calibration of volumetric glassware

GENERAL
CHEMISTRY
Laboratory Guide
This module provides a laboratory guidelines, safety declaration form, Lab Report guidelines and Laboratory manual for subject of General Chemistry (CLD 10004).
Mohd Zulkhairi Abdul Rahim

Laboratory Information
Before each lab session, you should prepare by reading the lab manual and the textbook required reading. We expect you to have a good understanding of the purpose, details of the procedure, the use of all chemicals and any significant hazards, and the underlying science of the experiment when you come to lab.

TABLE OF CONTENTS

Page

TITLE PAGE
INTRODUCTION
LABORATORY SAFETY GUIDELINES
SAFETY DECLARATION FORM
CHEMISTRY LABORATORY REPORT QUIDELINES

EXPERIMENT

3

1

Calibration of Volumetric Glassware

3

2

Empirical Formulas

3

3

Standardization of Sodium Hydroxide Solutions

10

4

Acids, Bases, Buffers and The Determination of pH

24

5

Redox Reactions

29

6

Qualitative Analysis of Common Ions

REFERENCES

103

APPENDICES

113

INTRODUCTION

This manual provides laboratory guidelines, safety declaration form, Chemistry Lab
Report guidelines and Laboratory manual for subject of Chemistry (CLB 10203).
The primary purpose of this manual is to compile all necessary information regrading laboratory component in one manual.
The manual contains four parts. Part 1 provides a description of the general requirements of a thesis produced in UniKL MICET, which should be useful to the student. Part 2 containing requirements for format and arrangement of research proposal which the student intends to undertake in the completed thesis. Part 3 explains the different parts of the thesis manuscript and how it is organized. Part 4 is on formatting of the thesis which the student will find necessary at the end of the thesis writing prior to submission.
Upon successful completion of this course, students will understand safety, transfer and measurement of chemicals, using physical properties to identify compounds, chemical reactions, paper chromatography, and pH. Students will gain experience in a variety of laboratory techniques to safely conduct chemical experiments and procedures.

Students will be able to independently perform accurate quantitative measurements, interpret experimental results, perform calculations on these results and draw a reasonable, accurate conclusion.
Students will synthesize, isolate, purify and characterize a series of compounds using modern methods. Students will demonstrate knowledge of proper use of modern instrumental techniques. Students will be able to design an experimental procedure.
Students will observe safe practices in the laboratory and will know how to respond in an emergency. Students will learn to gather hazardous materials information and will recognize and respond properly to potential hazards of handling chemicals and chemical waste. There may be shortcomings which we had overlooked but we pray that these should not hinder the process of producing a quality thesis. We welcome all suggestions and criticism, which could be later, included in future editions.

LABORATORY SAFETY GUIDELINES
General Guidelines

1. Conduct yourself in a responsible manner at all times in the laboratory.
2. Be familiar with your lab assignment before you come to the lab. Follow all written and verbal instructions carefully. If you do not understand a direction or part of a procedure, ask the instructor before proceeding.
3. No student may work in laboratory alone. The lab instructor or co-coordinator grant exceptions on a case by case basis.
4. When first entering a laboratory, do not touch any equipment, chemicals or other materials in the laboratory area until you are instructed to do so.
5. Do not eat, drink beverages or chew gum in the laboratory. Do not use laboratory glassware as containers for food or beverages.
6. Smoking is not allowed in any indoor area.
7. No music allowed in the laboratory. Radio (including walkman) and other entertainment devices are not permitted.
8. No cellular phone is allowed in this laboratory.
9. Perform only those experiments authorized by the instructor. Never do anything in the laboratory that is not called for the laboratory procedures or by your instructor. Carefully follow all instructions, both written and oral. Unauthorized experiments are prohibited.
10. Observe good housekeeping practices. Work areas should be kept clean and tidy at all times.
11. Horseplay, practical jokes, and pranks are dangerous and prohibited.
12. Always work in a well-ventilated area.
13. Bring only your laboratory instructions, worksheets and report to the work area. Other materials (books, purses, backpacks, etc) should be stored in the cabinet.
14. Know the locations and operation procedures of all safety equipment including the first aid kit, eyewash station, safety shower, spill kit and fire extinguisher.
15. Be alert and proceed with caution at all times in the laboratory. Notify the instructor immediately of any unsafe condition you observe.
16. Label and equipment instructions must be read carefully before use. Set up and use the

prescribed apparatus as directed in the laboratory instructions provided by your instructor.
17. Experiments must be personally monitored at all times. You will be assigned a laboratory station at which to work. Do not wander around the room, distract other students or interfere with laboratory experiments or others.
18. Write your name and equipment use every time you come in to the laboratory in the log book.
19. Defeating safety devices or using equipment in a manner other than that which is intended will be grounds for dismissal from the lab.

Clothing

1. Safety goggles and safety jacket must be worn whenever you work in lab.
2. Gloves should be worn whenever you use chemicals that cause skin irritations or need to handle hot equipment.
3. Mask should be worn every time you prepare the chemicals.
4. Safety shoes and hard hat should be worn at all times while in the laboratory.
5. Contact lenses should not be worn in the laboratory unless you have permission from your instructor. 6. Dress properly during a laboratory activity.
7. Long hair, dangling jewelry and loose or baggy clothing are a hazard in the laboratory. Long hair must be tied back and dangling jewelry and loose or baggy clothing must be secured.
8. Sandal, open-toed shoes, high heels or shoes with holes in the sols will not be worn in the lab.
9. Short and skirts are not permitted.
10. Instructor and laboratory assistant have a right dismiss to you from the laboratory if they found that you are not wearing proper safety clothing.

Handling Chemicals

1. Treat chemicals with respect and understand the chemicals you are using with Material Safety
Data Sheet (MSDS). The MSDS are available in the analytical room.

2. All chemicals in the laboratory are to be considered dangerous. Do not touch, taste or smell any chemical unless specifically instructed to do so.
3. Check the label on chemical bottles before removing any of the contents. Take only much chemical are you need. Smaller amounts often work better than larger amounts.
4. Label all containers and massing papers holding dry chemicals.
5. Never return unused chemicals to their original containers.
6. Never use mouth suction to fill a pipette. Use pipette bulb or pipette filler.
7. Acids must be handled with extreme care. Always add acids slowly to water, with slow stirring and swirling, being careful of the heat produced, particularly with sulfuric acid.
8. Handle flammable hazardous liquid over a pan to contain spills. Never dispense flammable liquids anywhere near a flame or source of heat.
9. Never take chemicals or other materials from the laboratory area.
10. Take good care when transferring acids and other chemicals from one part of the laboratory to another. Hold them securely and in the method demonstrated by the instructor as you walk.
11. All wastes generated during the course of an experiment must be disposed of according to the lab instructor’s directions.
12. Never mix chemicals in sink drains.
13. Sinks are to be used only for water and those solutions designated by the instructor.
14. Solid chemicals, metals, matches, filter paper, and all other insoluble materials are to be disposed of in the proper waste containers, not in the sink.
15. Checks the label of all waste containers twice before adding your chemicals waste to the container. 16. Cracked or broken glass should be placed in the special container for “broken glass”.
17. Keep hands away from your face, eyes, mouth and body while using chemicals. Wash your hands with soap and water after performing all experiments.

Personal Hygiene

1. Wash hands before leaving the lab and before eating.
Gloves should be removed before leaving the lab, using telephones, or entering common areas.

Accidents and Injuries

1. Report any accidents (spill, breakage, etc) or injury (cut, burn, etc) to the instructor immediately, no matter how trivial it may appear.
2. If you or your lab partners are hurt, immediately tell to the instructor.
3. If a chemical should splash in your eye(s), immediately flush with running water from the eyewash station for at least 20 minutes. Notify the instructor immediately.
4. Spills should be cleaned up immediately.

Handling Glassware and Equipment

1. Inserting and removing glass tubing from rubber stopper can be dangerous. Always lubricate glassware (tubing, thistle tubes, thermometer, etc) before attempting to insert it in a stopper.
Always protect your hands with tower or cotton gloves when inserting glass tubing into, or removing it from a rubber stopper.
2. When removing an electrical plug from its socket, grasp the plug, not the electrical cord.
3. Hands must be completely dry before touching an electrical switch, plug or outlet.
4. Examine glassware before each use. Never use chipped or cracked glassware.
5. Never use dirty glassware.
6. Do not immerse hot glassware in cold water; it may shatter.
7. Report damaged electrical equipment immediately. Look for things such as frayed cords, exposed wires and loose connections. Do not use damaged electrical equipment.
8. If you do not understand how to use a piece of equipment, ask the instructor for help.
9. Be careful when lifting heavy objects. Lift comfortably, avoid unnecessary bending, twisting, reaching out, and excessive weights, lift gradually and keep in good physical shape.
10. Do not transfer a glassware form one laboratory to another without permission from instructor.

Heating Substances

1. Do not operate a hot plate by yourself. Take care that hair, clothing, and hands are a safe distance from the hot plate at all times. Use of hot plate is only allowed in the presence of the teacher. 2. Heated glassware remains very hot for a long time. They should be set aside in a designated place to cool, and picked up with caution. Use tongs or heat protective gloves if necessary.
3. Never look into a container that is being heated.
4. Do not place hot apparatus directly on the laboratory desk. Always use an insulated pad. Allow plenty of time for hot apparatus to cool before touching it.
5. If leaving a lab unattended, turn off all ignition sources and lock the doors.

Ended the Experiments

1. At the end of the laboratory sessions, you should;


Shut-off main gas outlet



Turn-off the water inlet



Desk top, floor area and sink are clean



All equipment is cool, clean and arranged

2. All equipment use should be flushed using deionized water.

SAFETY DECLARATION FORM

The Dean/Head of Campus
Universiti Kuala Lumpur
Malaysian Institute of Chemical and Bioengineering Technology
Lot 1988, Vendor City Industrial Area
Taboh Naning, 78000 Alor Gajah
Malacca
Dear Sir,
SAFETY DECLARATION
I ………………………………..…………………………………………………………. ID No
………………………. declare that I have read and understood the safety rules and regulations in UniKL MICET. I hereby agree to abide by all the rules and regulations stated in the safety guidelines. 2. I hereby understood the contents and will disciplinary action will be taken against me, if I do not abide by the stated rules.

3. I am fully responsible for all my actions during laboratory sessions.

Thank you.
Yours faithfully,

……………………………….
Name:
Matrix No:
Subject:
Date:

CHEMISTRY LAB REPORT FORMAT

You should type your lab report, but you may draw or write by hand any tables, diagrams, or chemical equations as long as you do it neatly. Make sure that you check your document for any spelling errors. Each lab report is worth
100
points.
You should also read the student handbook on the subject of plagiarism. Your data and observations will be similar, but your interpretations should not be written identically. You may not copy another student's lab report in part or in its entirety. If you are found guilty of this infraction, you and the person from whom you copied will both lose points. In extreme cases or repeated offenses, both students may receive a zero for the lab.
Title
Use a separate title page. Include the title of the experiment, YOUR NAME, and the date.
Also clearly indicate the name(s) of your lab partner(s).
Purpose
Write a few sentences describing what you are supposed to learn by doing the experiment. You might write about learning the lab procedures themselves. Do not just copy word for word from the lab handout.
Materials
List the chemicals and equipment needed to perform the experiment.
Procedure
Briefly list or summarize the procedure. Again, DO NOT COPY DIRECTLY from the lab handout. Use your own words.
Data
Include any data, qualitative and quantitative, that you collect. This includes any observations.
You should include the proper units with any numbers, as well as use the proper number of significant figures based upon the lab equipment used. Remember, read the last known place and estimate one more digit. DO NOT place any calculations or data analysis in this section. It may be a good idea to reproduce here any data tables that you completed during the lab.
Data Analysis
Here is where any calculations or graphs are placed. Also show how you arrive at identifying any unknowns. Make sure that your graphs have titles, labeled axes with units, and legends.
Conclusions

This section is the most important one. Include the following:
· Percent error
· Sources of error (Don't just say that any errors were the fault of the equipment. Also don't use the generic excuse that you might have misread the measurements.)
· Identify any unknowns
· Answer the questions, "What did you learn?"
"Did I accomplish the purpose?"
"How would I improve the experiment next time?"
References
Write down any sources such encyclopedia, books, etc. that you used.

as

your

textbook,

the

Internet,

electronic

Appendix
Here is where you attach any material that you think is pertinent to the lab report. Also answer any questions here that are in the lab report. You do not have to re-write the questions, but label and number them appropriately.

EXPERIMENT 1
CALIBRATION OF VOLUMETRIC GLASSWARE

OBJECTIVES
• To calibrate a 10 mL volumetric pipette.
• To calibrate a 25 mL volumetric pipette.
• To calibrate a 100 mL volumetric flask.
• To calibrate a 50 mL measuring cylinder.

INTRODUCTION
Volumetric glassware is marked with TD for ‘to deliver’ and TC for ‘to contain’ and also with the temperature at which the calibration applies. For greatest accuracy, volumetric glassware should be calibrated to measure the volume that is actually contained in or delivered by a particular piece of glassware. The calibration is done by measuring the mass of water contained in or delivered by the glassware. The density of water at a particular temperature is used to convert mass into volume. Pipettes and burettes are calibrated to deliver specific volumes; whereas, volumetric flasks are calibrated on a to contain basis.

EXPERIMENTAL PROCEDURE
You must use three complete trials for each of the calibration.
1) Calibration of a volumetric pipette ( 10 ml and 25 ml) .
a) Obtain a transfer pipette. If distilled water does not drain uniformly but leaves droplets of water in the inner surface, the pipette should be cleaned. Use a cleaning solution or detergent to clean the pipette (ask your instructor). Please make sure you have read
(Laboratory techniques) or have been shown the correct way of pipetting by your instructor. b) The water used for calibration should be in thermal equilibrium i.e same temperature with the surroundings. The temperature should be recorded at uniform intervals. Weighing to the nearest milligram will be satisfactory and the use of a top loading balance is sufficient. Weighing bottles or small conical flasks with a stopper can be used as containers to hold the calibration liquid.
i. Weigh an empty weighing bottle to the nearest milligram.

ii. Fill the pipette to the mark with distilled water. iii. Drain the water by gravity (remove pipette bulb or pump) into the weighing bottle and cap the bottle to prevent evaporation. iv. Weigh the bottle again to find the mass of water delivered from the pipette.
v. Use the following equation to convert mass to volume:

True (actual) Volume = (grams of water) x (volume of 1 g of H2O in Table 4.1)

Notes:



Repeat the procedure another 2 times
Perform the above procedure on each of the pipette.

2) Calibration of a volumetric flask (100 ml).
a) Get a 100 mL volumetric flask. Make sure the flask is clean and dry. (Dry the flask by clamping the flask in an inverted position at room temperature if necessary. Do not dry a volumetric flask in an oven. Very seldom volumetric flasks need to be dried).
Weigh the flask to the nearest milligram.
b) Fill the flask to the mark with distilled water and weigh again.
c) Calculate the mass of water contained in the flask.
d) Convert the mass of water to volume
3) Calibration of a measuring cylinder (50 ml)
Calculate the mass of water contained in the cylinder and convert it to volume by the procedure you think most appropriate.

Table 1: Volume occupied by 1.000g of water weighed in air
Temperature,°C

Volume, mL at ToC

20

1.0028

21

1.0030

22

1.0033

23

1.0035

24

1.0037

25

1.0040

26

1.0043

27

1.0045

28

1.0048

29

1.0051

30

1.0053

31

1.0056

32

1.0058

33

1.0061

Standard Deviation ( σ )
Sample Standard Deviation

The standard deviation value represents the average distance of a set of scores from the mean.
It is a statistical measure of the precision for a series of repetitive measurements. The sample standard deviation, s, is the positive square root of the sample variance.

Where

is the sample,

is the mean of the sample and N is sample size

RESULTS
A) DATA:
i) 10 mL pipet
Trial 1

:

Temperature (˚C)

:

Actual volume (mL)

:

Average volume (mL)

:

Standard deviation, σ

:

Relative standard deviation (σ / x )

Trial 3

:

Mass of water (g)

Trial 2

:

Mass of container (g)

Trial 3

Trial 1

Mass of container + water (g)

Trial 2

:

ii) 25 mL pipet
Mass of container + water (g)

:

Mass of container (g)

:

Mass of water (g)

:

Temperature (˚C)

:

Actual volume (mL)

:

Average volume (mL)

:

Standard deviation, σ

:

Relative standard deviation (σ / x )

:

iii) 100 mL volumetric flask
Trial 1

:

Temperature (˚C)

:

Actual volume (mL)

:

Average volume (mL)

:

Standard deviation, σ

:

Relative standard deviation (σ / x )

Trial 3

:

Mass of water (g)

Trial 2

:

Mass of flask (g)

Trial 3

Trial 1

Mass of flask + water (g)

Trial 2

:

iv) 50 mL measuring cylinder

Mass of flask + water (g)

:

Mass of flask (g)

:

Mass of water (g)

:

Temperature (˚C)

:

Actual volume (mL)

:

Average volume (mL)

:

Standard deviation, σ

:

Relative standard deviation (σ / x )

:
[4 x 15 marks]

DISCUSSION
(Hints: Discuss on your findings and relate to your theory and objective of experiment)

[5 marks]

Conclusion
(Hints: Conclusion should contain summary of the results, sum up what you have learned from the lab. The conclusion should be one paragraph of 5 – 7 sentences).

[5 marks]

Appendix
Questions

1. Please tell in simplest way what a calibration is.

[4 marks]

2. Draw a flowchart for the calibration of 50 mL measuring cylinder.

[10 marks]

3. With reference to the capacity of the glassware you have chosen, give a set of reading to illustrate the meaning of good accuracy and poor precision.
.

[5 marks]
4. What does standard deviation, σ, indicate?

[5 marks]

5. A 50 mL pipet delivers 49.960g of water at 27 0 C. Calculate the volume delivered by this pipet at 28 0 C (given 1.000g of water weigh in air occupies 1.0048 mL at 28 0 C )
.

[6 marks]
Total Marks = 100

EXPERIMENT 2

EMPIRICAL FORMULAS

OBJECTIVE
To determine the empirical formula of a compound formed by a combination reaction.

INTRODUCTION
The empirical formula of a compound is the formula written with the smallest whole number ratio of moles of the elements in the compound. The percentage composition of a compound will enable to obtain the empirical formula. To determine the empirical formula experimentally, the following is required:


Determine the mass of each element in the compound.



Calculate the number of moles of each element.



Express the ratio of the moles of each element as integers.

For example, when a compound of nitrogen and oxygen is analyzed, it is found that the sample contains 0.500g N and 1.142g O.

The numbers of moles of these elements are:

0.500g X 1 mol N = 0.0357 mol N
14.0g N

;

1.142g X 1 mol O = 0.0714 mol O
16.0g O

The mole ratio of nitrogen : oxygen is 0.0357 : 0.0714. The mole ratio in integers is 1 : 2
(divide each number by the smaller number i.e 0.0357 = 1 ; 0.0714 = 2 )
0.0357
0.0357

Hence, the empirical formula is NO2. The formula NO2 states that 14.0 g (1mol) of nitrogen combines with 32.0 g (2 mol) of oxygen to form 46.0 g (1 mol ) of NO2.

In this experiment, a combination reaction of magnesium and oxygen is used to determine the empirical formula of magnesium oxide. The initial mass of the magnesium and the mass of the magnesium oxide formed are measured.

MATERIAL AND METHOD

Chemicals: Mg ribbon
A. Preparation of a clean crucible
1. Take a clean crucible and a lid and check the crucible for any crack or other flaws. If any is found, request for another one.
2. Place the crucible and lid on a clay triangle and heat with an intense flame for 5minutes.
3. Cool the crucible and lid to room temperature in a desiccator.
4. If the crucible is still dirty, move the whole setup to the fume cardboard and add 1-2 mL 6M HNO3 and gently heat the crucible, evaporate the acid to dryness.
5. Repeat heating with intense flame and subsequent cooling of the crucible.
6. When cool, remove the crucible and lid from the desiccator with tongs and weigh it with the analytical balance (±0.0001g). Record the weight. Do not handle the crucible with your fingers for the rest of the experiment (so that no dirt is transferred to it).

B. Combination Reaction of Magnesium (Mg) and Oxygen (O)
1. Prepare the sample:
a) Polish about 0.2-0.3 g of magnesium ribbon with steel wool.
b) Weigh the clean magnesium ribbon and record the weight to + 0.0001g. Curl the ribbon and place it in the crucible.
c) Weigh the crucible, lid and magnesium. Record this weight to + 0.0001 g.
2. Heat the Sample in Air
a) Place the crucible with the magnesium ribbon and lid on the clay triangle
b) Heat slowly and lift the lid occasionally (now and then) to allow access of air to the magnesium ribbon.

Caution: Rapid oxidation of the Mg will occur if too much air comes in contact with the
Mg and the Mg will burn brightly. If it does, place the lid immediately on the crucible, allow it too cool and repeat the experiment from Part A.
3. To Complete Reaction
a) Continue heating until there is no change in the appearance of the magnesium ash in the crucible.
b) Remove the lid. Continue heating the crucible and ash for about 30 seconds.
c) Remove the heat and cool the crucible to room temperature in a desiccator.
d) Take the mass of crucible, ash and lid and record it.

4. Test for Completeness of Reaction
a) Add a few drops of water to decompose any magnesium nitride that may have formed. b) Reheat the sample for 1 minute. Allow it to cool and repeat determination of the mass of the crucible, lid and ash.
c) If this mass is greater than ±1% from that recorded in Part B.3; repeat Part B.4.
d) Wash the cool crucible with a dilute solution of 6 M HCl. Discard the solution into the wash acids container. Rinse with tap water and then with deionised water.
5. Calculations
a) Determine the ratio of the number of moles of magnesium to the number of moles of oxygen. b) Determine the formula of the pure compound.

Notes:
You must complete 3 trials of the experiment.

RESULTS

A) DATA:

No

Trial 1

1

Mass of crucible + lid + Mg (g)

3

Trial 3

Mass of crucible + lid (g)

2

Trial 2

Mass of crucible + lid + Magnesium oxide:
1st measurement (g)
2nd measurement (g)
3rd measurement (g)

4

Mass of magnesium oxide (g)

(3 x 10 Marks)

Calculate the empirical formula of the compound for each trial.

(3 x 10 Marks)

The empirical formula of magnesium oxide: _____________________________

(10 Marks)
DISCUSSION

(Hints: Discuss on your findings and relate to your theory and objective of experiment)

(5 marks)

CONCLUSION
(Hints: Conclusion should contain summary of the results, sum up what you have learned from the lab. The conclusion should be one paragraph of 5 – 7 sentences).

(5 marks)

Appendix

Questions

1. What is the difference between an empirical formula and molecular formula?

(5 marks)

2. Draw a flowchart for the determination of empirical formula of magnesium oxide

(10 marks)
3. Write a word equation indicating the process that took place in the above experiment.

(5 marks)

4. A compound consisting of 62.02% C, 13.88% H and 24.10% N has a relative molecular molecular mass of 116.21. What is the molecular formula?

(10 marks)

Total Marks = 100
EXPERIMENT 3

STANDARDIZATION OF SODIUM HYDROXIDE SOLUTIONS

OBJECTIVES:




To understand the applications of neutralization titration
To standardize a sodium hydroxide solution with potassium hydrogen phthalate
To understand the calculation for titration in order to get the concentration of the standard solution INTRODUCTION

In volumetric analysis which involves a titration, the volume of reagent needed to react with an analyte is measured. Increments of the titrant from the burette are added to the analyte solution until the reaction is complete.

Standardization
The concentration of titrant is known if it is prepared by dissolving a weighed amount of pure reagent in known volume of solution. This reagent is called a primary standard. It should be >>99.9% pure and stable when heated. The titrant may not be available as a primary standard. However, the concentration of the titrant can be determined by titrating it with a weighed primary standard. This procedure is called standardization and we call the titrant a standard solution.

In this experiment your solution of NaOH will be standardized by titrating it against a very pure sample of potassium acid phthalate (KHC8H4O4) of known mass.

NaOH (aq)

+

KHC8H4O4 (aq)  NaKHC8H4O4 (aq) + H2O (l)

Consider the following reaction:

aA

+

bB

product

If VA ml of A with molarity MA, reacts completely with VB ml of B with molarity
MB in the titration, the relationship can be shown as,
MAVA
--------MBVB

a
= -------b

Practically, it is difficult to determine the equivalent point, thus a third substance called indicator that gives a change when the substance is added with slightly excess quantity is used.
When the change occurs, it is known as the end point. It is not equal to the equivalent point.
The equivalence point occurs when the amount of titrant added is sufficient for a stoichiometric reaction with analyte. What is actually measured is the end point which there is change in color of the indicators used in the titration.

Volumetric analysis also can be used to determine the concentration of unknown solution. It is based on the stoichiometry of the reaction involved. The reaction must rapid and complete. In this method, a standard solution from the burette is added to a solution of unknown concentration until the equivalent point is reached. To do this, you will accurately measure with a burette the volume of your standards base that is required to exactly neutralize the acid present in unknown.
MATERIALS AND METHODS

MATERIALS

Apparatus / Glassware
100 ml Erlenmeyer Flasks
50 ml Graduated Cylinder
Titration Set up
Spatula
Burette
Top loading balance  0.01g

Chemicals
50 % W/W NaOH
Potassium hydrogen phthalate, KHP
Distilled Water

METHODS

1)

Standardization of NaOH
i.

ii.

To each flasks add 50 ml of distilled water from a graduated cylinder and shake the flask gently until the KHP is dissolved. Add 2 drops of phenolphthalein to each flask. iii.

Titrate each of them with the NaOH solution to the end point with the first appearance of a faint pink color that persists for 15 second. The color will fade as
CO2 from the air dissolves in the solution.

iv.

11)

Weigh accurately three samples of about 0.7 g to 0.9 g solid KHP into each of the three clean, numbered Erlenmeyer flasks.

Calculate the average molarity, the standard deviation and relative standard deviation (S/X)

Diluted Standardization NaOH

i.

Prepare a diluted standardization NaOH (diluted 5 times) in a 250 ml volumetric flask.

ii.

Weigh accurately three samples of about 0.7 g to 0.9 g solid KHP into of the three clean, numbered Erlenmeyer flask.

iii.

To each flask add 50 ml of distilled water from a graduated cylinder and shake the flask gently until the KHP is dissolved. Add 2 drops of phenolphthalein to each flask.

iv.

Titrate each of them with NaOH solution to the end point with the first appearance of a faint pink color that persist for 15 seconds. The color will fade as CO2 from the air dissolves in the solution.

v.

Calculate the average molarity, the standard deviation and relative standard deviation (S/ X)

RESULTS
1)

STANDARDZIDATION OF NaOH SOLUTION

TRIAL 1
Mass of KHP (g)
Final burette reading, NaOH (mL)
Initial burette reading, NaOH (mL)
Volume of NaOH (M)
Molarity of NaOH (M)
Average molarity of NaOH (M)
Standard deviation, σ
Relative standard deviation (σ / x )

(Please show all calculation steps in your Jotter book)
11)
DILUTED STANDARDIZED NaOH SOLUTION

TRIAL 2

TRIAL3

each

TRIAL 1

TRIAL 2

TRIAL3

Mass of KHP (g)
Final burette reading, NaOH (mL)
Initial burette reading, NaOH (mL)
Volume of NaOH (M)
Molarity of NaOH (M)
Average molarity of NaOH (M)
Standard deviation, σ
Relative standard deviation (σ / x )

(Please show all calculation steps in your Jotter book)

[2x 30= 60 marks]
DISCUSSION

(Hints: Discuss on your findings and relate to your theory and objective of experiment)

[5 marks]

Conclusion
(Hints: Conclusion should contain summary of the results, sum up what you have learned from the lab. The conclusion should be one paragraph of 5 – 7 sentences).

[5 marks]

Appendix

QUESTIONS

1. Define the terms below:

i.

Titration

[2 marks] ii. Titrant

[2 marks]
2. What volume of a 1.420 M NaOH solution is required to titrate 25.00 mL of a 4.50 M
H2SO4 solution?

2NaOH + 2NaOH 

Na2SO4 + 2H2O 4

[6 marks]

3.

i)

What is the concentration of Solution and list four most common units of
Concentration?

[5 marks]

ii)

50 gram of NaOH is dissolved in 50 gram of water. What is concentration (%w/w) of the solution formed?

[5 marks] iii) Dilution is the procedure for preparing a less concentration solution from a more concentrated solution. Please show a flow chart and calculation step on how to prepared a diluted standardization NaOH (diluted 5 times) in a 250 ml volumetric flask

[10 marks]

EXPERIMENT 4

ACIDS, BASES, BUFFERS AND THE DETERMINATION OF pH

OBJECTIVES
1. To determine the pH of various concentrations of acids, bases and salts.
2. To learn the concept of buffers.

INTRODUCTION

In this experiment, the pH of various solutions of acids, bases and salts will be determined using a pH meter.
Strong Acids and Bases
When a strong acid such as HCI or a strong base such as NaOH dissolves in water, it dissociates completely into ions as shown in the equations below:+

HCI(g)

+ H O (1)

-

H O (aq) + CI (aq)

2

3

NaOH(s) + H O (1)

OH- (aq)

Na+(aq) +

2

+

For a strong acid or strong base, the hydronium ion concentration, [H O ] and
3

hydroxide ion concentration, [OH-], is equal to the concentration of the acid and base respectively. The hydronium ion concentrations are expressed logarithmically as pH values. The letter p means –log
10.

+

pH = -log [H O ]
10

3

+

The lower the pH, the higher the concentration of [H O ] and hence the higher the
3

acidity.

-14

Hydronium ion concentrations commonly range from 1.0 M to 10
+

0

to [H O ] by the ion product of water, K
3
+

-

M. [OH ] is related

-

-14

[H O ][OH ]= 1.0 x 10
3

w

which at 25 C has the value

K

w

=

, Therefore:
+

-

-log K = -log [H O ] – log [OH ] = 14.00 w 3

and pKw = pH + pOH = 14.00
+

-

For a neutral solution, [H O ] = [OH ] = 10-7 M or pH = pOH = 7
3

Solutions with pH values less than 7 are acidic and solution with pH values greater than
7 are basic.

Weak Acids and Bases
Weak acids and bases do not dissociate completely in solution unlike strong acids and bases. They ionize partially as shown for acetic acid and ammonia in the following equations. H O+(aq) + C2H3O2-(aq)

HC2H3O2(aq) + H2O(1)
NH3(aq)

3

NH4+(aq) + OH-(aq)

+ H2O(1)

The equilibrium constants for these reactions are known as dissociation or ionization constants. The expressions for the dissociation constants for acetic acid and ammonia are: Acetic acid: Ka = [H3O+][C2H3O2 -]
-5

[HC H O ]
2

3

2

= 1.75 x 10

Ammonia: K = [NH4+][OH-] b [NH ]
3

= 1.75 x 10 -5

K is the acid dissociation constant and K is the base dissociation constant. From a b

these expression, notice that the greater the percentage of ionization (the stronger the acid or base) the larger the K value.
Salts of weak acids and bases
Soluble salts will dissociate completely into their ions when they dissolve in water.
The anions of strong acids and the cations of strong bases do not react with water.
However, the anions and cations of weak acids and bases react with water and this process is called hydrolysis. For example, a solution of sodium acetate will be basic due to the hydrolysis of the acetate ion as shown in the following equation:
-

-

C H O (aq) + H O(1)
2

3

2

2

HC H O (aq) + OH (aq)
2

3

2

Buffer
Buffers occur in many biochemical systems. Buffers resist pH changes, that is, they prevent significant pH changes when an acid of base added to a system. They are solution comprising of a weak acid and its conjugate base (such as acetic acid and sodium acetate) or a weak base and its conjugate acid (such as ammonia and ammonium chloride).

This experiment contains two parts.
In Part A, you will measure the pH values of several solutions. Calculate the [H 3O+] and [OH] for each solution.
In Part B, you will compare the pH change in a buffer solution and a non buffer solution when a strong base is added.

MATERIALS AND METHODS

Materials

Chemicals:
0.1 M HC1
0.001 M HCI

0.05 M HC1
0.1 M CH3CO2H

0.1 M NaOH
0.1 M NaCl
0.1 NH4Cl

0.1 M NH3
0.1 NaOCOCH3
Vinegar sample, Coca Cola sample

Solution X and Y

1 M NaOH

Equipment: pH Meter

Methods
A. pH measurement using a pH meter
Before any pH measurement of a solution is made, make a rough guess of the pH value for the solutions, so that you get close to the correct value. Your technician will show you how pH is to be measured with the pH meter.
B. Buffers
Two solutions X and Y will be given to you.
-3

X = 1 x 10 M HCI

and

Y = 0.1M CH CO H + 0.1 M
3

2

NaOCOCH

3

(Y is a solution containing two compounds) 1. Transfer about 30mL of solution X to a small beaker.
2. Determine its pH with the pH meter.
3. With the pH electrode still in the solution, add 1 drop of 1 M NaOH, stir, and

read/ record the pH.
4. Add 9 more drops of 1 M NaOH, stir and read/record the pH again.
5. Add 15 more drops and read/record the pH.
6. Clean out the beaker and repeat the experiment with solution Y.

RESULTS
A. pH measurements
*pH
No

Solution

1

0.1 M CH CO H

5

0.1 M NaOH

6

0.05 M NaOH

7

0.1 M NH

8

0.1 M NaCl

9

0.1 NaOCOCH

10

0.1 NH 4 Cl

11

deionised water

12

tap water

13

Vinegar

14

3

0.001 M HCI

4

(recorded)

0.05 M HC1

3

(predicted)

+

[H O ]

0.1 M HCI

2

pH

Coca Cola

3

2

3

3

*Please show your predict pH calculation steps in your Jotter book- if any

B. Buffers

-

[OH ]

Solution X
Total drops of 1M NaOH

Solution Y pH Total drops of 1M NaOH

0

0

1

1

10

10

25

pH

25

* Compare the pH changes obtained for solution X and Y with the addition of
NaOH.
Explain if there is any difference

[60 marks]
DISCUSSION

(Hints: Discuss on your findings and relate to your theory and objective of experiment)

[5 marks]

Conclusion
(Hints: Conclusion should contain summary of the results, sum up what you have learned from the lab. The conclusion should be one paragraph of 5 – 7 sentences).

[5 marks]

Appendix

QUESTIONS
1) What is the pH of a solution where the hydrogen ion Concentration is 10-3M?

[4 marks]

2) In NaOH solution  OH-  is 2.9 x 10 -4 M. calculate the pH of the Solution.

[4 marks]

3) Strong acids dissociate completely to H+ ions in aqueous solutions. Examples of strong acids are HNO3.
HNO3 + H2O  H3O+ + NO3 –
Calculate the pH of HNO3 0.001 M

[4 marks]

4) Conversely, weak acid do not dissociate completely to H+ ions in aqueous solutions. All organic acids are weak acids. Example is nitrous acid,
HNO2 + H2O  H3O+ + NO2 –
Calculate pKa and pH of 0.01 M nitrous acid, HNO2. (Ka = 5.1 x 10 -4 M).

[7 marks]

5) Strong bases dissociate completely to OH- ions in aqueous solutions. Examples of

strong bases are NaOH .
NaOH  Na + (ag) + OH- (aq)
Calculate the pH of NaOH 0.1 mol dm-3

[4 marks]
6) Conversely, weak bases do not dissociate completely to OH - ions in aqueous solutions.
Ammonia (NH3) and all organic bases are weak bases. Example is NH2OH,
NH2OH + H2O  NH3OH+ + OH Calculate pKb and pH of 0.1 M NH2OH, (Kb = 9.1 x 10 -9 M).

[7 marks]

EXPERIMENT 5

REDOX REACTIONS

OBJECTIVES
• To observe and predict products of oxidation-reaction reactions.
• To determine the relative reactivity of a series of metallic elements.

INTRODUCTION
Redox reactions involve a transfer of electrons. If the oxidation number of any element has increased, the element has lost electrons and has been oxidized. In such reactions, another element must have gained electrons and has been reduced. It has also experienced a decrease in oxidation number. A reducing agent donates electrons for the reduction of another substance. An oxidizing agent is a substance that gains electrons and causes the oxidation of another substance in a redox reaction. This experiment focuses on a single replacement reaction. For the reaction,

Zn(s)

+

Pb (NO3)2(aq)



Pb(s)

+

Zn (NO3)2(aq)

Lead is reduced to its elemental form. Zinc is oxidized to its form. Therefore zinc metal is a more active reducing agent than lead metal. If no reaction occurs when a metal is put unto a solution, then the elemental form of metal in solution is a better reducing agent than the metal put into the solution.

For the reaction,

F2 (g)

+

-

2 Br (aq)



-

2F (aq)

+

Br2(1)

Fluorine is reduced and bromine is oxidized to its elemental form. Therefore fluorine is a more active oxidizing agent than bromine. If no reaction occurs when an oxidizing agent is put into a solution, then the oxidized form of the negative ion already in solution is a better oxidizing agent.
This experiment contains two parts: Part A and Part B
1. Part A
In this experiment, you will first determine the relative reducing activities of Cu,
Fe, H, Mg and Zn.
2. Part B you will observe and study several redox reactions
A redox reaction is divided into two half reactions, an oxidation and a reduction.

MATERIALS AND METHODS

MATERIALS
Chemicals:
Copper metal

0.1 M K2C2O4

Magnesium ribbon

6 M H2SO4 acid

Zinc metal
Steel wool
0.1 M CuSO4
3 M H2SO4 acid
0.1 M ZnSO4
0.01 M KMnO4
0.01 M K2Cr2O7
0.1 M Fe(NH4)2(SO4)2

METHODS

Part A
Metal as reducing agents
Label four test tubes 1-4 (15mL test tubes)
2
1. Put a 1 cm piece of copper metal into the first test tube.
2. Put a 2 cm long piece of magnesium ribbon into the second test tube.
2

3. Put 1 cm piece of zinc metal into a third test tube.
4. Put a small ball of steel wool into the fourth test tube.
5. Add a few millilitres of 0.1 M CuSO4 to each test tube. Note if there is any reaction. If a reaction occurs, write a balanced equation for the reaction. If not, write NR, for no reaction. 6. Discard the used metal squares in the waste container.
7. Rinse the tubes with deionised water, and test some new metal squares in the same way as before with 3 M H2SO4 acid.
8. Hold a lighted splint over any test tube in which bubbling occurs.
9. Clean the test tubes as before and test each metal square with 0.1 M ZnSO4 solution.

Part B
Redox reactions
1. Clean and label 4 small test tubes (3 mL test tubes)
2. Place 1 mL of each solution listed as Solution X in Table 5.1 into the test tubes.
3. Slowly add up to 1 mL of Solution Y until a permanent change is observed. Some reactions may be slow.
4. Record your observations for each mixture that shows a reaction.
5. Write balanced redox reactions for any reaction observed. Underline the oxidising agent in each reaction. The half reactions are given in Table 5.2.
6. Dispose of the test solutions into waste containers provided. Rinse the test tubes with tap water.

Table 5.1 Solutions for redox reactions
Tube

Solution X

Solution Y

1

0.01 M KMnO4

+ 2 drops 6 M H2SO4

Drops of 0.1 M Fe(NH4)2(SO4)2

2

0.01 M KMnO4

+ 2 drops 6 M H2SO4

Drops of 0.1 M K2C2O4

3

0.01 M K2Cr2O7 + 2 drops 6 M H2SO4

Drops of 0.1 M Fe(NH4)2(SO4)2

4

0.01 M K2Cr2O7 + 2 drops 6 M H2SO4

Drops of 0.1 M K2C2O4

Table 5.2 Half reactions
-

+

-

Mn04 (aq) + 8H (aq) + 5e
Purple



2+

Mn (aq) + light pink to colourless

2+
Cr2O7 (aq) + 14H (aq) + 6e 
Orange

2+

Fe (aq)



Light green to colourless
2-

C2O4 (aq) colourless 3+

Fe (aq)

3+

2Cr (aq) + 7H2O(1) dark green

+

e

red brown



-

2CO2 (g) + 2e colourless 4H20(l)

-

RESULTS
Part A
Metal as reducing agents
Test Tube No

0.1 M CuSO4

3 M H2SO4

0.1 M ZnSO4

1 (Copper)
2 (Magnesium)
3 (Zinc)
4 (Steel wool)

Metals as reducing agents
1. If a reaction occurs, complete and balance the equation. Write NR if there is no reaction observed. Assume elemental iron will react to form iron (II).
Balanced Equation
CuSO4(aq)

+

Cu(s)

CuSO4(aq)

+

Mg(s)

CuSO4(aq)

+

Zn(s)

CuSO4(aq)

+

Fe(s)

H2SO4(aq)

+

Cu(s)

H2SO4(aq)

+

Mg(s)

H2SO4(aq)

+

Zn(s)

H2SO4(aq)

+

Fe(s)

ZnSO4(aq)

+

Cu(s)

ZnSO4(aq)

+

Mg(s)

ZnSO4(aq)

+

Zn(s)

ZnSO4(aq)

+

Fe(s)

2. Arrange the four metals and hydrogen in the order of their activities beginning with the best reducing agent and ending with the poorest reducing agent.
i.
ii. iii. iv.
v.

Part B
Solution for redox reaction

Test tube

Solution X

Observation and Redox Reaction Equation (if any)
0.1 M Fe( NH4)2(SO4)

1

0.1 M K2C2O4

0.01 M KMnO4
+
2 drops 6 M H2SO4

2

0.01 M KMnO4
+
2 drops 6 M H2SO4

3

0.01 M K2Cr2O7
+
2 drops 6 M H2SO4

4

0.02 0.01 M
K2Cr2O7
+
2 drops 6 M H2SO4
(70 Marks)

DISCUSSION
(Hints: Discuss on your findings and relate to your theory and objective of experiment) [5 marks]

Conclusion
(Hints: Conclusion should contain summary of the results, sum up what you have learned from the lab. The conclusion should be one paragraph of 5 – 7 sentences).

[5 marks]
Appendix

1. The following reaction does not occur spontaneously:
2+

-

Pb (aq) + 2Cl (aq)  Pb(s) + Cl2(g)

2+

Which is the more active oxidising agent, Pb or C12 ?

(2 marks]

2.. Elemental Zn reacts with aqueous Ni (II) chloride to give zinc chloride and nickel.
a. Which is the more active metal (better reducing agent) zinc or nickel?

[3 marks]

b. Write down the oxidation and reduction half reactions and net ionic equation for the system. [6 marks]

3. When iron metal is placed in a solution containing lead (II) ion, iron (II) and lead metal are produced.

a. Which is the more active metal, iron or lead?

[3 marks]
b. Write down the oxidation and reduction half reactions and the net ionic equation for the system.

[6marks]

EXPERIMENT 6

QUALITATIVE ANALYSIS OF COMMON IONS

OBJECTIVES
* To perform specific test for identification of some common cations and anions.
* To identify the cations/anions present in unknown samples.

INTRODUCTION
[Qualitative analysis is a set of procedures to detect the presence of a specific element/ion] Ionic compounds are of 2 types: soluble and insoluble. Soluble salts have high solubility values whereas the insoluble salts have very low solubility values. Soluble salt dissociates into cation and anion when it is put in water. When cation and anion of an insoluble salt are present together in one sample, the 2 react forming the new salt which then appears as solid/precipitate because of its low solubility.
Soluble salts: all nitrates , all chlorides except of Pb2+, Hg22+ and Ag+ and (refer to lecture )
Insoluble salts: most sulphate except of Na+, K+ and… (Refer to lecture notes)
This experiment contains two parts: Part A and Part B

3. Part A
In this experiment, you will perform specific test for identification of some common cations
4. Part B
In this experiment, you will perform specific test for identification of some common anions

Part A

QUALITATIVE ANALYSIS OF COMMON CATIONS
The precipitating reagent used for the qualitative analysis of Pb2+, Hg22+ and Ag+ is chloride ion (Cl‾) in acidic solution. The Cl‾ ion is obtained from dissolving in water soluble chlorides eg. NaCl ; KCl.
The chlorides of Pb2+, Hg22+ and Ag+ (i) are white in colour and (ii) are insoluble in cold water.
Solutions that have these ions in them are used to show the above TWO (2) Properties.
The precipitation reactions for these cations that occur are as follows:
Ag+(aq)

+ Cl-(aq)



Hg22+(aq) + 2Cl-(aq) 
Pb2+(aq) + 2Cl-(aq)



AgCl(s)
Hg2Cl2(s)
PbCl2(s)

The PbCl2 dissolves in hot water. If the 3 chlorides are present together as a mixture, by heating the mixture in some water i.e making it hot and then filter, the Lead (II) chloride (PbCl2) can be separated from Ag Cl and Hg2Cl2.
PbCl2(s) HEAT  Pb2+(aq) + 2Cl-(aq)


The Pb2+ ion in solution is tested by adding a solution of K2CrO4. Pb2+ ion forms a yellow precipitate with the chromate ion, CrO42-.
Pb2+(aq) + CrO42-(aq)



PbCrO4(s) yellow The other two insoluble chlorides, AgCl and Hg2Cl2 can be separated by adding aqueous ammonia solution. AgCl dissolves forming the complex ion, Ag(NH3)2+.
AgCl(s) + 2NH3(aq)  Ag(NH3)2+(aq) + Cl-(aq)
Ammonia also reacts with Hg2Cl2 to produce finely divided metallic mercury (black in colour) and a compound, Hg(NH2)Cl (white in colour)

Hg2Cl2(s) + 2NH3(aq)  Hg(NH2)Cl(s) + Hg (l) + NH4+ (aq) + Cl-(aq) white black
As this reaction occurs, the solid appears to change colour from white to black to grey. The solution containing the Ag(NH3)2+ needs to be tested further to confirm the presence of silver. Dilute HNO3 acid is added to the solution which destroys the complex ion and reprecipitates silver chloride.
Ag(NH3)2+(aq) + 2H+(aq)  Ag+(aq) + 2NH4+(aq)
Ag+(aq) + Cl- (aq) 
AgCl(s)
________________________________________________
Ag(NH3)2+(aq) + 2H+(aq) + Cl-(aq)  AgCl(s) + 2NH4+(aq) white MATERIALS AND METHODS

MATERIALS

Chemicals:
Known solutions with ions (separately and together):

Known Solutions with ions

Label

containing Pb2+

I

containing Hg22+

II

containing Ag+

III

containing Ag+, Hg22+ and Pb2+

IV

6 M HCl acid
6 M acetic acid
0.1 M K2CrO4
6 M NH3
6 M HNO3
Unknown solution.

METHODS
1.0 Qualitative Analysis of Pb2+ ions
a) Put 1 mL of the solution I in a small test tube .Then add 2 drops of 6 M HCl. Mix thoroughly with your stirring rod. Note down your observation.
b) Place the test tube in boiling water bath. Leave the test tube in the bath for 3 minutes; while stirring occasionally with a glass rod. Note down your observation. c) Divide the hot solution into 2 test tubes.
d) Take one test tube to the tap and cool the tube under running water. Note down your observation.
e) To the other tube, add 1 drop of 6 M acetic acid and a few drops of 0.1 M K2CrO4
Note down your observation. (PbCrO4 is yellow precipitate)
2.0 Qualitative Analysis of Hg22+ ions
a) Put 1 mL of the solution II in a small test tube .Then add 2 drops of 6 M HCl.
Mix thoroughly with your stirring rod. Note down your observation.
b) Place the test tube in boiling water bath. Leave the test tube in the bath for 3 minutes; while stirring occasionally with a glass rod. Note down your observation. c) Filter OR centrifuge the solutions
d) Put all the precipitate together into a test tube. Add 1 mL of 6 M NH3 and stir thoroughly. Note down your observation.
e) Discard the solid that is formed into the proper waste container for solids.
3.0 Qualitative Analysis of Ag+ ions
a) Put 1 mL of the solution III in a small test tube .Then add 2 drops of 6 M HCl.
Mix thoroughly with your stirring rod. Note down your observation.
b) Place the test tube in boiling water bath. Leave the test tube in the bath for 3 minutes; while stirring occasionally with a glass rod. Note down your observation. c) Filter OR centrifuge the solutions

d) Put all the precipitate together into a test tube, add 1 mL of 6 M NH3 and stir thoroughly. Note down your observation.
e) Add HNO3 drop wise, while stirring, to the supernatant until it is acidic to litmus paper (blue to red) by dipping a glass rod into the test solution and touching it on the litmus paper. Note down your observation.
4.0 Qualitative Analysis of Known Mixture (Pb 2+, Hg22+ and Ag+ ions) and Unknown
Mixture
Test for Pb2+ ion
a) Put 2 mL of the solution IV (Known Solution) / Unknown Mixture in a small test tube .Then add 4 drops of 6 M HCl. Mix thoroughly with your stirring rod.
Note down your observation.
b) Filter the mixture OR centrifuge the mixture. Ask your instructor for instructions on how to use the centrifuge correctly. After centrifuge, in the test tube there will be a layer of clear solution (supernatant) on top of a thin layer of white precipitate.
c) Add 1 more drop of 6 M HCl to the solution from filtration or to the supernatant
d) Centrifuge again if any additional precipitate forms. Pour the supernatant into another test tube
e) Wash the precipitate from step (2) above with 1 to 3 mL of distilled water, i.e
f) add 1-3 mL of distilled water from a wash bottle to the precipitate in the test tube, stir, centrifuge/filter and pour out the liquid (this liquid may be thrown away)
g) Add 3 mL distill water to the precipitate. Place the test tube in boiling water bath.
Leave the

test tube in the bath for 3 minute; while stirring occasionally with a

glass rod. This should dissolve the PbCl2.
h) Centrifuge/filter the hot mixture. The precipitate is used for further tests for
Hg22+ and Ag+.
i) Add 1 drop of 6 M acetic acid and a few drops of 0.1 M K2CrO4 to the supernatant. A yellow precipitate of PbCrO4 will form if Pb2+ is present.
Test for Hg22+ ion

a) To the precipitate from step 4 (g) the previous test, add 1 mL of 6 M NH3 and stir thoroughly.
b) Centrifuge and decant the supernatant into a small test tube. A grey to black precipitate proves the presence of Hg22+. Discard into the proper waste container for solids.
Test for Ag+ ion
a) Add 6 M HNO3 dropwise, while stirring, to the supernatant until it is acidic to litmus paper (blue to red) by dipping a glass rod into the test solution and touching it on the litmus paper. A white precipitate confirms the presence of
Ag+.

Part B

QUALITATIVE ANALYSIS OF COMMON ANIONS
In this experiment, qualitative tests for SO42-, PO43-, CO32-, S2-, CI- and NO3- ions will be carried out. These ions: SO42-, PO43-, CO32-, S2-, CI‾ and NO3- are from the soluble salts.
These tests involve a particular chemical reaction for a particular anion. To test for a particular anion in a solution, selected chemicals are added to the solution. If the anion is present, a unique chemical reaction occurs which will produce a product with a particular color or form, or a gas that can be identified. If the reaction does not occur, the anion is absent.
Sulfide ion, S2The sulfide ion is an anion of a weak acid, H2S. When a strong acid is added to a solution that has S2- in it, gaseous H2S will be liberated. The reaction is represented by the chemical equation:

S2-(aq) + H+ (aq)
H2S(g)
The odour of H2S can be used for identification and can be further confirmed with the formation of a black lead sulfide, PbS.
H2S + Pb (OOCCH3)2

PbS (s) + 2 HOOCCH3 black Sulphate ion, SO42It forms a white finely divided precipitates with barium ion. The reaction is represented by the chemical equation:
Ba2+ (aq) + SO42- (aq)

BaSO4(s)

The precipitate, BaSO4, is not soluble in acidic solutions. (An acidic solution of Ba2+ ions is used to test for the presence of SO42- ion)

Carbonate ion, CO32Acidification of a solution that contains carbonate ions will liberate carbon dioxide which is colourless and odourless gas. The reaction is represented by the chemical equation:

CO32-(aq) + 2H+(aq)

H2O(I) + CO2(g)

The carbon dioxide gas is bubbled through an alkaline solution containing calcium ion, carbonate is produced and a white precipitate of calcium carbonate is formed. The reaction is represented by the chemical equation:
CO2(g) + 2OH-(aq) + Ca2+(aq)

CaCO3(s) + H2O(1)

Phospate ion, PO43The presence of the phospate ion is detected through the formation of a characteristic precipitate i.e. a yellow precipitate. An acidic solution is necessary.

Ammonium

molybdate solution is added to the test solution. An excess of ammonium ion is added to

favour formation of a yellow precipitate of ammonium phosphomolybdate. The reaction is as follows:3NH4+(aq) + 12MoO42-(aq)+ 3H+ + PO43-(aq)

21H+(aq) + (NH4)3PO412MoO(s). yellow Sulfide ion interferes with this test. If it is present in the test solution, it should be removed by adding HCI to the solution than boiling it to drive out the H2S gas.
Chloride ion, ClChloride ion is precipitated as a white precipitate of silver chloride with the addition of silver nitrate solution. The reaction is as follows:Ag+(aq) + CI-(aq)

AgCI(s)

The presence of chloride ion is further confirmed by adding ammonia solution to dissolve the silver chloride by the formation of soluble, complex ion, Ag(NH3)2+ and CI-. AgCI will precipitate out again with the addition of nitric acid.
Ag(NH3)2+(aq) + CI-(aq)

AgCI(s) + 2NH3(aq)

Ag(NH3)2+(aq) + CI-(aq) + 2H+(aq)

AgCI(s) + 2NH4+(aq)

Nitrate ion, NO3All nitrate salts are soluble in aqueous solutions. The identification of nitrate does not involve the formation of a precipitate. The ‘brown ring’ test is used. The brown ring is formed when the nitrate ion is reduced to nitric oxide, NO, by iron (II) in an acidic solution. The reaction is as follows:Conc.H2SO4
-

2+

+

3Fe3+(aq) + NO(aq) + 2H2O(1)

NO3 (aq) + 3Fe (aq) + 4H (aq)

The nitric oxide, NO, reacts with excess iron(II) ions to form a brown complex ion,
Fe(NO)2+ at the interface of the aqueous layer and the concentrated acid layer.
Fe2+ (aq) + NO(aq)

Fe(NO)2+(aq)

An excess of iron (II) ions is essential for this test

MATERIALS AND METHODS

MATERIALS

Chemicals:
0.5 M ammonium molybdate (8.8g/100mL)
0.02 M Ca(OH)2
0.1 M BaCI2
Lead acetate,[Pb(CH3COO)2] paper
Saturated FeSO4 (0.16 g mL-1 FeSO4-7H2O)
6M HCI, 6M HNO3, 6M NH3, 3M H2SO4
Concentrated H2SO4 acid
Known Solutions
Solutions containing ions

Label

0.1 M Na2S (freshly prepared)

I

0.1 M Na2SO4

II

0.1 M Na2CO3

III

0.1 M Na3PO4

IV

0.1 M NaCI

V

0.1 M AgNO3

VI

Unknown Solution
Note:


Acids and bases used in this experiment are corrosive. Be sure to wash them off immediately if you get them on your skin.

METHODS

1.0 Test for sulphide ion.
a) To about 2 mL of solution I, add an excess of 6 M HCI. You might notice a smell of rotten eggs with the H2S gas given off.
b) Moisten a piece of lead acetate paper and place it over the mouth of the test tube, then
c) heat the test tube gently in the fume hood. Formation of PbS will cause the paper to darken.
2.0 Test for sulphate ion.
a) To about 2 mL of solution II, add a few drops of 6 M HCI until the solution is acidic to litmus paper i.e. dip a glass rod in the test solution and touch it on the litmus paper.
b) Add 1 mL of 0.1 BaCI2. White precipitate of BaSO4 will form.

3.0 Test for carbonate ion.
a) Obtain a rubber stopper fitted with a bent glass tube for the 15 cm test tube. Fill the test tube with about 3 mL of solution III.
b) Insert the end of the bent glass tube into another test tube filled with 0.02M Ca(OH)2.
c) Lift the stopper and add a few drops of 6 M HCI to the solution.
d) Close the stopper immediately and heat the test tube gently over a low flame to boiling point to drive any CO2 gas produced into the Ca(OH)2 solution. Do not let any of the boiling liquid go through the tube into the Ca(OH) 2 solution.
e) A white precipitate formed in the Ca(OH) 2 solution indicates the presence of CO32ion.
4.0 Test for phosphate ion.
a) In a 15cm test tube, mix about 1 mL of 0.5 M (NH4)2MoO4 [ ammonium molybdate] reagent with 1 mL 6M HNO3.

b) Add 2 mL of solution IV. A yellow precipitate of (NH4)3PO4-12MoO3 formed slowly or after warming a few minutes in warm water bath at 60oC .
* If S2- ion is present in your unknown solution, acidify 2 mL of your unknown solution with

6 M HCI. Boil it to remove all the H2S gas. Then add it to the

clear molybdate solution.
5.0 Test for chloride ion.
a) Add a few drops of 6 M HNO3 to 2 mL of solution V to make it slightly acidic
(test with litmus paper as in 2).
b) Add 1 mL of 0.1M AgNO3. If no precipitate is formed, this will indicate there is no CI- ion in your unknown solution.
6.0 Test for nitrate ion.
a) Acidify (test with litmus paper) 2 mL of solution VI with 3 M H2SO4.
b) Add 1 mL of saturated FeSO4 (freshly prepared).
c) Hold the test tube at about a 45o angle and with a dropper (slowly and carefully), add 1 mL of concentrated H2SO4 acid down the side of the test tube. Do not draw concentrated H2SO4 acid into the bulb of the dropper.
d) Do not agitate the solution. The more dense concentrated H2SO4 acid will be below the aqueous layer. Avoid mixing the concentrated acid with the rest of the solution.
e) Allow the mixture to stand for a few minutes. A brown ring of Fe(NO)2+ at the interface confirms the presence of NO3- ion.
Note:
 Use a fresh 2 mL sample of your unknown solution for each test. Use 15 cm test tubes for the tests.

RESULTS
Part A
i) Qualitative Analysis of Common Cations (Known Sample)
Record Your Observation and write an equation as return in introduction if any

Solution

+

Cl─

Ppt. in water
And heat

Cool the
Hot solution

Hot solution
+ H+/ CrO4─

Ppt. + NH3, then
+ HNO3

Ppt. + NH3

I

II

III

IV

ii) Qualitative Analysis of Common Cations (UnKnown Sample)
Record Your Observation and write an equation as return in introduction if any

Solution

+

Unknown Sample

Cl─

Ppt. in water
And heat

Cool the
Hot solution

Hot solution
+ H+/ CrO4─

Ppt. + NH3, then
+ HNO3

Ppt. + NH3

Unknown solution code :
Cations present

:

Part B

QUALITATIVE ANALYSIS OF COMMON ANIONS

i) QUALITATIVE ANALYSIS OF COMMON ANIONS (Known Sample)
Anions
Known
Sample
S2-

SO42-

CO32-

PO42-

CI-

NO3-

Chemical change observed precipitate/gas/colour change

Equation for observed
Reaction
Yes / No

ii) QUALITATIVE ANALYSIS OF COMMON ANIONS (Unknown Sample)
Anions
Known
Sample

Chemical change observed precipitate/gas/colour change

Equation for observed
Reaction
Yes / No

S2-

SO42-

CO32-

PO42-

CI-

NO3-

Unknown solution code :
Anions present

:

(70 Marks)

DISCUSSION
(Hints: Discuss on your findings and relate to your theory and objective of experiment) [5 marks]

Conclusion
(Hints: Conclusion should contain summary of the results, sum up what you have learned from the lab. The conclusion should be one paragraph of 5 – 7 sentences).

[5 marks]

Appendix
1. Define what is meant by Qualitative analysis and Give three example of its application in industry

[5 marks]

2. Complete the following flow chart by referring theory as return in Qualitative analysis of common cations above (Part A)
Pb2+, Hg22+, Ag+
6 M HCl
White precipitate
AgCl, PbCl2, Hg2Cl2

Solutions with other ions Add hot water

White precipitate
AgCl, Hg2Cl2

Solution contains
_______D___________

Add______A_______

Add _____B______
______E______________ shows_____J______

____F________ shows ____H__________

Solution contains
______G___________
Add _____C______
____I______________ shows _______K_________

Fig. 2.1 Flowchart for qualitative analysis of Pb2+, Hg22+ and Ag+.
A, B, C – reagents that need to be added.
D, E, F, G, H – ions in solution or precipitate formed.
I, J, K – ion present
[15 marks]