1011 Lab Manual 2013 Print

School of Chemistry

CHM1011 Lab Manual Chemistry I

2013

CHM1011 - Student Laboratory Manual Introduction

T A B L E O F C O N T E N T S

(i) Introduction 2

(ii) Practical and Tutorial Timetable 2013 3

(iii) Aims of the Laboratory Course 4

(iv) Structure of the Laboratory Course 4

(v) General Information 4

(vi) Laboratory Hours and Rules 4

(vii) Proformas 5

(viii) Assessment 5

(ix) Resource Centre & Chemistry Mentors 5

(x) Learning Skills Unit (LSU) 6

(xi) Code of Practice - Teaching 6

(xii) OHS&E Information 7

Introduction – Health and Safety 9

Safety Quiz 9

Exercise 1 Determining the Waters of Hydration: CuSO4.xH2O 11

Exercise 2 Quantitative Analysis of Vitamin C contained in Ribena® 15

Exercise 3 Geometric Models: VSEPR Method 19

Exercise 4 Transformations of Copper 25

Exercise 5 Acid-Base Titration Curves 29

Exercise 6 IDEA#1: QC Calorimetric Analysis of Baking Powder 34

Exercise 7 IDEA#2: Timing the Iodine Clock 38

Exercise 8 The TEK Experiments 41

Periodic Table of Elements 44

(i) INTRODUCTION

Welcome to the CHM1011 laboratory experience. Chemistry is fundamentally an experimental science, and it is the joy of designing and testing ideas with experiments that both explore and impact on the world around us that has made chemistry so fascinating and innovative. This practical course is intended to introduce you to some of the most important laboratory skills and techniques used in chemistry, skills that should be useful to you even if chemistry doesn't become your major! However, we hope that this lab course will do more than just give you good technique. As practicing chartered chemists, we would like you to see how some of the theoretical principles of chemistry were uncovered, and how chemistry is applied on a day-to-day basis all over the world. But above all, we hope that you will enjoy the course, make new friends, and acquire some of the interest and fascination for chemistry that your lecturers and demonstrators have.

Dr. Chris Thompson Coordinator, 2013

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(ii) CHEMISTRY CHM1011 PRACTICAL TIMETABLE 2013 GROUPS A, B, C D, E, F G, H, I

Week Date LABORATORY EXERCISES

25/2 O-Week

1 4/3 1 Hour Induction Session

Ex 1 - Determining the Waters of Hydration: CuSO4.xH2O 2 11/3

Bays 1/2 (Blue/Violet) Bays 5/6 (Red/Orange) Bays 7/8 (Yellow/Green)

Ex 2 - Quantitative Analysis of Vitamin C contained in Ribena®

Ex 4 - Transformations of Copper

Ex 3 – Geometric Models: VSEPR

Method

3 18/3

Bays 7/8 (Yellow/Green) Bays 5/6 (Red/Orange) Bays 1/2 (Blue/Violet)

4 25/3 Bye Week – No Labs This Week

1/4 Mid-Semester Break (one week)


Method



5 8/4

Bays 1/2 (Blue/Violet) Bays 7/8 (Yellow/Green) Bays 5/6 (Red/Orange)



Method


6 15/4

Bays 5/6 (Red/Orange) Bays 1/2 (Blue/Violet) Bays 7/8 (Yellow/Green)

7 22/4 Bye Week – No Labs This Week

Ex 6 - IDEA#1 Analysis of Baking

Powder

Ex 7 - IDEA#2 Timing the Iodine

Clock

Ex 5 - Acid-Base Titration

8 29/4

Bays 7/8 (Yellow/Green) Bays 1/2 (Blue/Violet) Bays 5/6 (Red/Orange)


Clock



Powder

9 30/4

Bays 1/2 (Blue/Violet) Bays 5/6 (Red/Orange) Bays 7/8 (Yellow/Green)



Powder


Clock

10 7/5

Bays 5/6 (Red/Orange) Bays 7/8 (Yellow/Green) Bays 1/2 (Blue/Violet)

Ex 8 - The TEK Experiments 11 14/5

Bays 7/8 (Yellow/Green) Bays 1/2 (Blue/Violet) Bays 5/6 (Red/Orange)

12 21/5 Ex 8 - TEK Experiment Presentations

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(iii) AIMS OF THE LABORATORY COURSE

The laboratory course has three main aims.

The development of manual skills. There are many skills involved in handling laboratory instruments and equipment, in carrying out chemical transformations, and in making chemical measurements. These skills can only be learned by practice. As you become more skilled, you gain confidence both in what you do and in Chemistry as a subject. As you gain confidence, your ability to study and to master new topics improves.

The illustration of theoretical principles. Laboratory work serves to illustrate many of the somewhat abstract theoretical principles that you are taught in lectures and to demonstrate their practical application.

The gaining of experience in the practice of chemistry.

(iv) STRUCTURE OF THE LABORATORY COURSE

The class is divided into groups, so that at any laboratory session each group may be carrying out a different exercise. You should carefully check your program so that you know which exercise you will be doing and you should thoroughly read it through before the session. During the semester you will undertake eight practical exercises and a group presentation.

(v) GENERAL INFORMATION

Materials Required: During your lab session you will have access to all glassware, tools and chemicals that you will need. You will also need the following:

This laboratory manual. A note-book, ruler and calculator. Protective clothing (laboratory coat) to protect you and your ordinary clothing from

damage by chemicals or fire. Covered shoes with firm soles are required at all times when in the laboratory. Note:

You will be asked to leave the laboratory if you wear open-toe or other inappropriate footwear (including thongs, sandals and ballet shoes) during your lab session.

Safety glasses which must be worn at all times during lab. Students who wear contact lenses should purchase special wrap-around safety spectacles.

(vi) LABORATORY HOURS AND LABORATORY RULES

The laboratory class runs from 2.00 - 5.00 pm Monday to Friday afternoons and 10 am to 1 pm Monday, Wednesday and Friday mornings. The laboratory is closed during lunch hour, 12.45 - 1:45 pm. You may not start work until a Demonstrator or Supervisor is present in the laboratory.

Attendance: The lab component is a hurdle requirement for this unit. If you arrive late you will miss the safety briefing, and will not be permitted to commence the class.

Absent Due to Illness: You should provide your demonstrator with a copy of your medical certificate – this MUST be supplied. If you are absent more than once, you will need to complete an “in-semester special consideration form” at the front counter.

It is the responsibility of each student to ensure at the end of each session that the work spaces used are thoroughly cleaned and dried, and that all equipment is put away.

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NOTE: The School of Chemistry takes your well-being and safety, and that of everyone in the laboratory, very seriously. As such you can be asked to leave the laboratory for the following reasons (see page 7 Safety Rules and Safety Precautions):

Aggressive, offensive, or intimidating behaviour Behaviour which is deemed to be hazardous to yourself and/or class mates Wearing inappropriate footwear Not wearing appropriate safety equipment Using mobile phones in the laboratory

You may be allowed to continue the laboratory course but only after consultation with Dr. Chris Thompson (1st Year Coordinator).

(vii) PROFORMAS

In this unit, your lab report will be completed on a proforma sheet, available via MOODLE. You must print the proforma out in advance, complete pre-lab questions and the “Aim” and “Method” for the experiment, and then bring the proforma with you to your lab session. Your proforma must be submitted directly to your demonstrator at the end of the session.

(viii) ASSESSMENT

The laboratory assessment component constitutes 30% of the total mark for CHM1011.

To pass this subject you must pass the laboratory component of the course. Your assessment will be based on:

(a) The way in which you work in the laboratory, (b) Your attitude and attendance, (c) Your pre-lab exercises and (d) On the results that you achieve.

For each practical session each day's work is marked out of 20. Marking schemes are given in the proforma for each exercise.

You can keep track of your lab marks via MOODLE.

If you are unavoidably absent from any session you should inform both your demonstrator and the lab staff ASAP. If you have a medical certificate ensure you hand it in to your demonstrator (or to the lab manager) as soon as you can.

If more than one lab session is missed, a special consideration form should be filled in at the first year laboratory enquiries desk and a copy of a medical certificate or other verification attached. If you have serious difficulties of any kind throughout the year, you should also notify the appropriate Faculty Office in writing and talk to the Coordinator, as this will then be taken into account in assessing your performance in the final examinations. For more information please refer to the CHM1011 Unit Outline.

WARNING - A pass (50%) in lab program is required to pass CHM1011.

(ix) LUNCHTIME TUTORS IN THE CHEM ED LAB: 12-2PM DAILY

Our Chemistry Tutors are available 12-2pm daily to assist you with problem solving, lecture material, practical work or study methods. The Chem Ed Lab is up in Building 19, Level 1, Room 135C, next door to Dr. Chris Thompson’s Room. 15 minute appointments may be made with a Mentor between 12-1pm (Book in the diary at the front counter of the First Year Labs).

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Students with a limited background in chemistry are urged to contact Dr. Chris Thompson (1st Year Coordinator) as soon as possible. Special help is available, including suggestions for preliminary reading.

(x) LEARNING SKILLS UNIT, LIBRARY. The LSU in the library is an excellent place to find assistance for all your studies at Monash University. Learning skills advisers assist students in improving their academic language and approaches to learning, including:

academic English reading strategies study methods and exam preparation essay, report and thesis writing effective listening and note-taking writing for research projects problem-solving and critical thinking oral communication and presentation

Students with Language Difficulties: Inform your demonstrator and contact the LSU personnel at the Hargrave-Andrew Library. These guys can help!

(xi) CODE OF PRACTICE - TEACHING

Learning involves the participation of both staff and students; the responsibility to ensure that learning is conducted in the most efficient and effective manner is shared. This code of practice outlines the responsibilities of students and staff.

Responsibilities of students

All Monash students should observe the statutes, regulations and rules of the University. Students of the Faculty of Science also have responsibilities that include the following:

(i) to become familiar with the rules and regulations governing the degree in which they are enrolled, and to ensure that the subjects selected meet the degree requirements;

(ii) to maintain satisfactory progress; (iii) to become aware of the policies and practices of the Faculty of Science and

departments/ school from which they take subjects; (iv) to become aware of the rules and regulations concerning the use of university

computers and library facilities and observe the laws of copyright; (v) to meet assessment deadlines for work to be submitted; (vi) to submit their own original work for assessment, without plagiarising (copying the

work of any other person) or cheating; (vii) to attend all lectures, practical and laboratory classes and seminars for each subject

in which they are enrolled; (viii) to apply themselves to their studies to the best of their abilities; (ix) to conduct themselves in an orderly and proper manner in any class, library or

laboratory and not to affect adversely the working environment of others; (x) to take the initiative and consult with appropriate academic staff when problems

arise.

Responsibilities of staff

Staff of the faculty have responsibilities towards the students they teach which include:

(i) preparing and presenting material at an appropriate standard with the resources available;

(ii) informing students of the objectives, requirements and method of assessment to be used for the subject;

(iii) being available for reasonable periods of time during semesters, study weeks and examination periods so that students may discuss aspects of the subject with them;

(iv) assessing students’ work fairly, objectively and consistently;

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(v) being available to students after marked material has been returned and after the final results have been released so that students may receive feedback.

(xii) OH&S INFORMATION FOR UNDERGRADUATE STUDENTS

Safety Rules 1. Shoes must be worn - no student with bare feet will be permitted entry. 2. Thongs or sandals are not acceptable footwear. Shoes must cover the whole foot. 3. Eating, drinking & smoking in the laboratory are forbidden. 4. Long hair must be tied back. 5. Safety glasses must be worn at all times in the proper manner. (OH&S Regulations) 6. Visitors are not permitted except with the consent of the class supervisor. 7. Any injury must be reported to the demonstrator & to the preparation room staff. 8. Do not run in any part of the Chemistry Department. 9. No bunsen burners should be used unless specifically advised by your Demonstrator. 10. All accidents must be reported to your demonstrator. 11. Any intoxicated student will be required to leave the laboratory immediately.

Safety Precautions

Notify your demonstrator or the laboratory staff of any incident, hazard, or equipment fault that you become aware of. 1. Fire must be avoided. Most volatile organic compounds are inflammable & must never be used near a naked flame. Check with your demonstrator before working with solvents. Organic solvents may only be evaporated on a steam bath, NEVER on a hot plate. 2. Fire on the bench can often be extinguished by smothering with a fire blanket or damp rag. Fire extinguishers are provided for use by trained personnel ie. your demonstrator or laboratory staff. 3. Clothing on fire Get the person on the floor as quickly as possible & roll them to smother the fire. (Damage to respiratory passages & eyes may be caused if the person remains standing.) Fire extinguishers must not be used to extinguish burning clothing. A safety shower or fire blanket should be used. 4. Chemicals may constitute a hazard because of toxic or corrosive action as well as from fire. In many instances use of a dispenser or pipette filler will be recommended. 5. Chemicals on the skin or in the eyes should be removed immediately by washing with plenty of water.

A safety shower & eye wash facility is located in the corridor near the rear entrance. 6. If a chemical is spilt on the bench or floor consult your demonstrator about safe ways of cleaning it up. For most spills mops supplied near the sink on the sidewall can be used. 7. Chemicals with toxic or irritating vapours will be placed in the fume hoods & must be kept there. Some chemicals that need special care in handling are: • strong acids & alkali (bases) • toxic chlorinated solvents such as dichloromethane • bromine, phenols, cyanides, sulphides • highly inflammable solvents such as acetone

Your Roles & Responsibilities

Like all other members of Australian society, you have the right to a healthy & safe environment in which to work & study. Monash University has an extensive network of Occupational Health & Safety staff & facilities, whose purpose is to ensure that this goal is achieved. As a student at Monash University you also have certain roles & responsibilities in

relation to Occupational Health & Safety, in order to help ensure your own health & safety & that of other students, staff, visitors to the University & the general public. These are summarised below.

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A wide range of OHS resources are available to you, some of which are listed below, but you are also encouraged to seek out further information in this area. If you have any questions or concerns about any of this, or about any other OHS issues, please discuss them with any of the contact people listed below.

General

You must read & understand this information sheet before undertaking practical classes.

The School’s academic & technical staff will give you specific information relating to the location of OHS facilities in your lecture theatre or laboratory, & details of practical classes in which you are involved. You must listen to this information & comply with any health & safety instructions.

If you are ever unsure of the correct & safe procedure to be followed, ask one of the technical or academic staff. Do not rely on information from other students.

Evacuations

You should ensure you are aware of the location of fire exits, extinguishers & other emergency facilities in your laboratory or lecture theatre, & also of the point at which you should assemble after an evacuation. This information is available from the academic or technical staff in the area, & is also clearly signposted.

You should also ensure that you are aware of the meaning of the different tones produced by the alarm system. Drills will be held regularly to assist with this.

You also have a responsibility to ensure that if an alarm sounds you behave sensibly & do not panic. This includes heeding the alarms tones, listening carefully to all voice announcements, complying with instructions from Wardens, & if necessary evacuating in a calm & orderly fashion & proceeding directly to the nominated assembly point.

First Aid

You should not provide first aid or emergency medical assistance to any person unless you have received formal First Aid training.

However, you should be aware that most of the technical staff in your laboratory have been formally trained in First Aid, Asthma Management, etc.

You should also make sure that you are aware of the location of all emergency facilities in your laboratory (eye-washers, safety showers, etc). The technical staff will give you this information at your first laboratory session, & facilities are also clearly signposted.

Incident Reporting

In the past, some people have been reluctant to report incidents because they feel they might get someone else into trouble. However, the reporting process is not intended for that purpose! We would prefer you to think of it as a way of letting us know about an issue so we can make sure it is resolved before anyone gets into trouble.

So, if you witness or are involved in an incident which you feel may have safety or security implications, whether or not anyone was hurt, please report it as soon as practical to one of the contact people listed below.

KEY OHS CONTACT PEOPLE

Any of the School’s technical staff. Dr Boujemaa Moubaraki, Safety Officer, Bldg 23 or 9905-4798) Dr. Craig Forsyth, Radiation Safety Officer, Bldg 23 or 9905-4588 Prof. Steven Langford, Head of School (room 100C or 9905-4569) Sharon Lockhart, Safety Officer, Faculty of Science (9905-1627)


INTRODUCTION

HEALTH AND SAFETY

This first session is, in part, designed to introduce you to the First Year Laboratories, the

location of your assigned locker and to acquaint you with the laboratory safety features

and rules. Before arriving at the laboratories you should have attended an introductory

lecture covering key aspects of the laboratory program and you will also have seen a

safety video. Once you have viewed this video you should complete the safety quiz below.

Please let the laboratory supervisor know if you have any medical condition or

special circumstances that could affect your ability to participate in the laboratory

program.

SAFETY QUIZ

Name Demonstrator _______ __

It is important to familiarise yourself with the layout and facilities of the First Year

laboratories. Read carefully through the Introduction to this Manual. Enter the following

information into your laboratory notebook. Have this copy signed off by your

demonstrator.

1. From your locker find a locker number near to:

Nearest fire extinguisher Nearest CO2 extinguisher

Nearest safety shower Nearest emergency exit ____ __

Nearest fire blanket

2. In which direction is the nearest exit? EAST/WEST

3. Where is the evacuation assembly area for the First Year Chemistry laboratories?

______

4. Where is the eye wash equipment? ________________________________________

5. For which types of fires should the water extinguisher NOT be used?

______

______

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6. What are some potential hazards in a Chemistry laboratory?

___________

7. What is considered essential safety practice in the Chemistry laboratory?

______

______

______

8. What should you do if you spill any chemical on your skin?

______

______

9. Name six things not to do in the laboratory

a. b.

c. d.

e. f.

10. Where is the Resource Centre located?

______

11. What constitutes plagiarism in the laboratory? What is the penalty?

______

______

Your Signature: ____________________________ Date: _____________

Demonstrators Signature: _______________________ Date: _____________

CHM1011 - Student Laboratory Manual Ex.1: Waters of Hydration: CuSO4.xH2O

EXERCISE 1

DETERMINING THE WATERS OF HYDRATION: CuSO4.xH2O

Learning Outcomes

1. Use both top-loading and analytical balances to determine the accurate mass of a solid

2. Use a volumetric flask to accurately prepare a primary standard solution

3. Calculate the number of moles of an unknown using mass and molecular weight

4. Calculate the concentration of a solution using the formulae: A = εcl and c = n/V

1. INTRODUCTION

1.1 The Volumetric (or standard) Flask

A volumetric (or standard) flask is a precision instrument for making up a measured mass

or volume of sample to an accurately determined volume of solution. The volumetric flask

must be clean but need not be dry. Measure into it, by pipette, the volume to be diluted or

add into it, using a short-stemmed funnel, the solid to be dissolved. When making the

level finally up to the mark, use a dropping (Pasteur) pipette for accurate control. Mix the

partially diluted solution thoroughly, by inverting and shaking, before making up to mark -

then mix again. The stopper must be fitted carefully and the contents mixed thoroughly.

1.2 Making a Standard Solution

A standard solution is a solution for which the exact concentration of a reagent is known.

To prepare a standard solution, a known mass of a reagent is weighed out (1), and then

quantitatively transferred to a volumetric flask. The flask is then filled approximately ¾

with the solvent (2) and is thoroughly mixed (3), to ensure all of the reagent has dissolved.

The flask is then filled to the mark with the solvent (4), and the contents mixed again.

TIP: What is a quantitative transfer? It means ensuring that the entirety of your reagent

is transferred from your weighing vial into your volumetric flask.

(1) (2) (3) (4)

Figure 1: Making up a standard solution. Note that all of the solid must be transferred into

the volumetric flask

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In Part A, you will determine the water content in a sample of hydrated copper(II) sulfate

(CuSO4.xH2O, where x is an unknown integer), a crystalline substance that contains water

molecules. You will heat a sample of hydrated copper(II) sulfate to remove the "water of

crystallization". The mass of water is found by weighing before and after heating. This

information is then used to find x in the chemical formula.

You will need to make use of the concept of moles:

n = m / MR (1)

Where n = number of moles, m = mass in grams (g) and MR = molecular mass, in g/moles

Example:

CuSO4 - there is 1 Copper atom, 1 Sulfur atom and 4 Oxygen atoms

Therefore, the molecular mass of CuSO4 = Cu + S + 4xO = 63.55 + 32.07 + 4x16.00

= 159.62 g/moles

= 159.6 g/moles (in 4 significant figures)

In Part B, you will learn to accurately make up a solution in a volumetric flask, using UV-

visible spectroscopy to determine your accuracy.

You will be using your dehydrated copper(II) sulphate from Part A to make up your solution

in a 25mL volumetric flask, and calculating the concentration of your final solution using

the following formula:

c = n / V (2)

where the concentration (c, in moles/L) is equal to the number of moles (n) divided by the

volume in litres (V).

You will then measure the absorbance of your solution and calculate the concentration

using the Beer-Lambert Law:

A = εcl (3)

which states that the absorbance (A) of a solution is directly related to its concentration (c).

The terms ε and l refer to the molar extinction coefficient and path length, respectively.

All measurements and observations should be recorded directly into your notebooks or

your proformas. Show all calculations.

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2. EXPERIMENTAL PROCEDURE

Work in pairs and record all experimental data into your proforma.

2.1 Part A: The Formula of Hydrated Copper(II) Sulfate

1. Weigh out approximately 0.5 g of hydrated copper(II) sulfate in a weighing bottle (use

the top-loading laboratory balance for this) and screw on the lid. Note the colour of the

solid.

2. Take the capped bottle into the analytical balance room and accurately measure its

mass. Record all your data directly onto your proforma.

3. Transfer the copper(II) sulfate from the bottle into a crucible.

4. Reweigh the empty bottle (with lid) using the analytical balance. This procedure

enables you to accurately determine the mass of copper (II) sulfate used for the

experiment whilst eliminating the risk of spilling chemicals in the analytical room. You

will need to use this technique every time a solid needs to be weighed accurately in

subsequent experiments, so ensure you familiarise yourself with it.

5. Heat the crucible using a hotplate until the powder has

turned a pale blue colour, but do not heat so strongly

that it starts to blacken. A setting of 3-4 for about 20-

25 minutes with occasional mixing should provide

sufficient heating. Note: - there should be a maximum

of 4 casseroles per hot plate (Figure 2, right).

6. Remove the crucible from the hot plate and allow it to

cool briefly (2-3 minutes) and transfer the contents back into the weighing bottle and

recap. Make sure that you transfer as much of the solid into the weighing bottle as you

can as this mass will have direct bearing on the quality of your calculations.

7. Accurately weigh the capped weighing bottle, which now contains dehydrated

copper(II) sulfate.

2.2 Part B: Accurately Making a Standard Solution

1. Quantitatively transfer the dehydrated copper(II) sulphate into a 25 cm3 standard

(volumetric) flask using a short stemmed funnel. Make sure all the solid is transferred

into the flask, or your calculations will be affected.

2. Add deionised water until the flask is approximately ¾ full, and mix the contents to

dissolve the copper(II) sulphate. Your demonstrator should demonstrate the right

technique.

3. When all the solid is dissolved, make the solution up to the mark with water, and

THOROUGHLY MIX the solution. Note the colour of the solution.

4. Using a Pasteur pipette, transfer some of this solution into a cuvette cell. Add enough

so the cuvette is about ¾ full.

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5. Measure the absorbance (A) of your copper(II) sulfate solution at 808 nm, using an

ultraviolet-visible spectrophotometer. Your demonstrator will check your solution and

direct you to the instrument bay.

6. CLEAN ALL LABORATORY EQUIPMENT USED. 3. REPORT

Use your results to perform the required calculations and complete your proforma. Ensure

you have performed all necessary calculations, completed any tables, and answered the

questions. Your finalised proforma is due in to your demonstrator by the end of the

session.

Hazard Identification and Risk Assessment Exercise 1: Determining the Waters of Hydration of Hydrated Copper(II) Sulfate

Identify the Hazard (the Potential to do harm)

Determine the Risk

Control the Risk (Preventing an incident)

Disposal of waste

Copper Sulfate (solid) HARMFUL Irritating to eyes & skin. Harmful if swallowed. Very toxic to aquatic organisms.

LOW Do not breathe dust (wipe crucible if dirty DO NOT blow residue out)

Hazardous waste container

Copper Sulfate (solution) HARMFUL Irritating to eyes & skin. Harmful if swallowed. Very toxic to aquatic organisms.

LOW (small amount of solid, dissolved in water)

Do not breathe dust (use care when transferring solid to volumetric flask, do not blow solid from funnel or off benchtop)

Hazardous waste carboy in fumehood

Hotplate Burns

LOW Take care not to touch when hot, use tongs to remove crucible. Do not touch the hot plate

N/A

General Glassware Cuts

MEDIUM Handle with care, dispose of chipped/cracked glassware, collect broken glass using dustpan and brush. If cut occurs, see your demonstrator, seek 1st aid.

Place broken glass in the labelled broken glass bin

Additional comments:

Once prac work has commenced it is essential that safety glasses are worn at all times in the laboratory

when anyone is using or handling chemicals. This includes walking through other labs whilst leaving the

building.

Students should always wash their hands before leaving the building.

CHM1011 - Student Laboratory Manual Ex.2: Analysis of Vitamin C in Ribena®

EXERCISE 2

QUANTITATIVE ANALYSIS OF VITAMIN C CONTAINED IN RIBENA®

Learning Outcomes

1. Perform a titration in duplicate, and back calculate to determine concentration

2. Work out reaction stoichiometry from a reaction equation

3. Identify the Vitamin C molecule

4. Understand the role of NBS (N-bromosuccinamide), acetic acid and the iodide/starch

indicator in the determination of Vitamin C

1. INTRODUCTION

1.1 The Pipette

The pipette is a precision instrument for convenient delivery of an accurately determined

volume of liquid (an aliquot). You should never hold a pipette by the bulb because the

heat of your hand can cause a significant change in the volume it will deliver. Always hold

it by the upper stem.

Rinse the pipette twice with the solution to be used in the measurements, then draw the

liquid by suction to just above the mark and adjust until the bottom of the meniscus just

touches the mark. Using the first finger to seal the top allows greater flexibility than using

the thumb and is especially helpful when using graduated pipettes. Slowly

rotate the pipette to allow accurate control of the level.

The standard procedure for draining a pipette is to hold it vertical with its tip in

contact with the side of the vessel for 15 seconds after the continuous discharge

has ceased. At the end of 15 seconds remove the pipette together with any

liquid in the tip. Always use a pipette filler!

1.2 Reagent Bottles

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

in your work. These must be clean but need not be dry. It is useful, at the outset, to

use a measuring cylinder with water to locate and mark the 50, 100 and 150 cm3

levels with a water-resistant pen. A reagent bottle should be rinsed twice, before

use, with small (2-5 cm3) amounts of the solution to be collected to ensure that the

concentration is not changed by any water in the bottle.

1.3 Ascorbic Acid (Vitamin C)

Ascorbic acid (1) is commonly known as vitamin C. It was one of the first vitamins that

played a role in establishing the relationship between disease and its prevention by proper

diet. That disease was scurvy, and was prone to sailors and explorers as far back as the

16th century. The prevention of scurvy was as simple as eating fresh vegetables and fruit

(which contain ascorbic acid). It is a powerful biological antioxidant (reducing agent) which

keeps the iron, in the enzyme prolyl hydroxylase, in the reduced form thereby maintaining

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its enzyme activity. This enzyme is essential for the synthesis of normal collagen. In

scurvy, the abnormal collagen causes skin lesions and broken blood vessels. Vitamin C

cannot be produced in the human body and must be obtained from the diet (citrus fruits,

tomatoes, etc.) or by taking synthetic vitamin C.

In this experiment, the amount of vitamin C is determined quantitatively by titrating the test

solution (Ribena®) with N-bromosuccinimide (NBS) (2). Vitamin C is oxidised to the

diketone (3) by NBS according to the following chemical reaction:

(1) (2) (3)

Figure 1: The oxidation of ascorbic acid via NBS.

What is the correct stoichiometry for this reaction?

It is important to know that although vitamin C is very stable when dry, it is readily

oxidised by air when left in solution. Therefore, the method you are about to use has

been developed because it works in the presence of a variety of other oxidisable

substances with one important exception. It will not work in the presence of sulfur dioxide

or sulphite (both of which are commonly used as preservatives of fruit juices) because they

also react with NBS. Hence, only fruit juices free from artificial preservatives will give

reliable results.

In this experiment, a solution of vitamin C is acidified and mixed with a strong solution of

potassium iodide (KI). A solution of an oxidising agent (NBS) is then added slowly. Once

the vitamin C (1) is oxidised the I- begins to oxidise to iodine, I2, which in turn combines

with other I- ions to form the I3- ion. The indicator in this reaction uses the I3

- ion with

starch, which form a dark-blue complex, thus indicating the oxidation of the ascorbic acid

is complete.

2. EQUIPMENT

2.1 Chemicals

Vitamin C (ascorbic acid, C6H8O6) 1M acetic acid (CH3COOH)

Starch solution Ribena®

4% potassium iodide (KI) N-bromosuccinimide (NBS, C4H4BrNO2)

16


3. EXPERIMENTAL PROCEDURE Work in pairs and record all experimental data into your proforma.

3.1 Stock Solution of Ascorbic Acid

1. Using the top-loading laboratory balance, weigh about 0.1 g of ascorbic acid (Mw =

176.1 g.mol-1) into a weighing bottle.

2. Weigh the bottle and contents accurately in the analytical balance room.

3. Weigh the bottle accurately again after tipping the contents into a short stemmed glass

funnel that has been placed in the neck of a 100 mL standard volumetric flask.

4. Calculate the difference in weight to check that the accurate mass of ascorbic acid

delivered is between 0.08 and 0.12 g.

5. Add 1M acetic acid through the funnel to help carry the powder into the flask and make

up to the mark with 1M acetic acid.

6. Calculate the expected concentration of stock ascorbic acid solution in mol / L and

label the flask.

3.2 Diluted Solution of Ascorbic Acid

7. Accurately pipette 10.00 mL of the standard solution into a second 100 mL standard

volumetric flask and dilute to the mark with 1M acetic acid solution.

8. Calculate the concentration (in mol /L) of the diluted ascorbic acid and mark the flask.

The N-bromosuccinimide (Mw = 177.98 g.mol-1) solution has been prepared for you by the

lab staff. The following titration allows for the NBS to be standardised, that is, to determine

its concentration accurately from the standard (diluted) ascorbic acid solution.

3.3 Standardisation of NBS

9. Pipette 10.00 mL of the diluted ascorbic acid solution into a conical flask.

10. Pour in 5 mL of 4% KI solution and 10 drops of starch solution.

11. Titrate with the stock NBS solution until a permanent mauve colour just appears.

12. Repeat until at least two concordant results are obtained.

The following table is in your proforma. Do not write results directly into your lab manual.

Titration number: Rough 1 2 3

Final burette reading/mL ±0.05

Initial burette reading/mL ±0.05

Titre / mL

13. Calculate the concentration NBS (mol/L)

3.4 Quantity of Vitamin C (Ascorbic Acid) in Ribena®

14. Using your 2 mL graduated pipette, deliver 2.0 mL of Ribena® into a conical flask.

15. Pour in 5 mL of 4% KI solution, add 10 drops of starch solution and titrate quickly with

standard NBS solution until a permanent blue/purple colour just appears.

17


18

16. Complete the second table in your proforma and repeat until two concordant results

are obtained.

Titration number: Rough 1 2 3

Final burette reading/mL ±0.05

Initial burette reading/mL ±0.05

Titre / mL

17. CLEAN ALL LABORATORY EQUIPMENT USED

18. Calculate the amount/mol of ascorbic acid in the 2.0 mL sample of Ribena®.

19. Calculate the mass of ascorbic acid in the 2.0 mL sample.

20. Calculate the concentration of ascorbic acid in the Ribena® as mg ascorbic acid per

100 mL of juice.

4. REPORT

Use the proforma to complete your report. Ensure you have performed all necessary

calculations, completed the two tables, and answered the questions on your proforma.

Hazard Identification and Risk Assessment Exercise 2: Analysis of Vitamin C in Ribena®


Determine the Risk (the Probability that harm may result)


Disposal of waste

Ascorbic Acid (solid) Irritating to eyes, respiratory system & skin.

LOW

Handle with care. Wash skin immediately with water if spill occurs.

Acidic residue carboy in fume cupboard

N-Bromosuccinimide Toxic if ingested or inhaled (solid)

Compound is in solution and is dilute. LOW risk

Handle with care. Wash skin immediately with water if spill occurs.


Potassium Iodide (4%) Irritating to eyes. Will stain skin.

Solution is dilute. LOW risk Will still stain skin at this concentration

Handle with care. Rinse quickly with water if spilt on skin to avoid staining.


Acetic Acid (1M) Harmful to eyes.

Solution is dilute. LOW risk

Wash skin immediately with water if spill occurs.


General glassware Cuts

MEDIUM Handle with care, dispose of chipped/ cracked glassware, collect broken glass using dustpan and brush. If cut occurs, see your demonstrator, seek 1st aid.

Labelled broken glass bin

CHM1011 - Student Laboratory Manual Ex.3: Geometric Models - VSEPR Method

EXERCISE 3

GEOMETRIC MODELS:

VALENCE-SHELL ELECTRON-PAIR REPULSION (VSEPR) METHOD

Learning Outcomes

1. Construct models using a model kit to build and aid visualisation of chemical structures.

2. Manipulate common diagram types and their conventions to represent a range of

inorganic molecules.

3. Visualise molecules in three dimensions to enhance understanding of shapes and

stereochemistry.

4. Generate molecular geometries, based on bond numbers and angles.

READING: Chapter 5 of Blackman; CHM1011 Lecture Notes.

1. INTRODUCTION

Historically, a number of models have been proposed for the purpose of describing

molecular shapes and structure. One of the most easy to apply of these is the Valence-

Shell Electron-Pair Repulsion (VSEPR) model. The basis of the VSEPR model is the

consideration of the repulsive forces between pairs of valence electrons, and it is assumed

that only the electrons in the valence shell influence the molecular geometry.

Consider the two systems in figure 1. The figure on the left represents the centre atom in

a small molecule with two regions of negative charge (electron domains) surrounding it.

VSEPR theory predicts these two regions of negative charge will adopt a position that

minimises the repulsion, that is, maximises the distance between them. The figure on the

right represents a similar case with three electron domains. These configurations

ultimately determine the geometry of the molecule.

In the following exercise you will be presented with a molecular model set which you can

use to predict and analyse the shapes adopted by different small molecules.

2. EXERCISE

Your demonstrator will provide you with a list of molecules and your task is to construct

these examples and complete a description of their structural properties. In your proforma,

use tables in the same format as the one below to answer the first section of the exercise.

19


2.1 Molecular Formula: X2Y 1. Lewis Structures (Including all possible resonance

forms)

2. Perspective Drawing

3. Number of atoms bonded

to the central atom

3

4. Number of non-bonding

electron pairs on the

central atom

1

5. Parent Geometry

TETRAHEDRAL

6. Molecular geometry with

ideal bond angles

TRIGONAL

PYRAMIDAL

7. Hybridisation of central

atom

sp3

8. Polarity

YES

Each of these criterion are addressed below with respect to the example of the phosphite

ion, PO33-.

1. Lewis Structures

Follow the six steps below to construct a Lewis structure for a molecule or ion. An in depth

explanation of how to construct Lewis structures is contained in your CHM1011 lecture

notes.

1. Arrange the atoms to show how they are connected – atoms are grouped around a

central atom, which is usually the least electronegative. Do not enter any bonds at this

stage!

2. Determine the total number of valence electrons in the molecule or ion.

3. Place a bonding pair of electrons between each pair of adjacent atoms to give a single

bond.

4. Begin with the terminal atoms and add enough electrons to each atom to give all of the

atoms an octet.

5. Place any electrons left over on the central atom.

6. If central atom has fewer electrons than an octet, use lone pairs from terminal atoms to

form multiple bonds to the central atom in order to achieve an octet.

20


Using the phosphite ion as an example:

Step 1: Position 3 oxygen atoms symmetrically about a phosphorus atom.

Step 2: The number of valence electrons is 1x5 + 3x6 + 3 = 26. (ie. # of P &

O valence electrons, plus the negative charge.)

Step 3: Three sigma bonds (representing 2 electrons each) are introduced

for connectivity between the atoms.

(Note: We still need to incorporate 26 - 6 = 20 electrons into the structure.)

Step 4: Introduce lone pairs to the terminal atoms (the three oxygen

atoms) such that they possess an octet of electrons.

(Note: We still need to incorporate 20 - 18 = 2 electrons into the structure.)

Step 5: Place the two remaining electrons on the centre atom

(phosphorus).

Step 6: Finally, consider what this arrangement of electrons means for the

charge on each atom. Compare the valency of each atom with the number

of electrons surrounding it. Determine if there is a localised charge on

each atom.

In some cases, two or more resonance structures may exist. E.g. The nitrate ion has three

equivalent resonance forms;

2. Perspective Drawing

The 3-D models will serve as a visual guide to help you with your perspective structures.

Use the following guidelines to draw them correctly:

Consider the molecule such that the centre atom and at least one terminal atom are in the

plane of a piece of paper.

(a) Bonds lying in the plane of the paper are represented with a regular solid line.

(b) Bonds projecting 'into' the paper away from you are represented with a hatched wedge-

shaped line.

21


(c) Bonds projecting 'out of' the paper towards you are represented with a solid wedge-

shaped line.

3. Number of Atoms Bonded to the Central Atom

The phosphite ion has three atoms bonded to the central atom.

4. Number of Non-bonding Electron Pairs on the Central Atom

The phosphite ion has one electron pair on the central atom.

5. Parent Geometry

The parent geometry corresponds to the number of 'electron domains' about the central

atom. An electron domain may be a single bond, double bond, triple bond or a lone pair.

The fundamental principle governing geometries in VSEPR theory is that electron domains

render positions such that they are as far away from each other as possible due to the

electrostatic repulsion. Names for the different parent geometries are given table 1.

E.g. In the phosphite ion example there is a total of 4 electron domains (3 single bonds

and one electron pair). Therefore, phosphite will possess a tetrahedral parent geometry.

Table 1. A summary of different parent geometries and molecular geometries as described

by VSEPR theory.

# of Electron

Domains about

Central Atom

Parent Geometry

(& Hybridisation)

# of Atoms

Bonded about

Central Atom

Molecular Geometry

2 Linear (sp) 2 Linear

3 Trigonal Planar 3 Trigonal Planar

(sp2) 2 Bent

4 Tetrahedral

3 Trigonal Pyramidal

4 Tetrahedral

(sp3)

2 Bent

5 Trigonal Bipyramidal

4 Disphenoidal (See-saw)

3 T-shaped

5 Trigonal Bipyramidal

(sp3d)

2 Linear

6 Octahedral

5 Square Pyramidal

4 Square Planar

6 Octahedral

(sp3d2)

2 Linear

22


6. Molecular geometry with ideal bond angles The molecular geometry is dependent on the number of atoms bonded to the central atom

AND the parent geometry. The molecular geometry may be different to the parent

geometry!

E.g. In the PO3- ion example there are a total of 3 atoms bonded to the central atom, while

the parent geometry is tetrahedral. Therefore, the molecular geometry is trigonal

pyramidal.

7. Hybridisation of central atom

The hybridisation of the central atom is related to the parent geometry. See table 1 for the

labels given to different types of hybridisation. Refer to your CHM1011 lecture notes for

additional information regarding hybridisation.

8. Polarity

The polarity of a molecule relates to the strength of its dipole – an asymmetric distribution

of charge resulting in positive and negative poles. The uneven distribution of charge in a

molecule is a result of the different electronegativities, P, of the constituent atoms. Table 2

below provides some values for P for main group elements. Originally established by

Linus Pauling, he defined electronegativity as "the power of an atom to attract electrons to

itself."

Table 2. Electronegativities for the main group elements of the periodic table.

Group

1

Group

2

Group

13

Group

14

Group

15

Group

16

Group

17

H

2.2

Li

1.0

Be

1.6

B

2.0

C

2.6

N

3.0

O

3.4

F

4.0

Na

0.9

Mg

1.3

Al(III)

1.6

Si

1.9

P

2.2

S

2.6

Cl

3.2

K

0.8

Ca

1.0

(d-block

elements)

Ga(III)

1.8

Ge(IV)

2.0

As(III)

2.2

Se

2.6

Br

3.0

Rb

0.8

Sr

0.9

In(III)

1.8

Sn(IV)

2.0

Sb

2.1

Te

2.1

I

2.7

Cs

0.8

Ba

0.9

Tl(III)

2.0

Pb(IV)

2.3

Bi

2.0

Po

2.0

At

2.2

23


24

Consider the molecule H-F. Two electrons form a sigma bond between the nuclei, but

what is the average position of these nuclei? Is it precisely the midpoint of the bond, or

closer to one of the two atoms? From the table below we

can see that H and F have P values of 2.2 and 4.0

respectively. Therefore, the two electrons in the sigma bond

experience a stronger attraction toward the F atom than the

H atom, creating an uneven distribution of charge. As a result, we can consider the H

atom as being slightly positively charged and the F atom as being slightly negatively

charged. This is often represented using either of the diagrams above.

Let us consider two other examples, the triatomics CO2 and H2O. Carbon dioxide has two

electron groups, both bonded, so its molecular geometry is linear. Oxygen is more

electronegative than carbon, so if we considered the C=O bond alone, the bonding

electrons would be asymmetrically distributed toward the oxygen atom. But the bond is

not alone; there are two C=O bonds in CO2, opposite one another: The pull toward the

oxygen atom on the left exactly balances the pull toward the oxygen atom on the right.

The polar bonds cancel one another. The molecule itself is nonpolar – it does not have a

dipole.

Water has four electron groups, with two bonded atoms, so it has a bent molecular

geometry. Oxygen is more electronegative than hydrogen, so an isolated O–H bond has

its bonding electrons distributed toward the oxygen atom. Now let’s consider the overall

molecule: The pull toward the oxygen atom is not

balanced by a pull in the opposite direction. In

fact, both pairs of bonding electrons will be

distributed toward the oxygen atom, creating a

small, localised charge on that end of the molecule.

3. REPORT

For each of the molecules you have been allocated by your demonstrator, be sure to

complete the table on your proforma. Ensure you have answered the questions on your

proforma.

CHM1011 - Student Laboratory Manual Ex.4: Transformations of Copper

EXERCISE 4

TRANSFORMATIONS OF COPPER

Learning Outcomes

1. Perform a separation by filtration.

2. Perform simple calculations involving molarities and stoichiometry.

3. Know that solubility rules can simplify the task of planning the conversion of one

compound to another.

4. Understand the difference between a physical and chemical change.

1. INTRODUCTION

In this exercise, you will isolate metallic copper, after starting with a small amount of

copper(II) chloride. In a four step process you will explore a number of different chemical

transformations and experimental techniques. You will also need to be conscious of some

important safety elements of the exercise, including the use of concentrated acids and

bases.

CuO

Cu(OH)2 CuSO4 CuCl2 Cu

2. EQUIPMENT

2.1 Chemicals

Copper(II) chloride 4M NaOH (sodium hydroxide),

6M H2SO4 (sulfuric acid), Zinc dust

5.5M HCl (hydrochloric acid)

2.2 Apparatus

Balance Small beakers pH paper

Steam bath Long stem filter funnel Filter paper

Glass rod Spatula Wooden peg

Pasteur pipettes.

2.3 Safety aspects

Most of the chemicals used are relatively safe; however, sulfuric acid, hydrochloric acid,

and sodium hydroxide are corrosive and may be dangerous if splashed into the eyes or

onto skin. If contact occurs flush immediately with water and in the case of contact with

eyes, seek medical advice. Spills can be diluted with water and mopped up with an

absorbent cloth, which should be rinsed with water and disposed of in the normal rubbish

bin.

Residues can be safely discarded down the normal drain with copious amounts of water.

25


3. EXPERIMENTAL PROCEDURE Work in pairs and record all experimental data into your proforma.

3.1 Part A: Copper(II) hydroxide from copper(II) chloride

Use the top loading balance to weigh out between 0.16g and 0.18g of CuCl2.2H2O into a

yellow-capped weighing vial. Close the lid of the vial and use the analytical balance to

weigh the vial+CuCl2.2H2O to the nearest 0.001g. Add the CuCl2.2H2O to a 100mL

beaker (tap the mixture into the beaker - do not use water!) and put the lid back on the vial.

Re-weigh the vial on the analytical balance and record the mass to 0.001g. Determine the

mass of CuCl2.2H2O added to the beaker. Add approximately 10 mL of deionised water to

the beaker and stir to dissolve. Write an ionic equation for this dissolution.

Carefully add, in small amounts (dropwise) with stirring about 1mL of 4M

NaOH (Please note: a full Pasteur pipette contains 1mL approx.). Any

solid which forms is copper(II) hydroxide, Cu(OH)2. Record your

observations and write a net ionic equation for the change that has

occurred.

To ensure the reaction is complete, there must be an excess of NaOH added to the

solution. The presence of excess OH- can be tested by using a pH paper which turns blue

in basic solutions (solutions with an excess of OH- ions). Test your solution using a glass

stirring rod and the pH paper; immerse the glass rod in the solution and touch a small

square of pH paper on a watch glass. If the solution is not basic (paper does not turn blue),

add more NaOH solution until it is basic.

3.2 Part B: Copper(II) oxide from copper(II) hydroxide

Heat the contents of the beaker on a steam bath for about 5 minutes, stirring occasionally.

Any chemical change that occurs is the removal of a water molecule from Cu(OH)2 to give

copper(II) oxide, CuO. Record your observations and write a net ionic equation for the

change.

3.3 Part C: Copper(II) sulfate solution from copper(II) oxide

After 5 minutes, carefully remove the beaker from the steam bath using a

wooden peg and let it cool down to room temperature. Filter the copper oxide

precipitate through a fluted filter paper (see Figure 1: How to fold filter paper

for your funnel, right) into a 100 mL conical flask using a long stem funnel.

Transfer all the solid to the filter paper using a wash bottle, rinse the beaker

with deionised water into the filter paper. Discard all the filtrate.

Use the analytical balance to weigh out a clean and dry 100 mL

beaker to the nearest 0.001 g. Record the mass. Place the beaker

under the filter funnel. Pierce a hole in the filter paper using a

spatula and wash the precipitate through the hole with 3 mL

(approx. three full Pasteur pipettes) of 6M sulfuric acid (H2SO4).

26


Wash through the remaining solid on the filter paper with the acid filtrate (use a Pasteur

pipette). When all the solid has dissolved, wash the filter paper with about 5mL water.

Record your observations and write a net ionic equation for the reaction between copper

oxide and sulfuric acid.

3.4 Part D: Copper from copper(II) sulfate solution

Use a watch glass and the laboratory balance to weight out approximately 0.3g of zinc

dust, Zn. Carefully add the zinc dust in very small portions (use a spatula) to the coloured

solution. Allow the beaker to stand, stirring occasionally until the solution becomes

colourless. If the solution is not completely colourless, add a little more Zn until it is so.

Record all your observations and write an ionic equation.

Wash the precipitate twice by decantation. To do this, allow the precipitate to settle and

decant (pour off) most of the supernatant solution (solution above the precipitate). Add

approximately 10 mL of water, swirl the contents, allow the precipitate to

settle and decant the supernatant. Repeat with a further 10 mL water.

Add 1 mL 5.5M hydrochloric acid (HCl) (a full Pasteur pipette approx.) to

the solid and stir. Record your observations. You may or may not observe

a reaction here. If you do, write an ionic equation. If you do not, explain

why HCl was added to the solid.

TIP: Zinc reacts with HCl to evolve hydrogen gas, H2, but copper does not.

Wash the precipitate with deionised water twice by decantation. Decant most of the water

from the beaker.

3.5 Part E: Calculation of Percentage Yield

Place the beaker containing the recovered copper metal on the steam bath and heat it until

it appears dry. Remove the beaker from the steam bath using a wooden peg and let it cool

down to room temperature. Using the analytical balance, weigh the dry beaker plus the

copper and record the mass. Complete the calculations on the results sheet. Submit your

copper sample to your demonstrator for inspection.

4. REPORT

Make sure you complete the Exercise 4 proforma, and include answers to all questions.

Staple any relevant lab notes to the back page of your report.

27


28

Hazard Identification and Risk Assessment Exercise 4: Transformations of Copper


Determine the Risk


Disposal of waste

Sodium Hydroxide (4M) CORROSIVE Causes burns, damaging to eyes

LOW risk Avoid contact with skin and wash immediately with water if spill occurs.

Acidic residue container in fume cupboard

Sulfuric Acid (6M) CORROSIVE Causes burns, damaging to eyes.

MEDIUM risk

Use in fumehood. Handle with care. Wash skin immediately under water if spill occurs.


Hydrochloric Acid (5.5M) CORROSIVE Causes burns, damaging to eyes.

MEDIUM risk

Use in fumehood. Handle with care. Wash skin immediately under water if spill occurs.


Copper II Chloride Damaging to eyes, can irritate skin. Harmful if ingested or inhaled

LOW risk

Avoid contact with skin and wash with water immediately if spill occurs.

Residue container in fume cupboard

Zinc Powder Irritating to eyes and respiratory system

LOW risk Avoid inhalation of solid and spilling onto skin. Wash skin with water if it comes into contact.

Put filter paper with zinc residue in the rubbish bin

Steambath Burns

LOW risk Do not place hands directly over steam or directly on metal surface. Use wooden peg to remove hot objects.

N/A


MEDIUM risk

Handle with care, dispose of chipped/ cracked glassware, collect broken glass using dustpan and brush. If cut occurs, see your demonstrator & seek 1st aid.


Once prac work has commenced it is essential that safety glasses are worn at all times in the laboratory

when anyone is using or handling chemicals. This includes walking through other labs whilst leaving the

building.

Students should always wash their hands before leaving the building.

CHM1011 - Student Laboratory Manual Ex.5:pH Measurement & Acid-base Titrations

29

EXERCISE 5

ACID-BASE TITRATION CURVES

Learning Outcomes

1. Construct and interpret an acid-base titration curve.

2. Experiment with a standard pH meter.

3. Distinguish weak and strong acids/bases

4. Distinguish mono- and polyprotic acids/bases.

READING: Blackman, Chemistry. John Wiley & Sons 2011, Chapter 9 & 11.

1. INTRODUCTION

1.1 The definition of pH

The acidity of an aqueous solution is related to the “activity” of hydronium ions in the

solution, which can be approximated by the [H3O+]. As a result of the wide range of values

of [H3O+] commonly encountered, (1 to 10-14 M), the acidity is usually measured using the

pH scale. The pH of a solution may be calculated by:

pH -log10[H3O+] (1)

(Exact calculations use activities instead of concentrations, but not always practical.)

Water always participates in a self-dissociation equilibrium:

H2O + H2O H3O+ + OH- (2)

Described by the well known self-dissociation equilibrium constant:

[H3O+][OH-] Kw = 1 x 10-14 at 25 oC.

In a neutral aqueous solution [H3O+] = [OH–] and at 25 oC the pH of a neutral solution is

7.00. Acidic solutions have lower values of pH (high values of [H3O+]), whereas basic

solutions have higher values of pH (low values of [H3O+]).

1.2 Strong Acids vs Weak Acids

Acids that dissociate completely into H3O+ ions when dissolved in water are strong acids:

HX + H2O H3O+(aq) + X–

(aq) (3)

They include hydrochloric acid, HCl, nitric acid, HNO3, perchloric acid, HClO4, and sulfuric

acid, H2SO4 (for 1 proton only).

A weak acid is one that coexists at equilibrium with its corresponding (or “conjugate”)

base, as for example, with acetic acid (HOAc): In this example, the acidity constant, Ka,

can be determined from the concentrations of acid and base in solution using equation (5).

HOAc + H2O OAc–(aq) + H3O+

(aq) (4)

(5) aK

[H3O+][OAc-] = [HOAc]

The “p” notation (as in pH) is based on the definition, pX = – log X, where p is an

abbreviation for power. So we have:

pKa = -log10Ka (6)


30

If we combine equations (5) and (6) we find that for any weak acid, HA, with its conjugate

base, A–:

(7) 1.3 Measuring pH with a pH Meter.

+ logapKpH = [A-]10 [HA]

pH can be measured using a glass electrode

and a suitable reference electrode. Frequently,

the glass electrode and the reference electrode

are housed in the same

casing as a

combination electrode

(see Figure 1). The

electrical potential of

the glass electrode is

sensitive to changes in

pH, whilst the reference electrode remains at a fixed potential. The

two electrodes, together with the solution between them form an

electrical cell, and a voltmeter measures the difference between

the two electrodes and can be calibrated to display the result in pH

units (see Figure 2). Your demonstrator will explain the use of the

pH meter.

Figure 1 (Above): The pH meter – your demonstrator will explain how to use it.

Figure 2 (Left): Combination pH electrode

The combination pH electrode is fragile, so raise and lower it with care and do not use it

to stir the solution. The electrode must not be allowed to dry out, so keep it immersed in

distilled water between readings.

The readings of a pH meter are only valid if the meter has been calibrated with a buffer of

known pH appropriate to the range of measurement. Additionally, temperature has a

significant effect on pH measurements, so temperature should always be reported along

with pH.

Today we will use the pH meter to follow an acid-base titration. We will place an acid-filled

burette over a beaker containing the pH electrode and a solution of base. As acid is

titrated into the base, the pH decreases, slowly at first and then more quickly until after the

equivalence point when the rate of decrease begins to slow down again. The equivalence

point1 can be determined from a graph of pH against volume of acid added by finding the

1 Equivalence point - Point at which the number of moles of standard is equivalent to the number of moles of unknown.


31

point of maximum slope, which is the point at which the curve changes from convex to

concave (see Figure 3).

Figure 3: The titration curve above represents the titration of a weak acid with a strong

base. Notice the following points: (a) initially, pH ~ 3, (b) the equivalence point is

considered to be the point where the slope of the curve is greatest, (c) the volume

corresponding to half of the volume added at the equivalence indicates a point on the

curve where pH = pKa.

2. EXPERIMENTAL PROCEDURE

Work in pairs and record all experimental data into your proforma (see

point 5 below).

2.1 Part A: Titration of ~0.l M NaOH with ~0.l M HCl

1. Pipette 10 cm3 of NaOH into a 100 cm3 beaker and add 20 cm3

(measuring cylinder) of distilled water so that the electrodes are

immersed. Measure the pH and the temperature.

2. Fill the burette with 0.l mol dm-3 HCl (record the exact concentration of

the acid) and take an initial reading to ±0.01 cm3 (there is no need to start

exactly at 0.00!)

3. Arrange the burette, beaker, electrodes and pH meter so that titration can

be conducted while the electrodes remain in place (see Figure 4).

4. Add about 1 cm3 of acid from the burette. (Record the burette readings

to ±0.01 cm3). Swirl the solution thoroughly and measure the pH. Do

not remove the electrodes or STIR with them!


32

Experimental Setup

5. Plot the amount of acid added against the resulting pH. The reading should be plotted

as soon as it is measured. The graph scales should be pH 2 - 13 (y-axis), volume of

acid 0 - 20 cm3 (x-axis).

6. Repeat steps 4 and 5. As the pH begins to fall more steeply, the volume of acid per

addition should be reduced to about 0.1 cm3.2

7. Once you are well past the equivalence point (steepest point on the graph), use 1 cm3

aliquots again.

8. Mark the equivalence point on your graph. Read the pH of the solution at the

equivalence point (taken at the mid-point of the steepest part of the curve). Record the

volume of HCl required for equivalence.

9. Using the accurate concentration of HCl listed on the bottle, calculate the concentration

of NaOH used in the experiment.

2.2 Part B: Titration of ~0.1 M Na2CO3 with ~0.1 M HCl

The procedure is similar to that of the titration of NaOH with HCl. However, there are two

equivalence points, one after addition of approximately 10 cm3 of HCl, the other after

addition of about 20 cm3 of HCl. A total of 25 cm3 of HCl should be added altogether.

1. Pipette 10 cm3 of sodium carbonate solution into a clean beaker. Add 20 cm3

(measuring cylinder) of distilled water so that the electrodes are immersed. Measure

the pH and the temperature.

2. Titrate by addition of HCl from the burette (in a similar manner to Part A). Remember

there are two equivalence points in this titration. Plot the titration curve as you

proceed. The graph scales should be pH 2 - 13 (y-axis), volume of acid 0 - 30 cm3 (x-

axis).

3. Mark the two equivalence points on your graph.

4. Use the graph to determine the initial pH (solution of Na2CO3) and the pH of a solution

of sodium hydrogen carbonate (NaHCO3) at the first equivalence point.

5. Use the graph to determine the pH of a saturated solution of carbon dioxide (the 2nd

equivalence the point).

The first part of the titration can be described approximately by equation (8) and the

graph follows equation (9) where pKa2= -log10 (acid dissociation constant of HCO3-).

CO32–

(aq) + H3O+(aq) HCO3

– (aq) + H2O(l) (8)

(9)

10 + logapKpH = [CO3

2-][HCO3

-]

2 If you overshoot at this stage, repeat the titration. You do not need to add small portions slowly from the burette. Instead add an amount just less than the amount required to overshoot, and then add small amounts to precisely detect the end-point


33

6. Find the value of pH at the point where [HCO3-] = [CO3

2-], and hence determine an

approximate value of pKa for the hydrogen carbonate ion (pKa2 for carbonic acid.).

7. Between the first and second equivalence points the reaction can be described

approximately by equation (10) and the graph follows equation (11) where pKa1 = -

log10Ka1 (acid dissociation constant of H2CO3).

HCO3– + H3O+ H2CO3 + H2O (10)

(11) [H2CO3] pKa + log10pH =

[HCO3-]

8. Find the value of pH at the second halfway point (where [H2CO3] = [HCO3-]) and hence

determine an approximate value of pKa for carbonic acid (pKa1).

9. Compare the calculated mean of pKa1 and pKa2 with the observed pH at the first

equivalence point.

Clean-up

Thoroughly clean all your glassware with warm water. Make sure the pH electrode is

thoroughly rinsed and stored in a container of fresh deionised water.

Hazard Identification and Risk Assessment Exercise 5: pH Measurement and Acid-Base Titrations




Disposal of waste

Hydrochloric Acid (0.1M) CORROSIVE Causes burns, damaging to eyes.

Solution is dilute. LOW risk.

Handle with care. Wash skin immediately under water if spill occurs.

Corrosives residue carboy in fume cupboard

Sodium Hydroxide (0.1M) CORROSIVE Causes burns, damaging to eyes.




Sodium Carbonate (Na2CO3) (0.1M) Harmful if swallowed Irritating to eyes and skin


Wash skin immediately under water if spill occurs.





CHM1011 - Student Laboratory Manual Ex.6: QC - Baking Powder Analysis

34

EXERCISE 6

QUALITY CONTROL: CALORIMETRIC ANALYSIS OF BAKING POWDER

Learning Outcomes

1. Distinguish between exothermic and endothermic reactions

2. Investigate how calorimetry can be used to determine the enthalpy of a chemical

reaction

3. Devise a method utilising calorimetry to analyse a powder with an unknown composition

4. Manipulate data and determine unknown compositions using a calibration curve

READING: Blackman, Chemistry. John Wiley & Sons 2011, Chapter 8.

GUIDED INQUIRY PRAC

This exercise is a “Guided Inquiry Prac” which means that the whole procedure is NOT

written in the lab manual. For some parts of this experiment, you will have to design and

perform your own experiment.

You will be working in a group of 4 students. You need to assign roles within your

group and delegate the work that has to be done over the three hours.

Do not forget that at the end of the 3 hour session, you need to have finalised all

experimental work, plotted all graphs and answered all questions of the proforma.

1. INTRODUCTION

Baking powder or baking soda are common ingredients used for

cakes and other baked goods. Baking soda is pure sodium

bicarbonate (NaHCO3) whereas baking powder consists of

baking soda and one or more acid salts. These salts are usually

classified into two categories: the low-temperature acid salts

such as potassium tartrate or calcium dihydrogen phosphate and

the high-temperature acid salts such as sodium aluminium

sulphate or sodium aluminium phosphate. If the baking soda is

used instead, then an acidic ingredient, such as lemon juice

(which contains citric acid) or vinegar (which contains acetic

acid) must be added to the batter so that a

chemical reaction occurs.

C CCOO

OHHO

OOC KK

Potassium tartrate

HH

HOOC C

CH2COOH

OH

CH2COOH

Citric acid

CH3COOHAcetic acid

Figure 1: Chemical structure of molecules.

In this practical a mixture of Citric Acid (CA) and Sodium

Bicarbonate (SB) will be used to mimic the composition of baking

powder. Citric Acid (CA) is a triacid (pKa1 = 3.1, pKa2 = 4.7 and

pKa3 = 6.4). Sodium bicarbonate (NaHCO3) is an amphoteric

species: it can act as an acid and as a base (pKa (HCO3-/CO3

2-) =

10.4 and pKa (H2CO3/HCO3-) = 6.4).

Your goal is to determine the exact composition of Monash

University’s new Baking Powder.


35

2. BACKGROUND ABOUT CALORIMETRY

Calorimetry is used to quantify the heat released or absorbed by a chemical reaction or a

physical process, for example dissolution and phase transition. Chemical reactions that

raise the temperature (release thermal energy) are called exothermic reactions. Chemical

reactions that lower the temperature (absorb thermal energy) are called endothermic

reactions. Change in temperature due to chemical reactions can be formally described by

the change in enthalpy (H). The enthalpy H of a chemical species describes its internal

energy. If a reaction is exothermic, then the chemical species lose internal energy to its

surroundings (raising the temperature), and so H is negative. Conversely, if a reaction is

endothermic, the chemical species absorb energy from its surroundings (lowering the

temperature), and so H is positive. We will be investigating the thermodynamics of the

chemical reaction between citric acid and sodium bicarbonate.

The change in enthalpy (H) of a chemical reaction is related to the change in temperature

(T):

H = - m C T

where m is the total mass of solution (g) and C is the heat capacity (J.K-1.g-1).

Recording the evolution of the temperature of the system gives information about the

change in enthalpy of the system.

3. PART A: THE EXPERIMENT

Your demonstrator will show you a quick experiment. Record all observations and

determine the chemical reaction that occurs (Questions A1 to A3). Explain the

observations (Question A4).

4. PART B: CALIBRATION CURVE

In order to determine the composition of a mixture, scientists perform calibration curves

using mixtures of known compositions. Then they measure and plot the change in the

“significant parameter” (y axis) versus the composition (x axis). The choice of the

experimental data is very important as they have to cover the whole range of expected

compositions. Finally they perform the same experiment with the mixture of unknown

composition and using the calibration curve they deduce the composition of the mixture. In

the following example, the change in temperature (T) is plotted against the mass

percentage of solid A. Six mixtures

are prepared within the range 0-

50%. The unknown mixture has a

T of 5.0°C. Using the line of best

fit (line that matches most of the

experimental data), the mass

percentage of solid A can be

deduced and found as 25%. Figure 2: Example of a calibration curve.


36

In the experiment you will perform, the “significant parameter” will be T and it will be

plotted against the mass fraction of citric acid or the molar fraction of citric acid

. The mass fraction of Citric Acid (CA) is defined as the mass of citric acid over the

total mass: . The molar fraction of Citric Acid (CA) is defined as the number

of moles of Citric Acid over the total number of moles: . The molar ratio is

defined as the number of moles of Sodium Bicarbonate (SBC) over the number of moles of

Citric Acid (CA): .

In order to plot a “nice” and useful calibration curve, you need to choose the experimental

data with care. You need to determine the composition of each mixture you want to

investigate in order to cover the whole range of expected compositions. For each mixture

you will prepare, the total mass (mass of citric acid + mass of sodium bicarbonate)

has to be 2.00 g. The first composition you will investigate will be mCA = 0.87 g and mSB =

1.13 g. Then you are free to choose your own composition of mixtures and to investigate

as many mixtures as you wish. Explain your plan in the

proforma (Question B4).

4.1 Experimental Procedure

1. Prepare your calorimeter by nesting two foam cups

inside each other. Stand this arrangement in a

400 mL beaker to keep it from falling over.

2. Put 30 mL of water into the cup and place a lid on it.

3. Clean and label 3 weighing bottles (Citric Acid,

NaHCO3, Mixture).

4. Weigh out in the first container labelled “Citric acid”

roughly the desired amount of citric acid. Record

the exact mass.

5. Weigh out in the second container labelled

“NaHCO3” roughly the desired amount of sodium

bicarbonate. Record the exact mass.

6. Transfer both powders into the third container

labelled “mixture”. Be sure to transfer all powder

(you can tap the bottle to avoid powder sticking at the bottom).

Figure 3: Experimental setup

7. Insert the Pasco temperature sensor through a small hole in the lid of the cup. Get

ready to record data.

8. When everything is ready, click on the “Start” button on the program to begin data

recording. The temperature should be constant. If not, click on the “Stop” button, wait

for a minute and click again on the “Start” button.


37

9. After about 15 seconds have elapsed, lift the lid of the cup and quickly add the mixture

to the water. Ensure the lid is adequately put back on the cup (avoid gaps). Swirl the

beaker for 10-15 seconds. Make sure the temperature sensor is on the edge of the cup

while stirring so the temperature can be recorded at any time.

10. After overall 30 seconds have elapsed, leave the cup on the bench with the

temperature sensor in it. Let the computer record the temperature for 5 minutes for the

first experiment (for your subsequent experiments you can stop the experiment when

you want).

11. Click on the “Stop” button to end data recording.

12. Dispose of the reaction products in the sink. Rinse the cup with distilled water for use in

the next experiment.

13. Fill in the table given in the proforma.

14. Repeat from step 2 with different compositions of citric acid and sodium bicarbonate.

15. On the graph paper provided, plot:

- T (y-axis) versus the mass fraction of citric acid (x-axis),

- T (y-axis) versus the molar fraction of citric acid (x-axis),

- determine the line of best fit for both graphs.

5. PART C: DETERMINATION OF THE COMPOSITION OF THE BAKING POWDER

In the proforma, provide the method you plan to perform in order to determine the

composition of the Monash Baking Powder. The demonstrator must check and mark

this method before you start any further experimentation.

Then perform the experiment and determine the composition of the Monash Baking

Powder. If necessary you can perform further tests in order to get the composition. You will

have to explain these tests in your report.

Once you have recorded all the information needed, go to ‘Experiment’ and select ‘Delete

All Data Runs’ from the list.

Hazard Identification and Risk Assessment Exercise 6


Determine the Risk


Disposal of waste

Citric Acid Irritating to eyes, skin

LOW Wash with water if in contact with the skin Down the sink with plenty of water

Sodium Carbonate May cause eye burns Harmful if swallowed. Irritating to skin

LOW Wash skin immediately under water if spill occurs.

Down the sink with plenty of water


MEDIUM Handle with care, dispose of chipped /cracked glassware, collect broken glass using dustpan & brush. If cut occurs, see your demonstrator, seek 1st aid

Place broken glass in the labelled broken glass bin

CHM1011 - Student Laboratory Manual Ex.7: Timing the Iodine Clock

38

EXERCISE 7: Timing the Iodine Clock

Learning Outcomes

1. Explore the idea of “rate order”.

2. Design a simple experiment to determine the overall rate order of a chemical reaction.

3. Investigate the effects of experimental conditions on rate, such as concentration and

temperature.

4. Work in teams, delegate tasks, and optimise your time management and efficiency.

READING: Blackman, Chemistry. John Wiley & Sons 2011, Chapter 15.

1. INTRODUCTION

1.1 The Iodine Clock Reaction

The following chemical equation describes the so-called “Iodine Clock Reaction”.

H2O2(aq) + 2I-(aq) + 2H3O+(aq) I2(aq) + 4H2O(l) (1)

Practically, hydrogen peroxide is added to a mixture of potassium iodide and sulfuric acid.

This process occurs faster or slower depending on the concentration of each reactant. The

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

1.2 Measuring the Rate: This is done by means of the well-known gentle reaction

between iodine (I2) and a (relatively) small amount of thiosulfate ion (S2O32–) to form the

tetrathionate ion (S4O62–) (2)

I2(aq) + 2S2O32-

(aq) S4O62- (aq) + 2I-(aq) (2)

When hydrogen peroxide is added to the mixture, the iodine formed by reaction (1) reacts

instantly with thiosulfate by reaction (2), until all the thiosulfate is used up. The ‘excess’

iodine then begins to accumulate in solution and (in the presence of starch) the colourless

solution suddenly turns blue. (Iodine forms a complex with starch and this complex has a

brilliant blue colour). This kind of reaction is widely known as a 'clock reaction'.

1.3 Determining the Rate Order: During today’s experiment, the aim is to determine the

rate order of the chemical reaction. This is not covered in lectures until week 11, so today

we will keep this pretty simple.

Zero Order Kinetics occurs when the concentration of a reagent [A] is increased, and

the rate of the reaction doesn’t change! Rate ~ [A]0

First Order Kinetics occurs when the rate of the reaction increases proportionally with

concentration. Rate ~ [A]1. That is, if the concentration of one reagent is doubled,

the rate of the reaction doubles. If the concentration is tripled, the rate triples, and so

on …

Second Order Kinetics occurs when the rate of the reaction increases proportionally

with concentration. Rate ~ [A]2. That is, if the concentration of one reagent is

doubled, the rate of the reaction quadruples (22). If the concentration is tripled, the rate

increases by a factor of nine (32), and so on …


39

2. TODAY’S EXPERIMENT

2.1 The Aim: The aim of today’s experiment to determine the rate law for the iodine clock

experiment. Since there are three reagents in this chemical reaction, the rate law has the

general form:

rate = [H2O2]x[I–]y[H3O+]z (3)

You will need to determine x, y and z.

Some of the questions you should consider might be:

How does the rate change as [H2O2], [I-], and [H3O+] are increased?

(Can you test these simultaneously, or do you have to change only one

concentration at a time?)

How does the rate change as [S2O32-] is increased?

How does the rate change as temperature changes?

TIP: the success of the experiment depends on there being a large excess of the reactants

H2O2, I- and H3O+ for reaction (1), relative to the thiosulfate used for reaction (2).

2.2 The Demonstration: Your demonstrator is going to show you an example of how this

experiment works. Watch closely - you are going to have to design the rest of this

experiment yourself! Here is a basic description of the procedure your demonstrator is

following:

The following materials are added to a 250 cm3 conical flask, in the order shown:

1. 35 cm3 0.36 M H2SO4 (measuring cylinder)

2. 43 cm3 distilled water (measuring cylinder)

3. 10 cm3 0.025 M KI (pipette)

4. 10 cm3 0.0025 M thiosulfate (pipette)

5. 1 cm3 starch solution (measuring cylinder)

At this point, a stop-watch or timing should be ready!

6. 1 cm3 of 0.8 M hydrogen peroxide (FUMEHOOD! USING THE 2 cm3 PIPETTE

PROVIDED.

Time is measured from the moment the peroxide is added, and the mixture is swirled

gently 3 or 4 times. The solution is left to stand until a blue colour appears, due to an I3-

:starch complex. Determine the time elapsed (in seconds). There are a few important

points you should consider.

Is temperature important. If so, can you measure it?

When performing this experiment, it should be completed at least twice to verify

the timing.

When following the structions above, what is the total volume of the mixture?

Do you think this is important?


40

2.3 Inquiry Design Explore Answer

Once your demonstrator has completed the demonstration, you will complete the

experiment in your group. It is time to consider how you might design an experiment to

determine the overall rate order for the experiment.

Complete this experiment using the proformas provided.

The marking scheme can be found on your proforma.

All team members will receive the same mark /20.

There is a lot to do, so delegate your jobs accordingly!

Identification and Risk Assessment Exercise 7: Rate Law of an Iodine Clock




Disposal of waste

Sulfuric acid (0.36M) CORROSIVE Causes burns, damaging to eyes




Potassium Iodide (0.025M) Irritating to eyes Can stain skin

Solution is dilute. LOW risk Will still stain skin at this concentration

Handle with care. Rinse quickly with water if spilt on skin to avoid staining.


Thiosulfate (0.0025M) Irritating to eyes, & skin

LOW Avoid contact with skin, wash immediately with water if spill occurs.


Hydrogen Peroxide (0.8M) Harmful if swallowed Irritating to eyes and skin

LOW Avoid contact with skin as it can have a bleaching effect, wash immediately with water if spill occurs.





CHM1011 - Student Laboratory Manual Ex.8: The TEK Design Prac

EXERCISE 8: The TEK Experiments

Learning Outcomes

1. Explore a number of interesting but simple chemical experiments.

2. Research the underlying elements of thermodynamics/equilibria/kinetics in order to

explain the observed chemistry

3. Use a multimedia presentation (poster or vodcast) to explain the chemistry to your

peers.

READING: Blackman, Chemistry. Chapters 8, 9, 10, 11 & 15.

1. INTRODUCTION

This week’s activity is a little different. Our demonstrators are going to give you an

opportunity to sit back and focus on using your observation skills, although there will be

some opportunity to participate.

We have organised a ‘round-robin’ collection of short demonstrations for you to observe –

simple chemical experiments with some interesting underlying thermodynamics, equilibria

or kinetics, and in some cases all of the above.

Make sure you make note of the reagents used in the different reactions you will observe

in this weeks activities. Ask the demonstrator to clarify anything along the way if you are

unsure of the chemicals or equipment being used to demonstrate a reaction.

2. THE DEMONSTRATIONS

Here is the shortlist of some of the activities – have a think about which subtopic each one

explores. You will be provided with a proforma to take notes throughout the experiments.

At the end of the show, there will be a chance to break away into groups of 2 or 3 to start

planning.

Experiment 1: Elephants Toothpaste Experiment 4: The Rainbow Clock

Experiment 2: Fire without Matches Experiment 5: Colourful Cobalt

Experiment 3: Oscillating Reactions Experiment 6: Pharoah’s Serpent

Some questions to consider:

What did you observe (use all of your senses where appropriate)?

What is the chemical reaction?

What principles are being explored with this specific reaction?

What chemicals are involved?

What factors affect the rate of the reaction?

Is there a significant change in energy content of the products versus reactants?

Can we speed up/slow down the reaction?

Why do these chemicals react the way they do?

What are the components of the equilibrium/oscillation reaction?

41


3. PRESENTATIONS

Instead of the usual lab report this week, you will prepare either a poster (easy) or a

vodcast (more challenging) to present in Week 12. You will have access to the big screens

so come prepared with your presentation USB key. DO NOT USE LINKS TO YOUTUBE

CLIPS, but feel to embed footage you have recorded yourself in the lab. You DO NOT

need to print out your poster. Poster presentations should be accompanied by a 5 minute

oral presentation. Be prepared to answer questions from your classmates and tutors.

If you chose to do a Powerpoint presentation, it should contain only a single slide with the

relevant information on it. Every individual will need to present some aspect of the

chemical reaction being discussed. You will be assessed on the slide presentation as a

group so make sure everyone understands the chemical reaction discussed, and presents

part of the slide as well as being able to answer questions relating to the reactions on the

slide.

Take Photos & Record Footage! Don’t be afraid to use a camera during the show so that

you can add to the visual appeal of your presentation! Make it interesting so that others

want to read your poster and learn the material you have learnt. Remember, it is not

important to only learn interesting things but also make interesting the things that are to be

learnt.

The Assessment

Of course this exercise is all about exploring the chemistry, so your presentation will be

assessed with this in mind. Your demonstrator will be marking your posters. You will also

be judged by your peers.

Prizes! Every presentation will also be judged by your classmates. A prize will be given

out to the popular favourite!

Groupwork

Since you are working in groups, make sure you have swapped details (email, phone

number) so that you can collaborate over the following week. Everyone is expected to

know the material presented by their group.

Poster Template

We have created a Powerpoint template for you to download from MOODLE. Make sure

you use this template, as it has to right font size pre-set to look good via our projector

system in the First Year Labs. The poster should explain the topic in a clear, coherent and

correct way with text and graphics. They key idea of a poster is to communicate

information visually to other people even if you are not present to talk about it. However,

you should be able to discuss specific aspects covered in more detail if requested.

42


43

Oral Presentation

Presentations are the best way of communicating your research to a large group of people

quickly. You will be give 3-5 minutes as a group to present the reaction you have

observed. This type of presentation also gives you the opportunity to receive instant

feedback from your group leaders and peers and will allow others to assess how well you

have understood the material you have presented. Your presentation can include models,

demonstrations, samples or anything that will help communicate the key points to the

audience (but no hyperlinks to online material). The questions at the end of the

presentation (from your peers or tutor leader) may be directed at any member of the group

so ensure you understand the chemistry behind not only what you prepared but also the

information prepared and presented by any member of your group.

Presentation Marking scheme (/20)

Chemistry Content - accurate information

- explanation of principles

2

3

Organisation - clear theme

- logical flow of information

- relevant visual images

- overall layout and impression

2

1

1

2

Visual appeal of poster/Powerpoint (clear, easy to read, relevant to

theme)

2

Group work – equal participation from all members

- all members able to answer questions

2

2

Involvement and attentiveness to other presentations 3

CHM1011 - Student Laboratory Manual Periodic Table

44

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1011 Lab Manual 2013 Print