicator/buffer solution.

3. Determine the maximum buffer capacity for the bromocresol green.

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Topic 9: Determination of the percentages of sodium carbonate and sodium

hydroxide in a mixture

________________________________________________________________________

Introduction:

The total alkali (sodium carbonate and sodium hydroxide) present in a mixture is

determined by titration with hydrochloric acid using methyl orange as indicator.

Na2CO3 + 2HCl 2NaCl + H2O + CO2

NaOH + HCl NaCl + H2O

To a second portion of the mixture, excess barium chloride solution is added.

Na2CO3 + BaCl2 2NaCl + BaCO3

Thus sodium carbonate is removed from the mixture as insoluble barium carbonate. The

hydroxide ion remaining in solution reacts with the acid during titration. The

concentration of these is determined using phenolphthalein as indicator.

Apparatus:

Burette (50mL) Volumetric pipette (25mL)

Pipette filler Erlenmeyer flasks

Retort stand Burette clamp

Asbestos board

Materials:

Hydrochloric acid, HCl

Mixture of sodium carbonate, Na2CO3 and sodium hydroxide, NaOH in a solution

Barium chloride, BaCl2

Methyl orange indicator

Phenolphthalein indicator

Procedures:

1. Pipette 25 cm3

of the mixture solution into a conical flask.

2. Add a few drops of methyl orange indicator, and titrate with diluted hydrochloric acid

until the yellow color just changes to orange.


4. To another 25 cm3

of the mixture solution, add about an equal volume of barium

chloride solution and one or two drops of phenolphthalein.

5. Titrate the mixture with diluted hydrochloric acid, noting the burette reading when

the solution is decolorized (Take care to run in the acid slowly otherwise some of the

barium carbonate may be acted upon before the end-point is reached).


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Treatment of data:

(a) Methyl orange as indicator

Titration number 1

2 3

Final reading

Initial reading

Volume Used (cm3)

Average volume of hydrochloric acid required for titration = __________ cm3

(b) Phenolphthalein as indicator

Titration number 1

2 3

Final reading

Initial reading

Volume Used (cm3)

Average volume of hydrochloric acid required for titration = __________ cm3

Calculation:

1. Calculate the respective concentrations of sodium carbonate and sodium hydroxide in

the mixture in the units of g L-1

.

2. Determine the percentage of each alkali in the mixture.

(The molecular weights of sodium hydroxide and sodium carbonate are 40 and 106

respectively.)

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Topic 10: Ionic Measurement of the conductivities of strong and weak electrolytes ________________________________________________________________________

Introduction:

Solutions of electrolytes conduct electric current by migration of ions under the influence

of electric field. According to Ohms law,

V = I R

Where V, is the potential difference; I, is the current and R, is the resistance. The term

conductance (L) is generally used for dealing with electrolytes and this is defined as the

reciprocal of the resistance of the solutions, i.e.,

L (ohm -1 or ) = 1 / R

The conductivity (x) may be obtained from

x (-1 cm -1) = d / AR

Where A, and d are the area and separation between the electrodes of the conductivity

cell. The cell constant k of the conductivity cell is defines as

k = d / A

And therefore

x = k. L

Thus, conductivity is the reciprocal of the resistance in ohm of a 1 cm of liquid at a

specified temperature. The molar conductivity of an electrolyte solution is defined as

(-1 mol -1 cm 2) = x / C

Where C = molar concentration

For week electrolytes, the increase of molar conductivity with increasing dilution is

ascribed to increased dissociation of the electrolyte molecules to free ions. For strong

electrolytes, the molar conductivities are higher than those of weak electrolytes at high

concentrations. As the solutions become dilute, the molar conductivities also increase as

in the case of weak electrolytes but the variation is, in general, less steep than for weak

electrolytes.

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Conductivity meter Beaker (100 mL)

Volumetric flask (100 mL) Pasteur pipette

Volumetric pipette (50 mL) Wash bottle

Pipette filler

Materials:

0.1 M Sodium chloride, NaCl

0.1 M Acetic acid, CH3COOH

Distilled Water

Procedures:

1. Calibrate the conductivity meter with known cell constant. 2. Measure the conductivity of 50 cm 3 of 0.1 M acetic acid. 3. Repeat the measurement with 0.0500, 0.0250, 0.0125, and 0.00625M acetic acid

prepared by successive dilution.

4. Repeat with sodium chloride. 5. Measure the conductivity of the distilled water used.

Treatment of data:

Conductivity of distilled water = _________________

Conductivity of electrolytes = Conductivity of solution - Conductivity of distilled water

Tabulate the conductivities as follows:

C (Mol dm-3

) Conductivity (-1 cm-1 )

CH3COOH NaCl

0.1000

0.0500

0.0250

0.0125

0.00625

Calculate the molar conductivities and tabulate the results as follows:

C1/2

(Mol dm-3

) Molar Conductivity (

-1 mol

-1 cm

2)

CH3COOH NaCl

Plot a graph of versus C1/2 for each electrolyte.

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Calculation:

1. Show the steps (one example of calculation) in obtaining molar conductivity for each

electrolyte.

2. Determine the conductivity and molar conductivity at infinite dilution (o) for each electrolyte.

(a) CH3COOH

(b) NaCl

3. The molar conductivity of benzoic acid at infinite dilution is 385 -1 cm2 mol-1 and at

a concentration of 0.0050 mol dm-3

is 40.8 -1 cm2 mol-1. Calculate the pKa of benzoic acid.

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Physical chemistry experiments