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chapter19
#Chapter 19: Preparation of Solutions Without a Balance

The procedure in the section on Relative Standardization (p. 66) allows us to do something seemingly impossible – prepare solutions for volumetric analysis that allow students to get perfect results without using either a balance or volumetric glassware in the preparation. All that you have to do is make two solutions that are close, and then use several cycles of relative standardization to prefect the molarity ratio.

To measure volume, we can use marks on plastic water bottles as described in the entry for volumetric glassware in the Sources of Laboratory Equipment (p. 208) section.

##19.1 To make 0.05 M sulphuric acid (equivalent to 0.1 M HCl) for fifty students

1. Put 9.9 liters of water into a bucket.
2. Add 110 mL of battery acid. This may be accomplished easily by filling a 10 mL plastic syringe eleven times.

##19.2 To make 0.033 M citric acid (equivalent to 0.1 M HCl) for fifty students

1. Put 10 liters of water into a bucket.
2. Add 64 g of citric acid. In the absence of a balance, one can often have 1/8 of a kilogram (125 g) measured in the market. Dissolve this in 20 L of water to produce a 0.033 M solution.

##19.3 To make 0.1 M sodium hydroxide for fifty students

1. Put 10 liters of water into a bucket.
2. Add 40 g of caustic soda. In the absence of a balance, use a plastic syringe to find the volume of a plastic spoon. Fill the spoon with caustic soda and use it to add a total of 19 cm3 or mL caustic soda knowing the volume of each spoonful. Please read the safety note in Dangerous Chemicals (p. 37).

##19.4 To make 0.1 M sodium hydrogen carbonate for fifty students

1. Put 10 liters of water into a bucket.
2. Add 84 g of bicarbonate of soda. In the absence of a balance, find the volume of a spoon as above and add 39 cm3 or mL of bicarbonate of soda. Alternately, if 8.33 liters of solution is sufficient, measure this volume of water and then add one whole box of bicarbonate of soda. A box is 70 g.
chapter20
#Chapter 20: Relative Standardization

Preparing large volumes of solution is difficult with great accuracy. Relative standardization is a technique to correct the concentration of solutions so that they give the correct results for practical exercises. Note that this technique is only useful in educational situations where the purpose is to prepare a pair of solutions for titration that give an answer known by the teacher. In scientific research, the aforementioned technique – absolute standardization – is used because the concentration of one of the solutions is truly unknown.

All schools should use relative standardization to check the concentration of the solutions they prepare for the national examinations. This ensures that the tests measure the ability of the students to perform the practical, and not the quality of the school’s balance, water supply, glassware, etc. While useful for all schools, relative standardization is particularly helpful for schools with few resources, as it allows the preparation of high quality solutions with extremely low cost apparatus and chemicals.

##20.1 General Theory

The principle of a titration is that the chemical in the burette is added until it exactly neutralizes the chemical in the flask. If the two chemicals react 1:1, e.g.

$$\mathrm{HCl}_{(aq)} + \mathrm{NaOH}_{(aq)} \longleftarrow \mathrm{NaCl}_{(aq)} + \mathrm{H}_{2}\mathrm{O}_{(l)}$$

then exactly one mole of the burette chemical is required to neutralize one mole of the chemical in the flask. If the two chemicals react 2:1, e.g.

$$2\mathrm{HCl}_{(aq)} + \mathrm{Na}_{2}\mathrm{CO}_{3(aq)} \longleftarrow 2\mathrm{NaCl}_{(aq)} + \mathrm{H}_{2}\mathrm{O}_{(l)} + \mathrm{CO}_{2(g)}$$

then exactly two moles of the burette chemical is required to neutralize one mole of the chemical in the flask. Let us think of this reaction as a mole ratio.

$$\frac{\mbox{moles of }A}{\mbox{moles of }B} = \frac{n_{A}}{n_{B}}$$

Where $$n_{A}$$ and $$n_{B}$$ are the stoichiometric coefficients of A and B respectively.

moles = molarity × volume = M × V (so long as V is measured in liters)

$$\mathrm{moles} = \mathrm{molarity} \times \mathrm{volume} = M \times V \mbox{ (so long as V is measured in liters)}$$

By substitution,

$$\frac{(M_{A})(V_{A})}{(M_{B})(V_{B})} = \frac{n_{A}}{n_{B}}$$

A student performing a titration might rearrange this equation to get

$$M_{A} = \frac{(n_{a})(M_{B})(V_{B})}{(n_{B})(V_{A})}$$

or

$$M_{B} = \frac{(n_{B})(M_{A})(V_{A})}{(n_{A})(V_{B})}$$

As teachers, however, we care with something else: making sure that our students find the required volume in the burette. Solving the equation for $$V_{A}$$ we find that

$$V_{A} = \frac{(n_{a})(M_{B})(V_{B})}{(n_{B})(M_{A})}$$

As $$n_{A}$$ and $$n_{B}$$ are both set by the reaction, as long as we use the correct chemicals there is no problem here.

$$V_{B}$$ is measured by the students – it is the volume they transfer into the flask. As long as the students know how to use plastic syringes accurately, they should get this value almost perfectly correct.
The remaining term, $$M_{B}$$ is for the teacher, not the student, to make correct. If we prepare the $$M_{A}$$ solutions poorly, our students can do everything right but still get the wrong value for $$V_{A}$$. It is very important that we ensure that our solutions have the correct ratio of $$M_{B}$$ so that the exercise properly $$M_{A}$$ assesses the ability of our students.
Many people look at this ratio and decide that they therefore need to prepare both solutions perfectly, so that $$M_{B}$$ and $$M_{A}$$ are exactly what is required. This not true. The actual values for $$M_{B}$$ and $$M_{A}$$ are not important; what matters is the ratio $$M_{B}$$ to $$M_{A}$$!

For example, if the titration requires 0.10 M HCl and 0.10 M NaOH, our expected mole ratio is:

$$\frac{M_{HCl}}{M_{NaOH}} = \frac{0.10}{0.10} = 1$$

Preparing 0.11 M HCl and 0.09 M NaOH will cause the students to get the wrong answer:

$$\frac{M_{HCl}}{M_{NaOH}} = \frac{0.11}{0.09} = 1.22$$

However, preparing exactly 0.05 M HCl and 0.05 M NaOH results in the same molar ratio:

$$\frac{M_{HCl}}{M_{NaOH}} = \frac{0.05}{0.05} = 1$$

Thus the students can get exactly the right answer if they use the right technique even though neither solution was actually the correct concentration. How can we ensure that we have the correct molar ratio between our solutions? Titrate your solutions against each other. If the volume is not the expected value, one of your solutions is too concentrated relative to the other. You can calculate exactly how much too concentrated and add the exact amount of water necessary to perfect the ratio. This process is called relative standardization, because you are standardizing one solution relative to the other.

##20.2 Procedure for Relative Standardization

In some titrations the acid is in the burette and in some it is the base is in the burette. So let us not use “acid” and “base” to refer to the solutions, but rather “solution 1” and “solution 2” where solution 1 is the solution measured in the burette and solution 2 is measured by pipette (syringe). You should have prepared a bucket or so of each. The volume you have prepared is $$V_{1}$$ liters of solution 1 and $$V_{2}$$ liters of solution 2.

Titrate the solutions against each other. Call the volume you measure in the burette “actual titration volume” You know the desired molarity of each solution, so from the above student equations you can calculate the burette volume you expect, which you might call “theoretical titration volume.” After the titration, there are three possibilities. If the actual titration volume equals the theoretical titration volume, your solutions are perfect. Well done. If the actual titration volume is smaller than the theoretical titration volume, solution 1 is too concentrated and must be diluted. Use the ratio:

$$\frac{V_{1} \mbox{ (before dilution)}}{V_{1} \mbox{ (after dilution)}} = \frac{\mbox{actual titration volume}}{\mbox{theoretical titration volume}}$$

If the actual titration volume is larger than the theoretical titration volume, solution 2 is too concentrated and must be diluted. Use the ratio:

$$\frac{V_{2} \mbox{ (before dilution)}}{V_{2} \mbox{ (after dilution)}} = \frac{\mbox{theoretical titration volume}}{\mbox{actual titration volume}}$$

After diluting one of your solutions, repeat the process. After a few cycles, the solutions should be perfect. Remember that the volume “before dilution” is the volume actually in the bucket, so the amount you made less the amount used for these test titrations.
chapter21
#Chapter 21: Preservation of Specimens

See the Shika Express Biology manual for information regarding collection and preservation of particular specimens covered in the O-level syllabus.

- Mosses and lichens: Wrap in paper or keep in a closed container.
- Plants and parts thereof: hang in the sun until dry. Alternatively, press the plants using absorbent material and a stack of books.
- Insects: Leave exposed to air but out of reach by other insects until bacteria eat everything except the exoskeleton. If you want to preserve the soft tissue, store under methylated spirits.
- Fish, worms, amphibians, and reptiles: Store in methylated spirits (will makes specimens brittle) or a 10% formaldehyde solution (more poisonous and more expensive).
- Parts of mammals (e.g. pig eyes, bovine reproductive organs): store in 10% formaldehyde solution.

##21.2 Skeletons
Skin the animal and remove as much meat as possible. Bury the bones for several months. Exhume and assemble with wire and superglue.

##21.3 Living Specimens

Be creative! Figure out what the animal will eat, who will feed it, what it will drink, where it can hide, how it can be observed, etc.
chapter22
#Chapter 22: Dissection

See the Shika Express Biology manual for information regarding dissection of particular specimens covered in the O-level syllabus.

##22.1 Preparation of Specimens

Unless you want students to observe a beating heart, dead specimens are much easier to work with than unconscious ones. This also removes the problem of stunned animals waking up in the middle of their dissection.

- Flowers and other plant parts: No preparation required as long as the samples are relatively fresh. Store samples in closed plastic bags to minimize drying. If you intend to keep them for more than a day or two, submerge the bags in cold water to slow the rate of molding.
- Insects: Kill with household aerosol insecticide. Use specimens within one day of collection, unless you have refrigeration or freezer.
- Fish: Keep living until the day of the dissection. Then remove from water until they suffocate. Use immediately after death.
- Frogs: Able to breathe above and below water, frogs are hard to starve of oxygen. One option is to seal them in a container of methylated spirits and then rinse the dead bodies with water prior to dissection.
- Reptiles, birds, and mammals: For most organ systems, you can kill the animal by blunt trauma without ruining the lesson. Students can even bring animals caught and killed in homes. Snakes should be decapitated along with enough of the body to remove the fangs and venom sacks. Bury these deeply. Do not use animals killed by poison, or those that were found dead. For completely undamaged specimens, enclose the live animal in a cage (or a tin with adequate holes) and submerge in a bucket of water until drowned.
- Living specimens: If you really want to see that heart beating, use chloroform. This can be transferred from bottle to specimen jar via cotton ball, or perhaps made in situ by the reaction between propanone (acteone) and bleach. We have not yet attempted the latter – if you do, remember that the products are poisonous gases; indeed, that is the point. Note that if you use too little chloroform, the animal will feel the blade opening it up. If you use way too little, it may start squirming. If you use too much chloroform, however, you will simply kill the animal – you might as well have drowned it.

##22.2 Tools

For more, see the section on Sources of Laboratory Equipment (p. 208).

- Scalpels can be made using razor blades and tongue depressors. Make sure the razors are very sharp. If the blade is dull or floppy, the students will probably push too hard, and may cut themselves when the skin finally gives and the blade slips.
- Optical Pins from new disposable syringes are an easy option.
- Dissection trays can be prepared using cardboard or by making a 1 cm thick layer of wax on the bottom of a shallow tray or bowl. This surface will readily accept pins and is easy to clean.

##22.3 Procedure

This varies by species. The internet has many resources and there are many good books with very detailed instructions – alas, this manual is not yet one of them. A crude method follows:

1. Position the specimen on its back and make a clean, symmetric, and shallow incision down the full length of the underside.
2. Make additional perpendicular cuts at the top and bottom of the torso for an overall “I” shape. These cuts should only just penetrate the body cavity.
3. Open up skin “door” you have created, pinning them back onto the dissection tray.
4. Pick an organ system – circulation, digestion, nervous, etc – and, with the aid perhaps of a good drawing, remove other material to focus on the target anatomy.

You can teach many systems from one specimen – start with the most ventral (front) and move to the most dorsal (back).
Encourage students to sketch at various steps in the process. Also encourage them to identify anatomy for themselves, perhaps with the aid of thought provoking questions and discussion in groups.

##22.4 Cleanup and Carcass Disposal

Wash all blades, pins, and trays with soapy water. Rinse all tools to remove the soap and then soak for about fifteen minutes in bleach water. When finished, rinse again in ordinary water.
Bury all carcasses in a deep pit, below the reach of dogs. You may also add kerosene and burn, but this smells bad and costs money.
chapter23
#Chapter 23: Preparation of Culture Media

##23.1 Introduction

In microbiology, there are two basic types of media: solid agar media and a liquid broth media. From these, many types of media can be made. Generally, exact amounts of ingredients are not needed so if you want to make some agar plates or liquid cultures try with the resources you have. The recipes listed are a guideline to help you get started.

##23.2 Media Recipes

###23.2.1 Basic Agar (1.5%)
- 15 g/L agar, like gelatin or if you can find seaweed you can grind it up
- 10 g/L nutrient source, e.g.sugar, starch (potatoes), beans fruits like mango and papaya
- 1-2 g/L salts and phosphates. This varies with what you want to grow — experiment! (table salt is usually fine)
- 1 L water

Add and mix all the ingredients together and heat until boiling. Boil for ~15 minutes and make sure all the gelatin/agar is dissolved. Pour liquid into Petri plates (15-20 mL each). The plates should solidify ~45 degrees C. Cover and keep agar side up in a cool place if possible. If the plates do not solidify, try adding more gelatin or corn starch to thicken it up. You can also pour agar into test tubes/syringes to do oxygen tests (aerobic vs. anaerobic)

###23.2.2 Blood Agar
- 15 g/L agar/gelatin/ground sea weed
- 10 g/L nutrient source
- 15 mL sheep’s blood (other organizisms also work)
- 1 L water

Heat and boil agar, nutrient source and water for 15 minutes. After liquid has ceded, around ~45 degrees C (when you can leave your hand on the flask for a few seconds), add in blood until the mixture is blood red. Swirl in and pour into plates.

###23.2.3 Liquid Broths

- 10 g nutrient source
- 1 L water
- 1-2 g salts/phosphates

Mix together, heat, and boil. Distribute in test tubes.

##23.3 Things you can do after media preparation

- Agar-streaked plates! Swab something (back of throat, nose, belly button, door handle, etc) and gently rub onto the agar. Try not to gouge the agar.
- You can also do experiments to test the effects of salt concentrations, temperature, and nutrient concentrations.
- After all the plates solidify, incubate them at around 25-30 degrees C. Ideally the temperature remains constant. Check the plates after 24 hours for growth.
- For liquid broths you can inoculate test tubes with a sample from the environment. Incubate and check like agar plates. If there is growth the liquid will be turbid instead of clear like a control tube with only broth.
- You can use liquid cultures for wet mounts under microscopes as samples for agar plates or to allow students to see the difference between growth and no growth.

##23.4 What to use if you do not have plates or test tubes

- Use old water bottles or old plastic packaging for plates
- Use anything rigid and heavy for covered, e.g. building tiles
- Sealed/closed plastic syringes for test tubes
- Try to keep materials as sterile as possible but do not worry if there is contamination. Use contamination as a learning experience. Penicillin was contamination and it became a wonder drug.

##23.5 Things to do once you have cultures

- Take a sample from agar plate and drop hydrogen peroxide on it. Does it bubble? (Yes, it has catalase)
- Extract DNA from E. coli.
- Fermentation = use a liquid broth with peptone, acid-base indicator like phenol red, and inverted tube to trap gas and 0.5 – 1.0% of carbohydrate you want to test. If fermentation occurs (phenol red), the broth will turn yellow and gas should be collected in the tube. If the tube remains red, you can test for glucose production by adding a few drops of methyl orange. If the pH is below 4.4, it will remain red. If the pH is above 6.0, it will turn yellow.

##23.6 Guide to Identifying Common Microorganisms

- Pseudomonas aeroginosa: is green and smells like grape jelly (can grow in disinfectant)
- Serratia marcescens: grows pink-red between 25-32 degrees C (will be white otherwise)
- Escherichia coli: pale white/yellow, smells like inole
- Proteus spp: swarm on plates and smell like urine and brownies
- Bacillus subtillis: pale beige, smells a bit sweet
- Vibrio cholera: smells like buttery popcorn
- Staph vs Strep: Staph is catalase (+), strep is (-)
chapter24
#Chapter 24: Using a Microscope

##24.1 Parts of a Microscope

- Eyepiece: or ocular lens is what you look through at the top of the microscope. Typically, the eyepiece has a magnification of 10x.
- Body Tube: tube that connects the eyepiece to the objectives
- Objective Lenses: primary lenses on the microscope (low, medium, high, oil immersion) which are used to greater magnify the object being observed. A low power lens for scanning the sample, a medium power lens for normal observation and a high power lens for detailed observation. Normal groups of lens magnifications may be [4×, 10×, 20×] for low magnification work and [10×, 40×, 100×] for high magnification work. Some microscopes also use oil immersion lenses and these must be used with immersion oil between the lens and the cover slip on the slide. Oil immersion allows for a much greater magnification than air and typically ranges from 40x-100x.
- Revolving Nosepiece: houses the objectives and can be rotated to select the desired magnification.
- Coarse Adjustment Knob: a large knob used for focusing the specimen
- Fine Adjustment Knob: small knob used to fine-tune the focus of the specimen after using the coarse adjustment knob.
- Stage: where the specimen to be viewed is placed
- Stage Clips: used hold the slide in place
- Aperture: hole in the stage that allows light through to reach the specimen
- Diaphragm: controls the amount of light reaching the specimen
- Light Source: is either a mirror used to reflect light onto the specimen or a controllable light source such as a halogen lamp

##24.2 How to Use a Microscope
- Always carry a microscope with two hands! One on the arm and one on the base!
- Plug the microscope into an electrical source and turn on
- Make sure the stage is lowered and the lowest power objective lens is in place
- Place the slide under the stage clips with specimen above the aperture
- Look through the eyepiece and use the coarse adjustment knob to bring the specimen into focus
- If the microscope uses a mirror as the light source, adjust the mirror so enough light is reflected through the aperture onto the specimen
- You can adjust the amount of light reaching the specimen by opening and closing the diaphragm
- Once the object is visible, use the fine adjustment knob for a more precise focus
- At this point you can increase the magnification by switching to a higher power objective lens
- Once you switch from the low power objective lens, you should no longer be using the coarse adjustment knob for focusing because it is possible to break the slide and scratch the lenses
- If you switch objectives, use the fine adjustment to fine-tune the focus of the object If the high powered objective lenses on the microscope say oil then you can place a small drop of immersion oil on the cover slip then switch to the oil immersion lens. Only use the oil immersion lens with immersion oil and don’t use oil with any other objective that does not say oil.
- Once you have finished observing the specimen, lower the stage, remove the slide, and return to the lowest objective
- Clean the lenses with lens cleaner and lens paper (only use lens paper as other tissues will scratch the lenses)
- Wrap the cord around the base and cover the microscope for storage

##24.3 Making a Wet Mount

- Collect a thin slice (one cell layer thick is optimal) of specimen and place on the slide
- Place a drop of water directly over the specimen
- Place a cover slip at a 45 degree angle over the specimen with one edge touching the drop of water then drop the cover slip over the specimen. If done correctly, the cover slip will completely cover the specimen and there will be no air bubbles present.

##24.4 Staining a Slide

- Once you have completed the above process place one small drop of stain (ex. Iodine, methylene blue) on the outside edge of the cover slip
- Place the flat edge of a paper towel on the other side of the cover slip. The paper towel will draw the water out from under the cover slip and pull in the stain

##24.5 Magnification

The actual power of magnification is a product of the ocular lens (usually 10x) times the objective lens.

##24.6 Troubleshooting

1. _The Image is too dark!_
2. _There’s a spot in my viewing field, even when I move the slide the spot stays in the same place!_
Your lens is dirty. Use lens paper, and only lens paper to carefully clean the objective and ocular lens. The ocular lens can be removed to clean the inside.
3. _I can’t see anything under high power!_
Remember the steps, if you can’t focus under scanning and then low power, you won’t be able to focus anything under high power.
4. _Only half of my viewing field is lit, it looks like there’s a half-moon in there!_
You probably don’t have your objective fully clicked into place.
chapter25
#Chapter 25: Low Tech Microscopy

Microscopes are powerful tools for teaching biology, and many of their benefits are hard to replace with local fabrications. However, simple materials can be used to achieve sufficient magnification to greatly expands students’ understanding of the very small. They may view up close the anatomy of insects and even see cells.

##25.1 Water as a Lens

Water refracts light much the way glass does; a water drop with perfect curvature can make a powerful lens. A simple magnifier can be made by twisting a piece of wire around a nail and dipping the loop briefly into some water. Students can observe the optical properties of the trapped drop of water.

##25.2 Perfect Circles

Better imaging can be had if the drop is more per- fect in shape – the asymmetry of the wire twisting distorts the image. Search for a piece of thin but stiff plastic – water bottles work well. Cut a small piece of this plastic, perhaps 1×2 centimeters. Near one end, make a hole, the more perfect the better. The best hole-cutting tool is a paper hole punch, available in many schools. With care, fine scissors or a pen knife will suffice; remove all burrs.

![perfect-circle.jpg](images/perfect-circle.jpg)

##25.3 Slides

A slide and even cover slip may be made from the same plastic water bottles, although being hy- drophobic they will not have the same properties of glass when making wet mounts. Improvise a method for securing the punctured plastic over the slide; ideally the vertical spacing can be closely ad- justed to focus.

![slides.jpg](images/slides.jpg)

##25.4 Backlighting

On a bright day, there may not be any need for additional lighting, but in most classrooms the image will be too dim to be easily seen. The sun is a powerful light source, though not always con- venient. Flashlights are generally inexpensive and available; many cell phones have one built in the end. To angle the light into the slide, find either a piece of mirror glass, wrinkle-free aluminum foil, the metalized side of a biscuit wrapper, etc.

Experiment with a variety of designs to see what works best given the materials available to your school. If you use a slide of onion cells stained with iodine solution , your students should be able to see cell walls and nuclei.

##25.5 Simple Microscopes and Magnifiers

###25.5.1 Clear-Container Magnifiers

![clear-container-magnifier.jpg](images/clear-container-magnifier.jpg)

Any of these containers filled with water will make good magnifiers.

###25.5.2 Simple Microscope

![simple-microscope.png](images/simple-microscope.png)

Construct a small wooden box from plywood as shown (or use a small cardboard carton such as a light bulb box). Make a round hole of 2 cm diameter, at the top. Fit a small mirror (glass or polished metal) in the box, angled to reflect light up through the hole. Make a small hole (about 6 mm) in a strip of metal. Remove the round top from a pen-torch bulb and secure it in the strip using adhesive tape. Carefully cut off the tape where it may cover the lens. Bend the strip, then fix it to the side of the box, so that it can be moved up and down. Drawing pins or nails could be used for this. The object is focused by moving this strip. Note the eye should be placed as near as possible to the lens when viewing.

###25.5.3 Simple Compound Microscope

![compound-microscope.jpg](images/compound-microscope.jpg)

- Using 2 lenses together allows much greater magnification.
- Use a hand lens to make a water drop into a more powerful magnifier.
- Try using a hand lens with a lens from a torch bulb to make another simple compound microscope.

###25.5.4 Card Bridge Microscope

![card-microscope.jpg](images/card-microscope.jpg)

- Place a water drop in the card ‘bridge’.
- Place this on a sheet of glass as shown.
- Place the object you are looking at on the glass. This arrangement is most suitable for thin items, e.g. sections of leaves.
- Experiment with the angle of the mirror so that light shines up through the specimen.
- Use this arrangement with a hand lens to produce a compound microscope.
chapter26
#Chapter 26: Biology Practicals

##26.1 Introduction to Biology Practicals

###26.1.1 Format

The format of the Biology practical exam was revised in 2011 to keep up with the 2007 updated syllabus. As such, there will be no further Alternative to Practical exams, pending approval from the Ministry of Education. Prior to 2011, there were 3 questions in the Biology practical. Question 1 was required and the student could choose to answer either Question 2 or 3.

As of now, the Biology practical has 2 questions and students must answer both. Question 1 can come from any of the following topics: Nutrition, Movement, Transport of Living Things, Respiration, Reproduction, Coordination, Regulation or Growth. Question 2 is on Classification of Living Things. Each question is worth 25 marks, and students have 21 hours to complete the exam.

###Biology 1 Theory Format

The theory portion of the Biology exam comprises 100 marks, while the practical carries 50 marks. A student’s final grade for Biology is thus found by taking her total marks from both exams out of 150.

The theory exam for Biology contains 13 questions over 3 sections. Section A has 2 questions worth 10 marks each. Section B has 8 short answer questions, each having two items. This section weighs a total of 60 marks, and the mark allocation for individual questions is indicated at the end of each question. Section C has 3 long answer/essay questions, though students only need to answer 1 of them. The answer must be comprehensive and include as many points as possible. It is worth 20 marks.

Note This information is current as of the time of publication of this manual. Updated information may be obtained by contacting the Ministry of Education.

###26.1.2 Notes for Teachers

There are two sets of advance instructions. One set of advance instruction are given to teachers at least one month before the date of the exam. These instructions contain the list of specimen, apparatus, and other materials required for setting up the Biology practical questions. The instructions also suggest how many specimen to acquire for each candidate or group of candidates. It is imperative that the collection and storage of specimen for the practical be kept confidential.

The second set of instructions should be given 24 hours before the time of the practical. It includes how to label each specimen and which materials should be given to each candidate (or shared among candidates).

Usually if the instructions include a scalpel of some sort, this means students will be required to do some form of dissection. In most cases, the dissection is of a maize seed or bean to show whether is it a monocotyledon or dicotyledon. If the advance instructions include any form of glass apparatus or test tubes, there will most likely be a question on food tests.

###26.1.3 Common Practicals

**Food Tests** test a food solution for starch, sugars, fats, and protein

**Respiration** use lime water to test air from the lungs for carbon dioxide

**Transport** investigate osmosis by placing leaf petioles or pieces of raw potato in solutions of different solute concentrations

**Photosynthesis** test a variegated leaf for starch to prove that chlorophyll is necessary for photosynthesis

**Coordination** students look at themselves in the mirror and answer questions about the sense organs they see

**Movement** name bones and answer questions about their structure and position in the body

*Note* These are the most common practicals, but they are not necessarily the only practicals that can occur on the national exam. Food tests and Classification are by far the most common, but there are many eligible topics. Be sure to regularly look through Biology Past Papers (p. 162) to get an idea of the kind of questions that can occur.

##26.2 Food Tests

In this practical, students test a solution of unknown food substances for starch, protein, reducing sugars, non-reducing sugars, and lipids. They record their procedure, observation, and conclusions, then answer questions about nutrition and the digestive system.

This section contains the following:

- Preparation of Chemical Solutions
- Preparation of Food Solutions
- Performing the Food Tests
- Examination Room
- Student Report
- Sample Food Test Practical

###26.2.1 Preparation of Chemical Solutions

Always make sure the chemicals work before performing the food tests with students.

**Benedict’s Solution**

This solution can be bought at a chemical store already prepared or you can make it yourself.

**Using Sodium Carbonate:**

- Add 5 spoons of sodium carbonate (NaCO3).
- Add 3 spoons of citric acid.
- Add one spoon of copper sulphate.

**Using Bicarbonate of Soda:**

- Add 1 L of water to a cooking pot.
- Add a box (70 g) of bicarbonate of soda.
- Boil the mixture for 5-10 minutes. This makes sodium carbonate.
- Let cool and transfer to a plastic water bottle.
- Add 3 spoons of citric acid.
- Add one spoon of copper sulphate. Cap and shake to mix.

Label as: BENEDICT’S SOLUTION FOR FOOD TESTS
The solution may be stored in any plastic or glass bottle and will keep indefinitely.

**Copper (II) Sulphate**

- Add one spoon of copper (II) sulphate to a 1.5 L bottle.
- Add 1 L of water and shake until chemicals are fully dissolved.

Label as: 1% COPPER (II) SULPHATE SOLUTION FOR FOOD TESTS
The solution may be stored in any plastic or glass bottle and will keep indefinitely.

**Iodine Solution**

Make sure to use iodine tincture from a pharmacy. The tincture must not contain ethanol/alcohol/spirit.

- Add 1 part iodine tincture to 10 parts water. Example: In a 500 mL bottle, add 40 mL iodine tincture, then and 400 mL of water.
- Cap the bottle and shake.

Label as: IODINE SOLUTION FOR FOOD TESTS
The solution may be stored in any plastic or glass bottle and will keep indefinitely.

**Dilute NaOH**

- Using a PLASTIC teaspoon, add one level teaspoon of NaOH to a 500 mL water bottle. Caustic soda (NaOH) reacts with metal. DO NOT TOUCH.
SAFETY NOTE: Prepare about 100 mL of citric acid or ethanoic acid solution to neutralize sodium hydroxide spills on skin or lab tables. One spoon of citric acid in 100 mL of water is suitable. Ethanoic acid solutions are sold in stores as vinegar.
- Add 250 mL of water.
SAFETY NOTE: This reaction can cause the solution to become very warm. Avoid chemical burns by wearing gloves.
- Cap well and shake. This makes 1 M sodium hydroxide solution.

Label as: 1 M SODIUM HYDROXIDE SOLUTION FOR FOOD TESTS (CORROSIVE)
The solution will react with carbon dioxide in the air if not well sealed. Do not store in glass bottles with glass stoppers as these will stick. The solution may be stored in plastic bottles indefinitely.

**Dilute Acid**

Your school may have dilute hydrochloric acid or you may have to make it yourself.

**Using Hydrochloric Acid (HCl):**

- Add 1 part HCl to 9 parts water. Example: In a 1.5 L water bottle, add 900 mL of water, then add 100 mL of HCl.
- Shake well.

**Using Citric Acid:**

- Add 5 spoons of citric acid.
- Cap well and shake. This makes 0.5 M citric acid.

Label as: 0.5 M CITRIC ACID FOR FOOD TESTS
The solution may be stored in any plastic or glass bottle and will keep indefinitely.

**Sudan III Solution**

Using Sudan III solution takes a long time to show results. It may be replaced by iodine tincture solution for the lipids test.

- Combine 0.5 g of Sudan III powder with 100 mL of 70% ethanol solution (30 mL water and 70 mL ethanol).
- Place the solution in a warm water bath to help the Sudan III dissolve.
- Filter to remove any remaining solid.

Label as: SUDAN III SOLUTION FOR FOOD TESTS
The solution may be stored in any plastic or glass bottle and will keep indefinitely.

###26.2.2 Preparation of Food Solutions

For the NECTA and mock exams you may have to set up the food test solutions. The instructions will tell you which ones you’ll need to prepare in order to make ‘Solution X.’ Solution X consists of a mixture of at least 3 of the different food substances and is given to each student in at least 3 test tubes. Make sure food solutions are well-dissolved and colorless so that students don’t know what is in the mixture. You don’t need to measure the ingredients, but make sure to test the solutions before the practical.

**Reducing sugar**

Use glucose powder and dissolve in water. Make sure the substance is fully dissolved so that students don’t know what is in the mixture.

**Non-reducing sugar**

Use sugar and dissolve in water. Make sure the substance is fully dissolved so that students don’t know what is in the mixture.

**Lipids**

Mix sunflower oil with water. Shake immediately before use. Sunflower oil is best since it is liquid at room temperature.

**Protein**

Mix an egg white with water.

**Starch**

Save the water you use to boil potatoes, rice, or pasta. Make sure to remove the bits of food. You can also just mix flour in water, but it would be obvious.

###26.2.3 Performing the Food Tests

**Reducing Sugars Test**

- Add a small amount of Benedict’s solution to the food solution.
- Boil the solution and allow it to cool. Observe the colour changes from blue to green, yellow, then deep orange/brick red precipitate if reducing sugars are present.

Always do the reducing sugars test first because a non-reducing sugar will always test positive for a reducing sugar.

**Non-reducing Sugars Test**

- Add a small amount of dilute acid (HCl) to the solution.
- Boil the solution for about 30 seconds and allow it to cool.
- Add a small amount of NaOH to the solution and shake.
- Add a small amount of Benedict’s solution and boil.
- Allow the solution to cool and observe as the solution changes from green to yellow, then to deep orange/brick red precipitate if non-reducing sugars are present.

**Lipids Test**

- Add a small amount of Sudan III or iodine solution to the food solution and shake.
- A red ring will form at the top of the test tube if lipids are present.

Using Sudan III colours the whole solution red whether it contains lipids or not. Use iodine solution to get a more distinct result.

**Protein Test**

- Add an equal amount of sodium hydroxide (NaOH) to the solution and shake.
- Add a small amount of copper (II) sulphate to the solution and shake.
- Observe the solution turn violet/purple in colour if protein is present.

**Starch Test**

- Add a small amount of iodine solution to the food solution.
- Observe the solution turn blue-black in colour if starch is present.

###26.2.4 Examination Room

The NECTA practical exam is done in the school’s lab or any other suitable room. Heat sources (jiko, etc.) should be spread evenly in the exam room so that students don’t have to go far to heat their test tubes; this also cuts down on cheating. Spread students out and distribute supplies as you see fit.

**Each student gets:**

- 3 or more test tubes (to carry out 5 tests)
- A beaker containing Solution X
- A test tube rack (or a cut out water bottle with sand to hold the tubes)

*Note* Students may share the racks, but shouldn’t share the cut out bottles

**Each station should have:**

- Copper II sulphate
- Water
- Dilute acid (HCl, etc.)
- Dilute base (sodium hydroxide)
- Iodine solution
- Sudan III solution (can be replaced by iodine solution)
- Benedict’s solution

###26.2.5 Student Report

Food test data is recorded in a table containing four columns: Test for, Procedure, Observation and Inference.Students should write the Procedure using the passive voice in the past tense. For example, “A small amount of Benedict’s solution was added to the solution. Then the solution was boiled and allowed to cool.”

In the Observation column, the student should write what they observed using the past tense and passive voice. For example, “A violet colour was observed.”

In the Inferences column, the students should write what they saw in the past tense and passive voice. For example, “Reducing sugars were not (or were) present.” Note that every column is worth marks on the exam. Even if students fail to do the food tests correctly, they can still get marks for writing what they are testing for and what the procedure should be.

An example of a completed food test results table is given below. Assume the solution contains proteins, reducing sugars, non-reducing sugars and starch.

| **Food Tested** | **Procedure** | **Observation** | **Inference**|
| --------------- | ------------- | --------------- | -------------|
| Lipids | A few drops of Sudan III solution (or iodine solution) were added to solution X. The solution was shaken and allowed to stand.| A red ring did not form at the surface.| Lipids were not present. |
| Proteins | An equal amount of NaOH was added to solution X and shaken. A few drops of copper (II) sulphate were added to solution X and shaken again. | A violet colour was observed. | Proteins were present.|
| Reducing sugars | A small amount of Benedict’s so- lution was added to solution X. The solution was heated and al- lowed to cool. | A brick red precipitate was observed.| Reducing sugars were present. |
| Non-reducing sugars | A small amount of dilute acid was added to solution X. The so- lution was heated and allowed to cool. Then a small amount of NaOH solution was added, and the solution was shaken. Finally, a small amount of Benedict’s so- lution was added. The solution was boiled and let cool. | The solution changed from green to yellow, then to a deep orange/brick red precipitate. | Non-reducing sugars were present. |
| Starch | A few drops of iodine solution were added to solution X and shaken. | A blue-black colour was observed. | Starch was present. |

###26.2.6 Sample Food Test Practical

**You have been provided with solution B.**

1. Identify the food substances present in solution B by using the reagents provided. Tabulate your work as shown in the following Table:

| **Food Tested** | **Procedure** | **Observation** | **Inference** |
| --------------- | ------------- | --------------- | ------------- |
| | | | |
| | | | |

2. For each food substance identified in Question 1,

(i) Name two common sources.

(ii) State their role in the body of human being.

3. The digestion of one of the identified food substance in Question 1 starts in the mouth.

(i) Name this food substance.

(ii) Identify the enzyme responsible for its digestion in the mouth.

4. The digestive system of human being has several parts.

(i) Name the part of digestive system in which most of digestion and absorption of food takes place.

(ii) Explain how the named part in 4(i) is adapted for absorption of digested food substances.

**Additional Food Test Questions:** See Biology Past Papers (p. 162) for additional food test questions.

##26.3 Classification

The classification practical requires students to identify specimens of animals, plants, and fungi. The students must write the common name, kingdom, phylum, and sometimes class of each specimen. They also answer questions about the characteristics and uses of the specimens.
This section contains the following:

- Common specimens
- Where to find specimens
- Storage of specimens
- Sample practical with solutions

###26.3.1 Common Specimens

- **Fungi:** Mushroom, yeast, bread mold
- **Plants:** Fern, moss, bean plant, bean seed, maize plant, maize seed, pine tree, cactus, sugar cane, Irish potato, cypress tree, acacia tree, hibiscus leaf, cassava
- **Animals:** Millipede, centipede, grasshopper, lizard, tilapia (fish), scorpion, frog, tapeworm, liver fluke, cockroach, spider

###26.3.2 Where to Find Specimens

- Start collecting specimens several months before the NECTA exams, as some specimens can be hard to find in the dry season.
- Ask your students to bring specimens! Students are especially good at finding insects and other animals. You can even find primary school children to gather insects such as grasshopper and millipedes.
- Ferns, hibiscus, pines, and cypresses are used in landscaping. Try looking near nice hotelis or guestis. Ferns should have sori (sporangia) on the underside of their leaves.
- Moss often grows near water tanks and in shady corners of courtyards. It is hard to find in the dry season.
- Sugarcane, Irish potato, cassava, tilapia, bean seeds, and maize seeds can be found at the market. Yeast is available at shops.
- Mushrooms are hard to find in the dry season. However, they are available at grocery stores in large cities, and you may be able to find dried mushrooms at the market. You can also collect mushrooms in the rainy season and dry them yourself.
- Tapeworms and liver flukes may be acquired from butchers. Find out where livestock is slaughtered and ask the butchers to look for worms (minyoo). Liver flukes are found in the bile ducts inside the liver, while tapeworms are found in the intestines. You can also try going to a livestock fair/market (mnada) or talking to the local meat inspector (mkaguzi wa nyama).
- Grow your own bread mold. Just put some bread in a plastic bag and leave it in a warm place. But do it ahead of time – it can take two weeks to obtain bread mold with visible sporangia.

###26.3.3 Storage of Specimens

- Insects and mushrooms can be dried and stored in jars. However, they become brittle and break easily.
- A 10% solution of formaldehyde is the best way of storing specimens. Formaldehyde is often sold as a 40% solution. It should be stored in glass jars and out of the sun. Check specimens periodically for evaporation. Formaldehyde works because it is toxic; handle carefully.
- In a pinch, a 70% solution of ethanol can also be used to store insects, lizards, and worms. However, specimens sometimes decay in ethanol.

###26.3.4 Sample Classification Practical

**You have been provided with specimens L, M, N, O, and P.**

1. Identify the specimens by their common names.
2. Classify each specimen to the phylum level.
3. Further classification:
3.1. Write the classes of specimens L and M.
3.2. List two observable differences between specimens L and M.
4. Explain why specimen P cannot grow taller.
5. Write down two distinctive characteristics of the phylum to which specimen O belongs.
6. Reproduction:
6.1. List the modes of reproduction in specimens M and N.
6.2. What are two differences between these modes of reproduction?

###26.3.5 Sample Practical Solutions

1. Common names of specimens:
- L: maize plant
- M: bean plant
- N: yeast
- O: millipede
- P: moss

2. Classifaction by kingdom and phylum:

| Specimen | Kingdom | Phylum |
| -------- | ------- | ------ |
| L (Maize) | Plantae | Angiospermophyta |
| M (Bean) | Plantae | Angiospermophyta |
| N (Yeast) | Fungi | Ascomycota |
| O (Millipede) | Animalia | Arthropoda |
| P (moss) | Plantae | Bryophyta |

3. Further classification:
- Specimen L (maize plant): Class Monocotyledonae
- Specimen M (bean plant): Class Dicotyledonae
- Observable differences:

| Specimen | Vein structure | Root structure |
| -------- | -------------- | -------------- |
| L (maize) | Parallel veins | Fibrous roots |
| M (bean plant) | Net veins | Tap roots |

_The answers to this question should be differences between monocots and dicots that the student can see by observing the plants with their naked eyes. Hence answers such as “vascular bundles in a ring” are not correct._

4. Specimen P (moss) cannot grow taller because it has no xylem and phloem. If it grew taller, it would not be able to transport food and water throughout the plant.

5. Characteristics of phylum Arthropoda:
- jointed legs
- segmented body
6. Reproduction
6.1. Specimen M (bean plant) reproduces by sexual reproduction. Specimen N (yeast) reproduces by asexual reproduction.

| Method | Genetic variation | Parents | Gametes |
| ------ | ----------------- | ------- | ------- |
| Asexual reproduction | There is NO genetic variation between offspring | Requires only one parent | No gametes are involved |
| Sexual reproduction | There IS genetic variation between offspring | Usually requires two parents | Involves fusion of two gametes |

- Identify specimen X, Y, and Z by their common names.
- Classify specimens X, Y, and Z to the class level. (This means write the kingdom, phylum, and class.)
- Write the observable features of specimen X.
- List three observable differences/similarities between specimens X and Y.
- State the economic importance of specimen X.
- What characteristics are common among specimens X and Y?
- Why are specimens X and Y placed in different classes/phyla/kingdoms?
- Why are specimens X and Y classified under the same class/phylum/kingdom?
- What distinctive features place specimen X in its respective kingdom/phylum/class?
- How is specimen X adapted to its way of life?
- Suggest possible habitats for specimens X and Y.
- Which specimen is a primary producer/parasite/decomposer?
- For mushroom, yeast, bread mold, grasshopper, moss, tilapia, liver fluke, and tapeworm: Draw and label a diagram of specimen X.
- For tilapia: Draw a big and well-labeled diagram showing a lateral view of specimen X.

For maize and bean:

- Mention the type of pollination in specimen X [wind pollinated or insect pollinated].
- How is specimen X adapted to this type of pollination?
- Mention the type of germination [hypogeal or epigeal] in specimen X.

For bean seed:

- List three observable features of specimen X and state their biological importance.
- Split specimen X into two natural halves. Draw and label the half containing the embryo.

For fern:

- Observe the underside of the leaves of specimen X
- What is the name of the structures you have observed?
- Give the function of the structures named above.
- Draw specimen X and show the structures named above.

##26.4 Respiration

The purpose of this practical is to investigate the properties of air exhaled from the lungs. This section contains the following:

- Limewater (properties and preparation)
- Apparatus
- Sample practical with solutions

###26.4.1 Limewater

Limewater is a saturated solution of calcium hydroxide. It is used to test for carbon dioxide. When carbon dioxide is bubbled through limewater, the solution becomes cloudy. This is due to the precipitation of calcium carbonate by the reaction:

CO$$_2$$(g) + Ca(OH)$$_2$$(aq) → CaCO$$_3$$(s)

Limewater can be prepared from either calcium hydroxide or calcium oxide. Calcium oxide reacts with water to form calcium hydroxide, so either way you end up with a calcium hydroxide solution. Calcium oxide is the primary component in cement. Calcium hydroxide is available from building supply shops as chokaa.

To prepare lime water, add three spoons of fresh chokaa or cement to a bottle of water. Shake vigorously and then let stand until the suspended solids precipitate. Decant the clear solution. Chokaa produces a solution much faster than cement.

The exact mass of calcium hydroxide or calcium oxide used is not important. Just check whether some calcium hydroxide remains undissolved at the end – a sign that you have made a saturated solution. Test limewater by blowing air into a sample with a straw. It should become cloudy. If it does not, then the concentration of Ca(OH)2 is too low.

###26.4.2 Apparatus
Many books call for delivery tubes, test tubes, and stoppers. These are totally unnecessary. Add the limewater to any small clear container and blow into it with a straw.

If you use a delivery tube and pass it through a rubber stopper, do not use a single-holed stopper. This is what the pictures on NECTA practicals suggest, but it is a terrible idea. A single-holed stopper has no space for air to escape. So when a student blows air into the solution, the pressure in the test tube increases. The high pressure air then pushes limewater up the straw into the student’s mouth. Alternatively, the student blows the stopper out of the test tube. If you use a stopper, use a double-holed stopper so that the extra air has a place to escape.

Is a glass delivery tube stuck in a rubber stopper? Do not pull hard on it. Just soak the stopper in warm water for a few minutes. The rubber will soften and the tube will come out.

Are your test tubes and delivery tubes cloudy after the practical? Clean them with dilute acid. This will dissolve any calcium carbonate that has been deposited on the glass.

###26.4.4 Sample Respiration Practical

You have been provided with Solution B in a test tube. Use a delivery tube to breathe (exhale) into the solution until its color changes. (See diagram below.)

1. What is the aim of this experiment?
2. What is Solution B?
2.1. What changes did you observe after breathing into Solution B?
2.2. What can you conclude from these changes?
3. Breathe out over the palm of your hand. What do you observe?
4. Breathe out over a mirror. What do you observe?
5. Using your observations in the three experiments above, list three properties of exhaled air.
6. Explain why exhaled air is different from inhaled air. Where do the substances you identified in exhaled air come from?

###26.4.5 Sample Practical Solutions
1. The aim of this experiment is to test exhaled air for carbon dioxide.
2. Solution B is limewater.
- Solution B became cloudy (or milky).
- Conclusion: exhaled air contains carbon dioxide.
3. Air breathed out over the palm of the hand is warm.
4. Droplets of water condense on the mirror.
5. Conclusions:
- exhaled air contains carbon dioxide
- exhaled air contains water
- exhaled air is warm
6. Exhaled air contains the waste products of aerobic respiration. The carbon dioxide and water in exhaled air are products of respiration.

##26.5 Transport
The purpose of this practical is to investigate osmosis by observing the changes in a leaf petiole placed in a hypotonic solution (water) and a hypertonic solution (water containing salt or sugar). This section contains the following:

- Materials
- Sample practical with solutions

###26.5.1 Materials

The petiole is the stalk which attaches a leaf to a branch. The papaya leaf petioles in this practical should be soft petioles from young leaves, not stiff petioles from older leaves. Cut the petioles into pieces, and give each student two pieces of about 6 cm in length. Cylinders cut from a raw potato may be used instead of petioles. The hypertonic solution may be made with by mixing either salt or sugar with water. The hypotonic solution is tap water.

###26.5.2 Sample Transport Practical

**Instructions**

You have been provided with two pieces of a papaya leaf petiole, Solution A, and Solution B.
Use a razor blade to split the pieces of petiole longitudinally, up to a half of their length. You should have four strips at one end of each petiole, while the other end remains intact.
Place one petiole in solution A, and place the other petiole in solution B. Let the petiole sit for about
ten minutes, then touch them to feel their hardness or softness.
Draw a sketch of each petiole after sitting in its respective solution for ten minutes. Record your observations and explanations about the petioles in the table below.

| Solution | Observation | Explanation |
| -------- | ----------- | ----------- |
| A | | |
| B | | |

**Questions**

1. What was the aim of this experiment?
2. What was the biological process demonstrated by this experiment?
3. What is the importance of this process to plants?
4. Which solution contained:
-pure water
-a high concentration of solutes
5. What happened to the cells of the petioles in each solution? Illustrate your answer.
6. What would happen to the cells of the petioles in solution A if their cell walls were removed?

###26.5.3 Sample Practical Solutions
(Assume Solution A is pure water, and Solution B is a concentrated solution of water and salt.)

| Solution | Observation | Explanation |
| -------- | ----------- | ----------- |
| A | The petiole became hard (turgid) | Water diffused into the petiole cells |
| B | The petiole became soft (flaccid) | Water diffused out of the petiole cells |

1. The aim of the experiment was to investigate the effect of osmosis on plant cells.
2. The experiment demonstrated osmosis.
3. Importance of osmosis in plants:
3.1. Water enters plant cells by osmosis so that they become turgid. Turgor helps support the plant and hold it upright.
3.2. Water diffuses into the xylem from the soil via osmosis.
4. Solution identification
4.1. Pure water: Solution A.
4.2. High concentration of solutes: Solution B
5. [Illustrations]
6. The petiole cells would burst in Solution A if their cell walls were removed.

You can extend this experiment by giving students two pieces of meat in addition to the petioles. The piece of meat placed in pure water should expand and become soft due to the cells bursting. The piece placed in salt water should shrink and become hard due to water diffusing out of the cells. This experiment helps to teach the different effects of osmosis on plant and animal cells.

If your school has a good microscope, try observing plant cells under the microscope after letting them sit in hypotonic and hypertonic solutions.
You can add critical thinking questions to the practical that require the student to use their knowledge of osmosis. For example:

- Why does a freshwater fish die if it is placed in salt water?
- Why do merchants spray vegetables with water in the market?
- You can die if a doctor injects pure water into your bloodstream. Why?

##26.6 Photosynthesis
The purpose of this practical is to prove that chlorophyll is required for photosynthesis. This is done by using iodine to test a variegated leaf for starch. The parts of the leaf containing chlorophyll are expected to contain starch, while the parts lacking chlorophyll are expected to lack starch. This section contains the following:

- Procedure
- Cautions
- Materials and where to find them
- Sample practical with solutions

###26.6.1 Procedure
1. Use iodine tincture from the pharmacy without dilution.
2. Prepare hot water bathes. The water should be boiling.
3. While the water gets hot, send the students to gather small leaves. The best have no waxy coating and are varigated (have sections without green).
4. The leaves should be boiled in the hot water bath for one minute.
5. Each group should then move its leaf into their test tube and cover it with methylated spirit.
6. Each group should then heat their test tube in a water bath. Over time, the leaf should decolorize and the methylated spirit will turn bight green. The chlorophyll has been extracted and moved to the spirit. A well chosen leaf should turn completely white, although this does not always happen.
7. After decolorization, dips the leaves briefly in the hot water.
8. For leaves that turn white, students should test them for starch with drops of iodine solution.

###26.6.2 Cautions
Ethanol is flammable! It should never be heated directly on a flame. Use a hot water bath – place a test tube or beaker of ethanol in a beaker or bowl of hot water and let it heat slowly. The boiling point of ethanol is lower than the boiling point of water, so it will start boiling before the water. If the ethanol does catch fire, cover the burning test tube with a petri dish or other non-flammable container to extinguish the flame.

###26.6.3 Materials and Where to Find Them
- Variegated leaf: this is a leaf that contains chlorophyll in some parts, but not in others. Often variegated leaves are green and white or green and red. Look at the flower beds around the school and at the teachers’ houses – they often contain variegated leaves. Test the leaves before the practical, as some kinds are too waxy to be decolorized by ethanol. Also, check for chlorophyll by looking at the underside of the leaves; the leaves you use have at least a small section of white on their undersides, signifying a lack of chlorophyll.
- Source of heat: anything that boils water – Motopoa is best, followed by kerosene and charcoal
- Ethanol: use the least expensive strong ethanol available; this is probably methylated spirits unless your village specializes in high proof gongo.

###26.6.4 Sample Photosynthsis Practical

**You have been provided with specimen G.**

1. Identify specimen G.
2. Make a sketch showing the color pattern of specimen G. Carry out the following experiment:
2.1. Place specimen G in boiling water for one minute.
2.2. Boil specimen G in ethanol using a hot water bath. Do not heat the ethanol directly on a flame.
2.3. Remove specimen G from the ethanol. Dip it in hot water.
2.4. Spread specimen G on a white tile and drip iodine solution onto it. Use enough iodine to cover the entire specimen.
2.5. Make a sketch showing the color pattern of specimen G at the end of the experiment.
3. What was the aim of this experiment?
4. Why was specimen G
4.1. Boiled in water for one minute
4.2. Boiled in ethanol
4.3. Dipped in hot water at the end of the experiment
5. What was the purpose of the iodine solution?
6. Why was the ethanol heated using a hot water bath?
7. What can you conclude from this experiment? Why?

###26.6.5 Sample Practical Solutions
1. Specimen G is a variegated leaf.
2. Drawing: See diagram above.
3. The aim of this experiment was to investigate whether chlorophyll is required for photosynthesis.
4. Specimen G was:
4.1. boiled in water to kill the cells and stop all metabolic processes.
4.2. boiled in ethanol to decolorize it (to remove the chlorophyll).
4.3. dipped in hot water to remove the ethanol. (If ethanol is left on the leaf it will become hard and brittle.)
5. The purpose of the iodine solution was to test for starch.
6. The ethanol was heated using a hot water bath because ethanol is flammable.
7. The experiment shows that chlorophyll is required for photosynthesis. We know this because the parts of the leaf containing chlorophyll also contained starch, which is a product of photosynthesis. Thus, the parts of the leaf containing chlorophyll performed photosynthesis. The parts of the leaf lacking chlorophyll lacked starch. Hence, these parts of the leaf did not perform photosynthesis.

**To test if light is required for photosynthesis:**

Take a live plant, and leave it in the dark for 24 hours to destarch all leaves. Then, cover some of its leaves with cardboard or aluminum foil, while leaving others uncovered. Let the plant sit in bright light for several hours. Give each group of students one leaf that was covered in cardboard, and one leaf that was uncovered. Have them use the procedure above to test for starch. They should find that the covered leaf contains no starch, while the uncovered leaf contains starch.

A cool variation on this experiment is to cover leaves with pieces of cardboard that have letters or pictures cut out of them. The area where the cardboard is cut out will perform photosynthesis and produce starch. When the students do a starch test, a blue-black letter or picture will appear on the leaf.

**To prove that oxygen is a product of photosynthesis:**

This experiment requires a water plant. Basically, place a live water plant under water*, then cover it with an inverted funnel. Place an upside-down test tube filled with water on top of the funnel. Let the plant sit in bright light until the water in the test tube is displaced and the test tube fills with gas. Use a glowing splint to test the gas – if it is oxygen, it will relight the splint.

**Note:** some books suggest putting sodium bicarbonate (baking soda) in the water.
chapter27
#Chapter 27: Chemistry Practicals

##27.1 Introduction to Chemistry Practicals

###27.1.1 Format
The format of the Chemistry practical exam was revised in 2011 to keep up with the 2007 updated syllabus. As such, there will be no further Alternative to Practical exams, pending approval from the Ministry of Education. Prior to 2011, students only had to answer 2 of the 3 questions, including Question 1.

As of now, the Chemistry practical has 3 questions and students must answer all of them. Question 1 is on Volumetric Analysis and Laboratory Techniques and Safety. Question 2 is taken from Ionic Theory and Electrolysis/Chemical Kinetics, Equilibrium and Energy. Question 3 is on Qualitative Analysis.

Question 1 is worth 20 marks, while Questions 2 and 3 carry 15 marks each. Students have 21 hours to complete the exam.

Students are allowed to use Qualitative Analysis guidesheet pamphlets in the examination room.

###Chemistry 1 Theory Format

The theory portion of the Chemistry exam comprises 100 marks, while the practical carries 50 marks. A student’s final grade for Chemistry is thus found by taking her total marks from both exams out of 150. The theory exam for Chemistry contains 3 sections. Section A has 2 questions and is worth 20 marks - Question 1 is 10 multiple choice and Question 2 is 10 matching. Section B has 9 short answer questions, each having two items, for a total of 54 marks. Section C has 2 essay questions without items for a total of 26 marks. Students are required to answer all questions.

**Note** This information is current as of the time of publication of this manual. Updated information may be obtained by contacting the Ministry of Education.

###27.1.2 Notes for Teachers

There are two sets of advance instructions. One set of advance instructions are given to teachers at least one month before the date of the exam. These instructions contain the list of apparatus, chemicals, and other materials required for preparing the Chemistry practical questions. The instructions also give suggestions on the amount of chemicals that should be available for each candidate to use.

The second set of instructions should be given 24 hours before the time of the practical. It includes which chemicals and apparatus should be given to each candidate (or shared among candidates) for each of the three practical questions. These instructions also state how to label each solution and/or compound.

The bottom of the 24 Hours Advance Instructions also states that the Laboratory Technician or Head of Chemistry Department should perform some of the experiments immediately after the last session of the examination. It is only required to perform the titration and chemical kinetics experiments. This is required to be done for every school and is used as a reference for the markers in case the water, chemicals, and apparatus are not the same at every school. This is enclosed and submitted together with the students’ test papers and may be used as a marking scheme. It is also advised that any notes, comments or concerns for the markers be included at this time.

###27.1.3 Common Practicals
**Volumetric Analysis** determine the concentration of a solution of a known chemical by reacting it with a known concentration of another solution

**Qualitative Analysis** systematically identify an unknown salt through a series of chemical tests

**Chemical Kinetics and Equilibrium** observe changes in chemical reaction rates by varying condi- tions such as temperature and concentration.

**Note** These are the most common practicals, but they are not necessarily the only practicals that can occur on a NECTA exam. Although the updated exam format lists Questions 1 and 3 as Volumetric Analysis and Qualitative Analysis respectively, Question 2 can come from a variety of topics which may not yet have been used in older past papers. Be sure to regularly check the most recent past NECTA papers to get a good idea of the types of questions to expect.

##27.2 Volumetric Analysis
This section contains the following:

- Volumetric Analysis Theory
- Substituting Chemicals in Volumetric Analysis
- Properties of Indicators
- Sample Practical Question

###27.2.1 Volumetric Analysis Theory

Volumetric Analysis is a method to find the concentration (molarity) of a solution of a known chemical by comparing it with the known concentration of a solution of another chemical known to react with the first.

For example, to find the concentration of a solution of citric acid, one might use a 0.1 M solution of sodium hydroxide because sodium hydroxide is known to react with citric acid.

The most common kinds of volumetric analysis are for acid-base reactions and oxidation-reduction reactions. Acid-base reactions require use of an indicator, a chemical that changes color at a known pH. Some oxidation-reduction reactions require an indicator, often starch solution, although many are self-indicating, (one of the chemicals itself has a color).

See also the sections on Preparation of Solutions (p. 63), Preparation of Solutions Without a Balance (p. 65) and Relative Standardization (p. 66) in Laboratory Techniques.

The process of volumetric analysis is often called _titration._

##27.2.2 Substituting Chemicals in Volumetric Analysis

**Theory**

The volumetric analysis practical exercises sometimes call for expensive chemicals, for example potassium hydroxide or oxalic acid. As the purpose of exercises and exams is to train or test the ability of the students and not the resources of the school, it is possible to use different chemicals as long as the solutions are calibrated to give equivalent results. For example, if the instructions call for a potassium hydroxide solution, you can use sodium hydroxide to prepare this solution. It will not affect the results of the practical – if you make the correct calibration. How to calibrate solutions when substituting chemicals is the subject of this section.

Technically, only two chemicals are required to perform any volumetric analysis practical: one acid and one base. The least expensive options are sulfuric acid, as battery acid, and sodium hydroxide, as caustic soda. To substitute one chemical for another in volumetric analysis, the resulting solution must have the same normality (N).

- For all monoprotic acids (HCl, ethanoic acid), the normality is the molarity.
_Example: 0.1 M ethanoic acid = 0.1 N ethanoic acid_
- For diprotic acids (sulfuric acid, ethandiotic acid), the normality is twice the molarity, because each molecule of diprotic acid brings two molecules of H$$^+$$.
_Example: 0.5 M sulfuric acid = 1.0 N sulfuric acid_
- For the hydroxides and hydrogen carbonates used in ordinary level (NaOH, KOH, NaHCO$$_3$$, the normality is the molarity.
_Example: 0.08 M KOH = 0.08 N KOH_
- For the carbonates most commonly used (Na$$_2$$CO$$_3$$, Na$$_2$$CO$$_3$$·10H$$_2$$O, K$$_2$$CO$$_3$$, the normality is twice the molarity.
_Example: 0.4 M Na$$_2$$CO$$_3$$ = 0.8 N Na$$_2$$CO$$_3$$_

**Substitution Calculations**

When instructions describe solutions in terms of molarity, calculating the molarity of the substitution is relatively simple. For example, suppose we want to use sulfuric acid to make a 0.2 M solution of ethanoic acid. 0.2 M ethanoic acid is 0.2 N ethanoic acid which will titrate the same as 0.2 N sulfuric acid. 0.2 N sulfuric acid is 0.1 M sulfuric acid, and thus we need to prepare 0.1 M sulfuric acid.

When instructions describe solutions in terms of concentration $$^\text{g}/_\text{L}$$, we just need to add an extra conversion step. For example, suppose we want to use sodium hydroxide to make a 14.3 $$^\text{g}/_\text{L}$$ solution of sodium carbonate decahydrate. 14.3 $$^\text{g}/_\text{L}$$ sodium carboante decahydrate is 0.05 M sodium carbonate decahydrate which is 0.1 N sodium carbonate decahydrate. This will titrate the same as 0.1 N sodium hydroxide, which is 0.1 M sodium hydroxide or 4 $$^\text{g}/_\text{L}$$ sodium hydroxide, and thus we need to prepare 4 $$^\text{g}/_\text{L}$$ sodium hydroxide to have a solution that will titrate identically to 14.3 $$^\text{g}/_\text{L}$$ sodium carbonate decahydrate.

**Common Substitutions**

To simplify future calculations, we have prepared general conversions for the most common chemicals used in volumetric analysis. Remember to check all final solutions with relative standardization to ensure that they indeed give the correct results.

Required Chemical Low Cost Alternative Substitution Method Molarity Example ConcentrationExample
Hydrochloric Acid Sulfuric Acid (Battery Acid) If you are required to prepare an X molarity solution of HCl, prepane a X×0.5 molarity solution of battery acid The instructions call for 0.12 M HCl. Instead, prepare 0.06 M sulfuric acid -
Ethanoic (Acetic) Acid Sulfuric Acid (Battery Acid) If you are required to prepare an M molarity solution of ethanoic acid, prepare a M×0.5 molarity solution of sulfuric acid The instructions call for 0.10 M ethanoic acid. Prepare 0.075 M sulfuric acid. -
Ethandioic (Oxalic) Acid dehydrate Sulfuric Acid (Battery Acid) If you are required to prepare an M molarity solution of ethandioic acid, prepare an M molarity solution of sulfuric acid. If you are required to prepare a C concentration solution of ethandioic acid, prepare a C/126 molarity solution of sulfuric acid. The instructions call for 0.075 M ethandioic acid. Prepare 0.075 M sulfuric acid. The instructions call for 6.3 $$g /L$$ ethandioic acid. Prepare 0.05 M sulfuric acid.
Potassium Hydroxide Sodium Hydroxide (Caustic Soda) For M molarity potassium hydroxide, make M molarity sodium hydroxide. For C concentration potassium hydroxide, make C×40/56 concentration sodium hydroxide. The instructions call for 0.1 M potassium hydroxide. Just prepare 0.1 M sodium hydroxide. The instructions call for 2.8 g/L potassium hydroxide. Prepare 2 g/L sodium hydroxide.
Anhydrous Sodium Carbonate Sodium Carbonate Dehydrate (Soda Ash) For M molarity anhydrous sodium carbonate, make M molarity sodium carbonate decahydrate. For C concentration anhydrous sodium carbonate, make C×286/106 sodium carbonate decahydrate. The instructions call for 0.09 M anhydrous sodium carbonate. Make 0.09 M sodium carbonate decahydrate. The instructions call for
5.3 g/L anhydrous sodium carbonate. Make 14.3 g/L sodium carbonate decahydrate.
Anhydrous Sodium Carbonate Sodium Hydroxide (caustic soda) For M molarity anhydrous sodium carbonate, make M×2 molarity sodium hydroxide. For C concentration anhydrous sodium carbonate, make C×2×40/106 sodium hydroxide. The instructions call for 0.09 M anhydrous sodium carbonate. Make 0.18 M sodium hydroxide. The instructions call for 5.3 g/L anhydrous sodium carbonate. 4.0 g/L sodium hydroxide.
Sodium Carbonate Decahydrate Sodium Hydroxide (caustic soda) For M molarity sodium carbonate ecahydrate, make M×2 molarity sodium hydroxide. For C concentration sodium carbonate decahydrate, make C×2 ×40 /286 sodium hydroxide. The instructions call for 0.09 M sodium carbonate decahydrate. Make 0.18 M sodium hydroxide. The instructions call for 14.3 g/L sodium carbonate decahydrate. Make 4.0 g/L sodium hydroxide.
Anhydrous Potassium Carbonate Sodium Carbonate decahydrate (soda ash) For M molarity potassium carbonate, make M molarity sodium carbonate decahydrate. For C concentration potassium carbonate, make C×286/122 concentration sodium carbonate. The instructions call for 0.08 M anhydrous potassium carbonate. Prepare 0.08 M sodium carbonate decahydrate. The instructions call for 6.1 g/L anhydrous potassium carbonate. Prepare 14.3 g/L sodium carbonate decahydrate.
Anhydrous Potassium Carbonate Sodium Hydroxide (caustic soda) For M molarity potassium carbonate, make M×2 molarity sodium hydroxide. For C concentration potassium carbonate, make C×2 ×40 /122 concentration sodium hydroxide. The instructions call for 0.08 M anhydrous potassium carbonate. Prepare 0.16 M sodium hydroxide. The instructions call for 6.1 g/L anhydrous potassium carbonate. Prepare 4.0 g/L sodium hydroxide.

- In volumetric analysis experiments with two indicators, it is not possible to substitute one chemical for another as the acid/base dissociation constant is critical and specific for each chemical. It is still possible to substitute sodium carbonate decahydrate for anhydrous sodium carbonate with the above conversion.
- These substitutions only work for volumetric analysis. In qualitative analysis, the nature of the chemical matters. If the instructions call for sodium carbonate, you cannot provide sodium hy- droxide and expect the students to get the right answer!

###27.2.3 Properties of Indicators

**Acid-base Indicators**
These indicators are chemicals that change colors in a specific pH range, which makes them suited to use in acid-base reactions. When the pH of changes from low pH to high pH or from high to low, the color of the solution changes.
Four common acid-base indicators are methyl orange (**MO**), phenolphthalein (**POP**), bromothymol blue (**BB**), and universal indicator (**U**)

- Methyl Orange, **MO**, is always used when titrating a strong acid against a weak base. The pH range of **MO** is 4.0 - 6.0 and thus no color change is observed until the base is completely neutralized. If you use **MO** with a weak acid, the color might start to change before completely neutralizing the acid.
- Phenolphthalein, **POP**, is always used when titrating a weak acid against a strong base. The pH range of **POP** is 8.3 - 10.0, and thus no color change is observed until the weak acid is completely neutralized. If you use **POP** with a weak base, the color might start to change before completely neutralizing the base.
- Bromothymol Blue, **BB**, is used in the same manner as methyl orange.
- Universal indicator, **U**, is not suitable for volumetric analysis involving either weak acids or bases as it changes color continuously rather than in a limited pH range. It is very useful for tracking the pH continuously over a titration, perhaps by performing two titrations side by side, one with a standard indicator and another with universal indicator.

Any indicator can be used when titrating a strong acid against a strong base. Universal indicator, however, will not produce very accurate results. No indicator is suitable for titrating a weak acid against a weak base. In some experiments, more than one indicator may be used in the same flask, for example when titrating a mixture of strong and weak acids or bases.

**Colors of Indicators** The colors of the above indicators in acid and base are:

| Indicator | Acid | Neutral | Base |
| --------- | ----- | ----- | ----- |
| Methyl Orange | Red | Orange | Yellow |
| Phenolphthalein | Colorless | Colorless | Pink |
| Bromothymol Blue | Yellow | Blue | Blue |
| Universal Indicator | Red, Orange, Yellow | Yellow/Green | Green, Blue, Indigo |

Titration is finished when the indicator starts a permanent color change. For example, when methyl orange turns orange, the titration is finished. If students wait until methyl orange turns pink (or yellow) they have overshot the endpoint of the titration, and their volume will be incorrect. Likewise, POP indicates that the titration is finished when it turns light pink. If students wait until they have an intensely pink solution, they will use too much base and get the wrong answer.

Note that light pink **POP** solutions may turn colorless if left for a few minutes. This is due to carbon dioxide in the air reacting to neutralize bases in solution.

**Note on technique** Students should use as little acid-base indicator as possible. This is because some acid or base is required to react with the indicator so that it changes color. If a lot of indicator is used, students will add more acid or base than they need.

**Other Indicators**
Starch indicator is used in oxidation-reduction titrations involving iodine. This is because iodine forms an intense blue to black colored complex in the presence of starch. Thus starch allows a very sensitive assessment of the presence of iodine in a solution.

It is important to add the starch indicator close to the end point when there is an acid present. The acid will cleave the starch and that will prevent the starch from working properly. Students using starch should use a pilot run to get an idea when to add the starch indicator.

**Preparation of Indicators**

- **Methyl orange (MO):** if you have a balance, weigh out about 1 g of methyl orange powder and dissolve it in about 1 L of water. Store the solution in a plastic water bottle with a screw on cap and it will keep for years. If it gets thick and cloudy, add a bit more water and shake. If you do not have a balance, add half of a small tea spoon to a liter of water.
- **Phenolphthalein (POP):** Dissolve about 0.2 g of phenolphthalein powder in 100 mL of pure ethanol; then add 100 mL water with constant stirring. If you use much more water than ethanol, solid phenolphthalein will precipitate. Store **POP** in a plastic water bottle with a screw on cap. We recommend making **POP** in smaller quantities than **MO** as it does not keep as well, mostly due to the evaporation of ethanol. If the solution develops a precipitate, add a bit of ethanol and shake. We do not recommend using purple methylated spirits as a source of ethanol for making **POP**. You can distill purple spirits to make clear spirits. For clear methylated spirits, use 140 mL of spirit and 60ml of water, as spirits generally are already 30% water.
- **Starch:** place about 1 g of starch in 10 mL of water in a test tube. Mix well. Pour this suspension into 100 mL of boiling water and continue to boil for one minute or so. Alternatively, use the water leftover after boiling pasta or potatoes. If this is too concentrated, dilute it with regular water.
- *Note:* The authors have never prepared bromothymol blue or universal indicator from powder, but suspect their preparation is similar to methyl orange.

Note that the exact mass of indicator used is not very important. You just need to use enough so that the color is clearly visible. Students use very little indicator in each titration, and a liter of indicator solution should last you a long time.

The Volumetric Analysis practical consists of an acid that is being titrated acid against a base until neutralization, in order to determine the concentration of the base.

On NECTA practical exams, titrations are done four times: a pilot followed by three trials. The pilot is done quickly and is used to determine the approximate volume needed for neutralization to speed up the following trials.

Ex: If the pilot gives an end point of 25.00 mL, then for the three subsequent trials, 20.00 mL can quickly be added from the burette. Then begin to add solution slowly until the endpoint is reached.

Results from the pilot are not accurate and are not included when doing calculations. Students should also know that not all three trials are always used in calculating the average volume used. Values of trials must be consistent and within ± 02 cm$$^3$$ of each other to be valid for average volume determination.

**Volumetric Analysis Using Burettes Preparation**

1. After washing Burettes thoroughly, rinse the Burette with 3 mL of the acidic solution that will be used during the titration (Acid usually goes in the Burette).
- Cover the entire inside surface of the Burette.
- Discard 3 mL of solution properly when finished
- Why? This prevents dilution of acid by water.

2. After washing the flask thoroughly, rinse the flask with 3 mL of solution that will be used during the titration (Base usually goes in the flask).
- Cover the entire inside surface of the flask.
- Discard 3 mL of solution properly when finished.

**Procedure**

1. Clean the burette with water.
2. Rise the burette with the acid that will be used for the titration.
3. Fill the burette with the acid. Let a little run out through the stopcock.
4. Record the initial burette reading.
5. Use a syringe to transfer the base solution into a conical flask.
6. Record the volume moved by the syringe.
7. If you are using an indicator, add a few drops to the flask.
8. Slowly add the acid from the burette to the flask. Swirl the flask as you titrate. Be careful. Avoid acid drops landing on the sides of the flask.
9. Stop titration when the slight color change become permanent. This is the end point.
10. Record final reading of the burette.
11. Repeat for remaining titrations.

**Notes**
Burettes tell you the volume of solution used, not the volume present.

- Ex: Initial Reading - 4.23 mL
- Final Reading - 20.57 mL
- You used 16.34 mL of acid during the titration.

- Always read burettes at eye level.
- Always read from the bottom of the meniscus. In a plastic apparatus, there is often no meniscus.
- Burettes are accurate to 2 decimal places. Students should estimate to the nearest 0.01 mL
- For Acid-Base indicators: The less indicator used, the better. To change color, the indicator must react with fluid in the burette. If you add too much, it uses more chemical than necessary for neutralization, creating an indicator error.
- For starch indicators: use 1 mL. Starch is not titrated; indicators are, and you must use more to get a good color change.

**Volumetric Analysis Without Using Burettes**

Use plastic syringes instead of burettes.

As of late 2010, the most precise syringes available are the 10 mL NeoJect brand - you should use these (A titration with 2 plastic syringes is more accurate than a titration with a burette and a cheap glass pipette).

If use of these syringes is new to you, read Use of a Plastic Syringe to Measure Volume (p. 58) before continuing.

If students are using syringes in place of burettes, they require two syringes for the practical: one to use as a burette (for acid) and one to use as a pipette to transfer base into the flask. It may be useful to label the different syringes “burette” /“flask” or “acid”/“base”.

**Preparation (without burettes)**

1. Clean the “pipette” syringe with water.
2. Rinse the “pipette” syringe with base solution that will be put into the flask.
3. Use the “pipette” syringe to transfer base into the flask. To do this accurately, first add 1 mL of air to the syringe and then suck up the base beyond the desired amount. Push back the plunger until the top of the fluid is at the required volume.
4. Record the total volume transferred (multiple transfers with the 1 syringe may be required to react the desired volume).
5. If you are using indicator, add a few drops to the flask.
6. Clean the “burette” syringe with water.
7. Rinse the “burette” syringe with the acid solution that will be used for titration.

**Procedure (without burettes)**

1. Add 1 mL of air to the syringe and suck up the acid beyond the 10 mL mark. Slowly push back the plunger until the top of the fluid is exactly at the 10 mL line.
2. Slowly add acid from the “burette” syringe into the flask. Swirl the flask as you titrate. Be careful. Make sure the acid lands in the base, avoid acid drops landing on the sides of the flask.
3. Stop titration when the slight color change become permanent. This is the end point.
4. Often, more than 10 mL of acid will need to be used. This is not a problem. Once 10 mL is finished in the syringe, students should just fill it up again and continue the titration.
5. Record final volume of acid transferred by the “burette” syringe.

**Notes for when using syringes in place of burettes**

- Students must record their results in a manner that is consistent with traditional reporting.
- On rough paper, students should calculate the volume of solutions used during titration. If they only used one syringe and the initial volume in the syringe was 10.00 mL and the final volume was 2.55 mL, the student used 7.45 mL of solution. If they used two full syringes and then part of a third (which had the initial reading of 10.00 mL and a final reading of 4.65 mL), the student used 5.35 mL + 10.00 mL + 10.00 mL = 25.35 mL total.
- In the table of results, the student should write 25.35 mL for Volume Used. If they had used a burette, the initial reading would have been 0.00 mL and the final reading would have been 25.35 mL. This is what they should write in their table of results.
- When using a syringe as a burette, students should write 0.00 mL as the Initial Volume and then, for the Final Volume, they should write the number they calculated for the total volume used.

###27.2.5 Common Calculations in Titration Experiments
All NECTA practical experiments require students to determine some unknown in the titration procedure. Common calculations that the problem statement will ask for include:

- Concentration (molarity) of an acid or base
- Relative atomic mass of unknown elements in an acid or base
- Percentage purity of a substance
- Amount of water of crystallization in a substance

**Concentration of an Acid or Base**
The problem statement may have the student find either the unknown molarity (moles per litre) or concentration (grams per litre) of the acid or the base. As an example, the following steps are used to calculate the unknown concentration of an acid:

1. _Calculate the average volume of acid used._
Remember to not use the pilot trial or any trials that are not within ± 0.2 cm$$^3$$ of each other.
2. _Calculate the number of moles of the base used._
$$\text{Molarity} = \frac{\text{number of moles}}{\text{volume of solution}}$$
These values can usually be taken from the solutions listed on the test paper. Also be sure that the units of volume of solution are in litres or dm$$^3$$.

3. _Write a balanced chemical equation for the reaction._
The chemical equation can also be written as an ionic equation.

4. _Calculate the number of moles of acid used from the mole ratio taken from the balanced chemical equation._
Both ionic and full formulae equations give the same mole ratio.

5. _Work out the molar concentration of the acid._

The molar concentration can be determined using the calculated number of moles of acid (found in the previous step) and the average volume of acid used (found in step 1), using the equation in step 2.

Alternatively, the following equation can be used:
$$\cfrac{C_AV_A}{C_BV_B} = \cfrac{n_A}{n_B}$$
where:

C$$_A$$ is the molar concentration of the acid.
V$$_A$$ is the volume of the acid used.
n$$_A$$ is the number of moles of the acid used.
C$$_B$$ is the molar concentration of the base.
V$$_B$$ is the volume of the base used.
n$$_B$$ is the number of moles of the base used.

Similar steps are used to calculate the unknown concentration of a base. Repeat steps 1 through 5, but with the following changes:

- Step 2: Calculate the moles of the acid used.
- Step 4: Calculate the moles of the base from the mole ratio.
- Step 5: Find the molar concentration of the base, either using the molarity calculation or the equation above.

**Relative Atomic Mass of Unknown Elements**

Atomic mass of unknown elements, as well as molecular mass of compounds with unknown elements may need to be calculated in the problem statement. Most unknown elements will be a metal of a basic compound. As an example, the following steps are used to calculate the relative atomic mass of an unknown metal element of a metal carbonate:

- _Calculate the average volume of acid used._
Remember to not use the pilot trial or any trials that are not within ± 0.2 cm$$^3$$ of each other.
- _Calculate the number of moles of the acid used._
These values can usually be taken from the solutions listed on the test paper. Also be sure that the units of volume of solution are in litres or dm$$^3$$.
- _Write a balanced chemical equation for the reaction to get the mole ratio._
- _Determine the number of moles of the metal carbonate used._ This can be taken from the balanced chemical equation.
- _Work out the molecular concentration of the metal carbonate solution._ Use the formula as shown in step 2.
- _Calculate the mass of the metal carbonate in one litre of solution._ This can be done using the following ratio:

$$\frac{\text{mass given in problem statement}}{\text{volume given in problem statement}} = \frac{\text{mass of unknown metal}}{\text{one litre}}$$

Make sure the units correspond because sometimes the problem statement will be expressed in dm$$^3$$ or cm$$^3$$.

- Using the molarity of the solution and the mass of the metal carbonate per litre of solution, work out the relative molecular mass of the metal carbonate.

The following equation can be used to calculate molar mass:

$$\text{molar mass} = \frac{\text{mass per litre}}{\text{molarity}}$$

- Calculate the relative atomic mass of the metal based on the formula of the carbonate.

Use the total molar mass of the compound found in step 7 and the molar mass of each element in the compound to find the molar mass of the unknown element. Some problem statements may require the student to identify the unknown element from its molecular mass.

Similar steps should be followed if the unknown element is of an acidic compound. Just replace the steps that include the metal carbonate solution with the acid solution.

**Percentage Purity of a Substance**

Problem statements that require the student to find percentage purity will usually contain one solution in the list provided that specifically states it is impure or that it is a hydrated compound (seems very low in concentration). Again, it is possible to determine percentage purity of an acid or a base. As an example, the following steps are used to calculate the percentage purity of a base:

- _Determine the average volume of the acid used._
Remember to not use the pilot trial or any trials that are not within ± 0.2 cm$$^3$$ of each other.
- _Calculate the number of moles of the acid used._

$$\text{Molarity} = \frac{\text{number of moles}}{\text{volume of solution}}$$

These values can usually be taken from the solutions listed on the test paper. Also be sure that the units of volume of solution are in litres or dm$$^3$$.

- _Write a balanced chemical equation for the reaction to get the mole ratio._
- _Determine the number of moles of base used in the reaction._ This can be taken from the mole ratio from the previous step.
- _Calculate the mass of the base used in the reaction._

The mass can be determined by the number of moles calculated and the following relationship:

$$\text{mass} = \text{number of moles} \times \text{molar mass}$$

- _Work out the percentage purity of the base solution sample._
The following equation for percentage purity should be used:

$$\text{percentage purity} = \frac{\text{mass of pure substance in sample}}{\text{mass of the impure sample}} \times 100\%$$

It is very important to note that when calculating percentage purity, the amount of volume in the concentration of base must be equal to the volume of concentration of acid used. For example, if there was 0.424 g of sodium carbonate in 25 cm$$^3$$ of solution reacting with a 250 cm$$^3$$ solution of acid, the mass of sodium carbonate must be converted to know the mass in 250 cm$$^3$$. Therefore, 250 cm$$^3$$ of base solution will contain 4.24 g, not 0.424 g.

The value for the mass of the impure sample comes from the list of provided solutions and the mass of the pure sample will come from the calculations.

Similar steps can be followed to find the percentage purity of a the acid solution sample. Instead of finding the mass of the base, use the calculated moles of acid used to find the mass of acid in the actual reaction.

**Amount of Water of Crystallization**

Water of crystallization is the water that is bound within crystals of substances. Most hydrated substances and solutions contain water of crystallization. Problem statements that ask students to determine the amount of water of crystallization will have a solution with a formula similar to [base] · XH$$_2$$O, and they have to solve for X. As an example, the following steps are used to determine the number of molecules of water of crystallization in a hydrated base compound sample:

- _Calculate the average volume of the acid used._
Remember to not use the pilot trial or any trials that are not within ± 0.2 cm$$^3$$ of each other.
- _Calculate the number of moles of the acid used._

$$\text{Molarity} = \frac{\text{number of moles}}{\text{volume of solution}}$$

These values can usually be taken from the solutions listed on the test paper. Also be sure that the units of volume of solution are in litres or dm$$^3$$.

- _Write a balanced chemical equation for the reaction to get the mole ratio._
- _Calculate the number of moles of the base used._
This can be determined from the mole ratio in the previous step.
- _Determine the molar concentration of the base._

The molarity can be calculated using the volume of base used in the experiment and the equation from step 2.

- _Calculate the relative molecular mass (R.M.M.) of the base compound._ The following equation can be used to calculate molar mass:

$$\text{molar mass} = \frac{\text{mass per litre}}{\text{molarity}}$$

- _Determine the number of molecules of water of crystallization in the sample._

Using the relative atomic masses of the various atoms in the base compound, subtract the mass of the compound from the total mass of the hydrated compound. Water molecules always have a total molecular mass of 18 $$^\text{g}/_\text{mol}$$, so the remaining mass will be composed of multiples of 18. For example, if a hydrated carbonate (Na$$_2$$CO$$_3$$·XH$$_2$$O) has a total mass of 286 g, the molecules of water can be determined as follows:

$$2\text{Na} + \text{C} + 3\text{O} + x(2\text{H} + \text{O}) = 286$$
$$(2 \times 23) + 12 + (3 \times 16) + x[(2 \times 1) + 16] = 286$$
$$106 + 18x = 286$$
$$18x = 180$$
$$x = 10$$

Therefore, in this example, there are 10 molecules of water of crystallization in the hydrated sodium carbonate (Na$$_2$$CO$$_3$$ · 10H$$_2$$O) sample.

###27.2.6 Sample Practical Question

The following is a sample practical question from 2012. You are provided with the following solution:

TZ: Containing 3.5 g of impure sulphuric acid in 500 cm$$^3$$ of solution; LO: Containing 4 g of sodium hydroxide in 1000 cm$$^3$$ of solution; Phenolphthalein and Methyl indicators.

**Questions:**

(A)

1. What is the suitable indicator for the titration of the given solutions? Give a reason for your answer.
2. Write a balanced chemical equation for the reaction between TZ and LO.
3. Why is it important to swirl or shake the contents of the flask during the addition of the acid?

(B)

1. Titrate the acid (in a burette) against the base (in a conical flask) using two drops of your indicator and obtain three titre values.

(C)

1. ________________ cm3 of acid required and ________________ cm3 of base for complete reaction.
2. Showing your procedures clearly, calculate the percentage purity of TZ. (20 marks)

**Discussion**

This practical question requires students to know and understand how to use volumetric analysis apparatus and technique. Since this question involves the titration of sulphuric acid (strong acid) and sodium hydroxide (strong base), either phenolphthalein or methyl orange are acceptable indicators to use. An explanation of suitable indicators can be found in Acid-base Indicators (p. 98).

Make sure that students create a table for the first pilot titration and three titre values, for a total of four titrations. Only the titre values (not the pilot) that are within ±0.02 ml of each other will be used to calculate the average titrated volume. Students should also be swirling the contents of the volumetric flask in order to thoroughly mix the acid and base together. The titration is complete only when there is permanent color change in the indicator.

Note that although this procedure states the number of drops of indicator and how many number of titre values, it does not indicate what volume to use in the flask. The typical volume is 25 ml, but students can use any volume as long as they are consistent for each trial.

The practical question for volumetric analysis will always ask students to either determine percentage purity, molecules of crystallization of water, unknown concentration of one of the solutions, or molar mass of one of the solutions. See Common Calculations in Titration Experiments (p. 102) for more explanation on various volumetric analysis calculations.

##27.3 Qualitative Analysis

This section contains the following:

- Overview of Qualitative Analysis
- Local Materials in Qualitative Analysis
- The Steps of Qualitative Analysis
- Hazards and Cleanliness
- Sample Practical: Preparation of Copper Carbonate for Qualitative Analysis

###27.3.1 Overview of Qualitative Analysis

The salts requiring identification have one cation and one anion. Generally, these are identified separately although often knowing one helps interpret the results of tests for the other. For ordinary level in Tanzania, students are confronted with binary salts made from the following ions:

- Cations: NH$$^{+4}$$, Ca$$^{2+}$$, Fe$$^{2+}$$, Fe$$^{3+}$$, Cu$$^{2+}$$, Zn$$^{2+}$$, Pb$$^{2+}$$, Na$$^{+}$$
- Anions: CO$$^{2−}$$, HCO$$^{−}$$, NO$$^{-}$$, SO$$^{2−}$$, Cl$$^{−}$$

At present, ordinary level students receive only one salt at a time. The teacher may also make use of qualitative analysis to identify unlabeled salts.

The ions are identified by following a series of ten steps, divided into three stages. These are:

- Preliminary tests: These tests use the solid salt. They are: appearance, action of heat, action of dilute H$$_2$$SO$$_4$$, action of concentrated H$$_2$$SO$$_4$$, flame test, and solubility.
- Tests in solution: The compound should be dissolved in water before carrying out these tests. If it is not soluble in water, use dilute acid (ideally HNO$$_3$$) to dissolve the compound. The tests in solution involve addition of NaOH and NH$$_3$$.
- Confirmatory tests: These tests confirm the conclusions students draw from the previous steps. By the time your students start the confirmatory tests, they should have a good idea which cation and which anion are present. Have students do one confirmatory test for the cation they believe is present, and one for the anion you believe is present. Even if several confirmatory tests are listed, students only need to do one. When identifying an unlabelled container, however, you might be moved to try several, especially if you are new to this process.

###27.3.2 Local Materials in Qualitative Analysis

For all low-cost local material substitutes, consult the section on Sources of Laboratory Equipment (p. 208).

- **Heat Sources:** Motopoa burners cost nothing to make (soda bottle caps) and consume only a small amount of fuel. They give a non-luminous flame ideal for flame tests and still produce enough heat for the other tests.
- **Test Tubes:** Plastic test tubes suffice.
- **Litmus paper:** Rosella flowers give very good results.
- **Low-cost sources of chemicals:** (see Shika Express Chemistry companion manual).

Share expensive chemicals among many schools. A single container of potassium ferrocyanide, for example, can supply ten or even twenty schools for several years. Schools should consider bartering 10 g of one chemical for 10 g of another. Another alternative is for all of the schools in a district or town to pool money to buy one container of each required imported reagent, and then divide the chemicals evenly.

###27.3.3 The Steps of Qualitative Analysis

####Appearance

Three properties of the salt may be observed directly: colour, texture, and smell.

**Colour:** While most salts are white, salts of transition metals are often colored. Thus colour is an easy way to identify iron and copper cations in salts.

**Texture:** Carbonates and hydrogen carbonates generally form powders although sometimes they can form crystals. Sulphate, nitrates, and chlorides are almost always founds as crystals.

**Smell:** Some ammonium salts smell distinctly like ammonia. Some, however, have no smell. Therefore the smell of ammonia can confirm the presence of ammonium cations, but its absent can not be used to prove the absence of ammonium.

*Materials:* Soda bottle caps, table salt, bicarbonate of soda, soda ash (sodium carbonate), copper (II) sulphate, ammonium sulphate, locally manufactured iron (II) sulphate, locally manufactured iron (III) sulphate, locally manufactured copper (II) carbonate

**Preparation**

1. Place a small amount of each sample in a different soda bottle cap for observation.

**Activity Steps**

1. Look at the samples. Describe their colour, texture, and smell. Do not touch or inhale the salts.

**Results and Conclusion**

- **Colour**
*White:* Copper and iron absent
*Blue Copper* cation present
*Green:* Iron (II) or copper present
*Light Green:* Iron (II) present
*Yellow or Red-Brown:* Iron (III) present

- **Texture**
*Powder:* Carbonate or hydrogen carbonate anion present
*Crystals:* Sulphate, chloride, or nitrate anion probably present
*Wet Crystals:* Chloride or nitrate anion present

- **Smell**
*Smell of ammonia:* Ammonium cation present
*No smell of ammonia:* Inconclusive – some ammonium compounds have no smell

**Clean Up**

1. Collect salts for use another day. Do not mix.
2. Wash and return soda bottle caps.

*Notes* Wet crystals are the result of the salt absorbing water from the atmosphere. Qualitative analysis salts with this property are not locally available. However, caustic soda (sodium hydroxide) has this property, so samples of caustic soda can be used to show the absorption of water from the air and how this changes the appearance of the salt. Note that caustic soda burns skin, blinds in eyes and corrodes metal, so care is required.

####Action of heat

Many salts thermally decompose when heated. When these salts decompose, they produce gases that may be identified to identify the anion of the salt. After decomposition, many salts also leave a residue that may identify the cation.

Materials soda bottle caps, motopoa, matches, long handled metal spoons, steel wool, sand, beaker, water, table salt, copper (II) sulphate, bicarbonate of soda, locally prepared copper (II) carbonate*, soda ash (sodium carbonate), locally prepared zinc carbonate

**Hazards and Safety**

- Ammonium nitrate explodes when heated. For this reason, ammonium nitrate should never be used in qualitative analysis when the Action of Heat test is used.
- Test tubes should be pointed away from the student holding them and from other students by holding them at an angle. This will prevent injuries due to splashing chemicals, and will also minimize inhalation of any gases produced. Teach students to never to fill test tubes or any other container more than half.

**Preparation**

1. Fill a beaker with water.
2. Make a small pile of sand on the table for resting the hot spoon.
3. Place a small amount of each sample in a different soda bottle cap.
4. Add motopoa to another soda bottle cap to use as a burner.

**Activity Steps**

1. Light the motopoa. Note that the flame will be invisible.
2. Place a very small amount of a sample on the spoon. Generally, the smallest amounts of sample give the best results because they are easier to heat to a hotter temperature.
3. Heat the sample strongly, observing all changes.
4. Place the hot spoon on the sand to cool.
5. Once the spoon has mostly cooled, dip it in the beaker of water to remove the rest of the heat.
6. Use the steel wool to remove all residue from the spoon.
7. Repeat these steps with each sample.

**Results and Conclusion**

- **Gas released**
- *Brown gas* Nitrogen dioxide, nitrates present, confirmed
- *Colourless* gas with smell of ammonia Ammonia, ammonium present, confirmed
- *Colourless* gas with no smell Very likely carbon dioxide, especially if the compound decomposes near the start of heating, carbonate or hydrogen carbonate present
- *No change* Salt probably a chloride, sulphate (very high temperatures are required to decompose many sulphates), or sodium carbonate
- **Residue**
- *No residue* Ammonium cation present
- *Black residue* Copper cation probably present
- *Red residue* when hot, dark when cool Iron cation present Yellow residue when hot, white when cool Zinc cation present Red residue when hot, yellow when cool Lead cation present
- **Sound**
- *Cracking sound* Sodium chloride or lead nitrate present

**Clean Up**

1. Thoroughly remove all residues from the spoons.

Notes Sodium carbonate is the only carbonate used in qualitative analysis that does not thermally decompose. Therefore a white powder that does not decompose when heated is probably sodium carbonate.

####Action of dilute H$$_2$$SO$$_4$$:

Carbonates and hydrogen carbonates react with dilute acid. Sulphates, chlorides and nitrates do not. Therefore reaction with dilute acid is useful test to help identify the anion. Sulphuric acid is used because it is the least expensive.

**Materials:** dilute sulphuric acid, droppers*, bicarbonate of soda, table salt

**Hazards and Safety**

- Use only a few drops of acid. These are all that are necessary and using more can be dangerous.

**Preparation**

1. Place a small amount of each sample in a different soda bottle cap.
2. Fill droppers with 1-2 mL dilute acid.

**Activity Steps**

1. Add a few drops of acid to each sample. Observe the results.

**Results and Conclusion**

Bubbles of gas Carbon dioxide produces; carbonate or hydrogen carbonate anion present No bubbles of gas Carbonate and hydrogen carbonate absent

**Clean Up**

1. Neutralize spills of dilute sulphuric acid with bicarbonate of soda.
2. Mix the remains from the reactions together so the extra bicarbonate of soda can neutralize the acid used to test table salt. Dilute the resulting mixture with a large amount of water and dispose down a sink, into a waste storage tank, or into a pit latrine.

**Notes** You can confirm that the gas produced is carbon dioxide by testing to see if it extinguishes a glowing splint. To do this, light a match, use about 0.5 mL of acid (rather than a few drops), and see if the gas released will extinguish the match.

####Action of concentrated H$$_2$$SO$$_4$$:

Concentrated sulphuric acid can convert chloride anions to hydrogen chloride gas and some nitrates to nitrogen dioxide. Because both of these gases are easy to detect, the addition of concentrated acid is used to distinguish between nitrates, chlorides, and sulphates. The concentrated acid used in this experiment should be about 5 M, similar to battery acid.

**Materials** battery acid, droppers, spoons, test tubes, test tube rack, test tube holder, heat source, hot water bath, table salt (sodium chloride), gypsum (calcium sulphate), ammonium sulphate, blue litmus paper, beaker, water

**Hazards and Safety**

- Use battery acid or another source of 5 M sulphuric acid for this experiment. Do not use fully concentrated 18 M sulphuric acid directly from either industry or laboratory supply. 18 M is too concentrated and very dangerous to use.
- Concentrate acid reacts violently with carbonates and hydrogen carbonates. The previous test – the addition of dilute acid – will detect carbonates and hydrogen carbonates. If that test is positive, do not test the sample with concentrated sulphuric acid.
- Test tubes should be pointed away from the student holding them and from other students by holding them at an angle. This will prevent injuries due to splashing chemicals, and will also minimize inhalation of any gases produced. Teach students to never to fill test tubes or any other container more than half.

**Preparation**

1. Place a small amount of each sample in a different soda bottle cap.
2. Add about 1 mL of air to each dropper syringe (no needle!) and then 2 mL of battery acid. Distribute the dropper syringes in the test tube racks so they stand with the outlet pointing down. The goal is to prevent the battery acid from reacting with the rubber plunger.

**Activity Steps**

1. Light the heat source and start heating the hot water bath. The water in the hot water bath should boil.
2. Use the spoon to add a small amount of a sample to a test tube.
3. Add two drops of battery acid to the sample to make sure there is no violent reaction.
4. Add just enough battery acid to cover the sample. Avoid spilling drops of acid on the inside walls of the test tube.
5. If a brown gas is released, stop at this step.
6. Moisten the blue litmus paper by quickly dipping it in the water of the hot water bath.
7. Place the litmus paper over the mouth of the test tube to receive any gases produces. If the litmus paper changes colour, stop at this step.
8. Hold the test tube in the hot water bath and heat for a while. Stop heating before the acid in the test tube boils. If the litmus paper changes colour before the acid boils, this is a useful result. If the acid boils, fumes from the acid itself will change the colour of the litmus paper – this result is not useful, and acid fumes are dangerous.

**Results and Conclusion**

**Bubbles with a few drops of acid:** Carbonate or hydrogen carbonate anion present
**Brown gas produced:** Nitrate anion present
**Litmus changes to red:** Hydrogen chloride gas produced; chloride anion present
**No effect observed:** Sulphate anion probably present

**Clean Up**

1. Fill a large beaker half way with room temperature water. This will be the waste beaker.
2. Pour waste from the test tubes into the waste beaker.
3. Fill each test tube half way with water and add this water to the waste beaker.
4. Return unused battery acid from the droppers to a well-labelled storage container for future use. Immediately fill each dropper (syringe) with water and transfer this water to the waste beaker.
5. Slowly add bicarbonate of soda to the waste beaker until addition no longer causes bubbling. This is to neutralize the acid in the waste.
6. Dilute the resulting mixture with a large amount of water and dispose down a sink, into a waste storage tank, or into a pit latrine.
7. Thoroughly wash all apparatus, including the test tubes and droppers, and return them to the proper places.

####Flame test

Some metal ions produce a characteristically coloured flame when added to fire.
Materials soda bottle caps, motopoa, metal spoons, beaker, steel wool, water, table salt (sodium chloride), gypsum (calcium sulphate), copper (II) sulphate, ammonium sulphate

**Preparation**

1. Fill a beaker with water.
2. Place a small amount of each sample in a different soda bottle cap.
3. Add motopoa to another soda bottle cap to use as a burner.

**Activity Steps**

1. Light the motopoa. Note that the flame will be invisible.
2. Place a small amount of sample on the edge of the spoon. For some spoons, it is better to hold the spoon by the wide part and to place the sample on the end of the handle.
3. Hold the sample into the hottest part of the flame, 1-2 cm above the motopoa. If necessary, tilt the spoon so that the sample touches the flame directly. Do not spill the sample into the flame.
4. Dip the hot end of the spoon into the beaker of water to cool it and remove the sample. If necessary, clean the spoon with steel wool.
5. Repeat these steps with each sample.

**Results and Conclusion**

- *Blue or green flame:* Copper present, confirmed
- *Golden yellow flame:* Sodium present, confirmed
- *Brick red flame:* Calcium present
- *Bluish white flame:* Lead present
- *No flame colour:* Copper and sodium absent; calcium and lead probably absent; cation is probably ammonia, iron, or zinc

**Clean Up**

1. Collect unused samples for use another day.
2. Wash and return all apparatus.

####Solubility

**Materials** soda bottle caps, two spoons, test tubes, test tube rack, hot water bath, heat source, distilled (rain) water, table salt (sodium chloride), soda ash (sodium carbonate), gypsum (calcium sulphate), powdered coral rock (calcium carbonate) or locally manufactured calcium carbonate or locally manufactured copper (II) carbonate

**Preparation**

1. Fill a beaker with water.
2. Place a small amount of each sample in a different soda bottle cap.

**Activity Steps**

1. Light the heat source and start heating the hot water bath. The water in the hot water bath should boil.
2. Decide which spoon will be used for transferring samples and which will be used for stirring.
3. Use the transfer spoon to transfer a very small amount of a sample to a test tube.
4. Add 3-5 mL of distilled water to the test tube.
5. Use the handle of the stirring spoon to thoroughly mix the contents of the test tube.
6. If the sample does not dissolve, heat the test tube in the water bath until the contents of the test tube are almost boiling (small bubbles rise from the bottom). Mix.
7. Repeat these steps with each sample.

**Results and Conclusion**

*Sample dissolves in room temperature water:* Soluble salt present
*Sample dissolves only in hot water:* Calcium sulphate or lead chloride present
*Sample does not dissolve in even hot water:* Insoluble salt present

**Solubility Rules**

- All Group I (sodium, potassium, etc) and ammonium salts are soluble (sodium borate is an exception but not relevant to qualitative analysis)
- All nitrates and hydrogen carbonates are soluble
- Most chlorides are soluble (silver and lead chlorides are exceptions, although the latter is soluble in hot water)
- Carbonates of metals outside of Group I are generally insoluble (note that aluminum and iron (III) carbonate do not exist)
- Lead sulphate is insoluble and calcium sulphate is soluble only in hot water. Magnesium sulphate is completely soluble while sulphates of the Group II metals heavier than calcium (strontium and barium) are insoluble. All other sulphates used in qualitative analysis are soluble]

**Table of Solubility for Qualitative Analysis**

| | Ammonium | Sodium | Copper | Iron | Zinc | Calcium | Lead |
| ----| -------- | ------ | ------ | ---- | ---- | ------- | ---- |
| Nitrate | O | O | O | O | O | O | O |
| Chloride | O | O | O | O | O | O | ∆ |
| Sulphate | O | O | O | O | O | ∆ | X |
| Carbonate | O | O | X | X | X | X | X |
| Hydrogen Carbonate | O | O | - | - | - | - | - |

KEY:

- O = soluble at room temperature
- ∆ = soluble only when heated
- X = insoluble in water
- – = salt does not exist

**Clean Up**

1. Collect all unused (dry) samples for use another day.
2. Unless copper carbonate is used, none of the salts listed in the materials section of this activity are harmful to the environment.
3. Dispose of solutions in a sink, waste tank, or pit latrine.
4. Dispose of solids and liquid wastes with precipitates in a waste tank or pit latrine – never dispose of solids in sinks.
5. If using copper carbonate, collect all waste containing copper carbonate and filter to recover the copper carbonate. Save for use another day.
6. If you do this activity with a lead nitrate or lead chloride, collect these wastes in a separate container. Add dilute sulphuric acid dropwise until no further precipitation is observed. Neutralize with bicarbonate of soda. Dispose this mixture in a waste tank or a pit latrine. The lead sulphate precipitate is highly insoluble will not enter the environment.
7. Wash and return all apparatus.

**Notes** Calcium carbonate or copper carbonate are recommended qualitative analysis salts to use as examples of insoluble salts. If these are difficult to get, other insoluble compounds may be used for teaching this specific step (but not for other parts of qualitative analysis). Examples of other insoluble compounds include sulphur power, manganese (IV) oxide from batteries, and chokaa (calcium hydroxide, which is only slightly soluble so a significant precipitate will remain).

**Materials** soda bottle caps, two spoons, test tubes, test tube rack, beakers, medium droppers (5 mL syringes without needles), large droppers (10 mL syringes without needles), caustic soda (sodium hydroxide), table salt (sodium chloride), ammonium sulphate, copper (II) sulphate, locally manufactured iron (II) sulphate, locally manufactured iron (III) sulphate, locally manufactured zinc sulphate, distilled (rain) water

**Preparation**

1. Fill a 500 mL water bottle about half way with distilled (rain) water.
2. Add one level tea spoon of caustic soda and then wash the spoon.
3. Label the bottle “1 M sodium hydroxide – corrosive”
4. Place a small amount of each sample in a different soda bottle cap.
5. Pour some of the sodium hydroxide solution into a clean beaker.
6. For each small dropper syringe, suck in about 1 mL of air and then add about 4 mL of sodium hydroxide solution. Distribute the dropper syringes in the test tube racks so they stand with the outlet pointing down. The goal is to prevent the sodium hydroxide from reacting with the rubber plunger.

**Activity Steps**

1. Decide which spoon will be used for transferring samples and which will be used for stirring.
2. Use the transfer spoon to transfer a very small amount of a sample to a test tube.
3. Use the large dropper syringe to add 3-5 mL of distilled water to the test tube.
4. Use the handle of the stirring spoon to thoroughly mix the contents of the test tube.
5. Use the small dropper to add a few drops of sodium hydroxide solution to the test tube.
6. Observe the colour of any precipitate formed. Also waft the air from the top of the test tube towards your nose to test for smell.
7. If a white precipitate forms, use the stir spoon to transfer a very small quantity of the precipitate to a clean test tube. Add 1-2 mL of sodium hydroxide directly to this sample to see if the precipitate is soluble in excess sodium hydroxide solution.

**Results and Conclusion**
*No precipitate and smell of ammonia:* Ammonium cation present, confirmed
*No precipitate and no smell:* Sodium cation probably present
*Blue precipitate:* Copper (II) cation present
*Green precipitate:* Iron (II) cation present
*Red-brown precipitate:* Iron (III) cation present
*White precipitate not soluble in excess NaOH:* Calcium cation present
*White precipitate:* soluble in excess NaOH Lead or zinc cation present

**Clean Up**

1. Save all waste from this experiment, labeling it “basic qualitative analysis waste, no heavy metals” and leave it in an open container. Over time atmospheric carbon dioxide will react with the sodium hydroxide to make less harmful carbonates. After 2-3 days, dispose of the waste in a waste tank or a pit latrine.

####Addition of NH$$_3$$ solution:

This test is very similar to the addition of sodium hydroxide solution. The useful difference is that zinc forms a precipitate in ammonia that is soluble in excess ammonia whereas lead forms a precipitate in ammonia that is not soluble in excess ammonia. Therefore, this test is mainly used to separate lead and zinc. Neither lead salts nor ammonia are locally available in Tanzania. Because the process of this test is the same as the addition of NaOH and the results so similar, students can adequately learn about the Addition of NH$$^3$$ test by practicing the Addition of NaOH. For the national exam, a small amount of ammonia solution can be obtained.

Note also that the addition of ammonia to a solution of copper (II) will produce a blue precipitate that dissolves in excess ammonia to form a deep blue solution. This is a useful conformation of the presence of copper, but such conformation is generally unnecessary because the flame test for copper is so reliable.

If you have ammonia solution, store it in a well-sealed container to prevent the ammonia from escaping. A good container for this is a well labeled plastic water bottle with a screw on cap.

####Confirmatory Tests
Every cation and anion has at least one specific test that can be used to prove its presence. Not all of these tests are possible with local materials, but many of them are. The following list shows how to confirm each possible cation and anion.

####Confirmatory Tests for the Cation

**Ammonium**

- Example salt: ammonium sulphate
- Procedure: add sodium hydroxide solution and heat in a water bath
- Confirming result: smell of ammonia
- Reagents: NaOH solution as used above

**Calcium**

- Example salt: calcium sulphate
- Procedure: Two options
1. flame test
- Confirming results:
1. flame test: brick red flame
2. addition of NaOH: white precipitate insoluble in excess
- Reagents:
1. none
2. NaOH solution

**Copper**

- Example salt: copper sulphate
- Procedure: flame test
- Confirming result: blue/green flame
- Reagents: none

**Iron (II)**

- Example salt: locally manufactured iron sulphate (keep away from water and air)
- Procedure: addition of sodium hydroxide solution and then transfer of precipitate to the table surface
- Confirming result: green precipitate that oxidizes to brown when exposed to air
- Reagent: sodium hydroxide solution from above

**Iron (III)**

- Example salt: locally manufactured iron sulphate (oxidized by water and air)
- Procedure: addition of sodium ethanoate solution
- Confirming result: yellow to red solution
- Reagent: slowly add bicarbonate of soda to vinegar; stop adding when further addition does not cause bubbles; label the solution “sodium ethanoate for detection of iron (III)”

- Example salt: no local sources for safe manufacture, consider purchasing lead nitrate
- Procedure: Three options
1. flame test
2. addition of dilute sulphuric acid
3. addition of potassium iodide solution
- Confirming results:
1. flame test: blue/white flame
2. addition of dilute sulphuric acid: white precipitate
3. addition of KI solution: yellow precipitate that dissolves when heated and reforms when cold
- Reagents:
1. none but a very hot flame, e.g. Bunsen burner, is required
2. dilute sulphuric acid used in Step 5 above
3. obtain pure potassium iodide by evaporating iodine tincture until only white crystals remain; do this outside and do not breathe the fumes; it might also be possible to use the KI solution prepared for electrolysis in the chapter on ionic theory

**Sodium**

- Example salts: sodium chloride, sodium carbonate, sodium hydrogen carbonate • Procedure: flame test
- Confirming result: golden yellow flame
- Reagents: none

**Zinc**

- Example salt: locally manufactured zinc carbonate or zinc sulphate
- Procedure: addition of 0.1 M potassium ferrocyanide solution
- Confirming result: gelatinous gray precipitate
- Reagents: no local source of potassium ferrocyanide – consider collaborating with many schools to share a container; only a very small quantity is required

####Confirmatory Tests for the Anion

**Hydrogen carbonate**

- Example salt: sodium hydrogen carbonate
- Procedure: add magnesium sulphate solution and then boil in a water bath
- Confirming result: white precipitate forms only after boiling
- Reagent: dissolve Epsom salts (magnesium sulphate) in distilled (rain) water*

**Carbonate**

- Example salt: sodium carbonate
- Procedure for soluble salts: addition of magnesium sulphate solution
- Confirming result: white precipitate forming in cold solution
- Reagent: dissolve Epsom salts (magnesium sulphate) in distilled (rain) water
- Note that insoluble salts that effervesce with dilute acid are likely carbonates. None of the other anions described here produce gas with dilute acid. Note also that all hydrogen carbonates are soluble.

**Chloride**

- Example salt: sodium chloride
- Procedure: Three Options
1. addition of silver nitrate solution
2. addition of manganese (IV) oxide and concentrated sulphuric acid followed by heating in a water bath
3. addition of weak acidified potassium permanganate solution followed by heating in a water bath
- Confirming results:
1. silver nitrate: white precipitate of silver chloride
2. manganese (IV) oxide: production of chlorine gas that bleaches litmus
3. acidified permanganate: decolourization of permanganate
- Reagents:
1. Silver nitrate has no local source but may be shared among many schools as only a very small amount is required.
2. Manganese dioxide may be purified from used batteries and battery acid is concentrated sulphuric acid. Note that careful purification is required to remove all chlorides from the battery powder. This method is useful because of its low cost, but remember that chlorine gas is poisonous! Students should use very little sample salt in this test.
3. Prepare a solution of potassium permanganate, dilute with distilled water until the colour is light pink, and then add about 1 percent of the solution’s volume in battery acid. Note that this solution will cause lead to precipitate, and will also be decolourized by iron II, so it is not a perfect substitute for silver nitrate. This final option is also not yet recognized by examination boards, i.e. NECTA

**Sulphate**

- Example salt: copper sulphate, calcium sulphate, iron sulphate
- Procedure: addition of a few drops of a solution of lead nitrate, barium nitrate, or barium chloride
- Confirming result: white precipitate
- Reagents: none of these chemicals have local sources. Because lead nitrate is also an example salt, it is the most useful and the best to buy. The ideal strategy is to share one of these chemicals among many schools. Remember that all are quite toxic.

**Notes** Emphasize to students that they need to carry out only one confirmatory test for the cation, and one for the anion. If the test gives the expected result, then they can be sure that the ion they have identified is present. If the test does not give the expected result, they have probably made a mistake, and they should revisit the results of their previous tests and choose a different possibility to confirm.

###27.3.4 Hazards and Cleanliness

Qualitative analysis practicals are full of hazards, from open flames to concentrated acids. Here are some ways to reduce the risk of accidents:

- Teach students how to use their flame source before the day of the practical.
- Have students hold their test tubes at an angle pointed away from them and other students to prevent splashing chemicals and minimize inhalation of any gases produced.
- Teach students never to fill test tubes or any other container more than halfway in order to minimize spills and boiling over of chemicals during heating.
- Teach students that if they get chemicals on their hands, they should wash them off immediately, without asking for permission first.
- Teach students to tell you immediately when chemicals are spilled. Sometimes they hide chemical spills for fear of punishment. Do not punish them for spills – legitimate accidents happen. Do punish them for unsafe behavior of any kind, even if it does not result in an accident.
- Practicals involving nitrates, chlorides, ammonium compounds, and some sulphates produce harmful gases. Open the lab windows to maximize airflow.
- Make absolutely sure that students clean their tables and glassware before they leave.

###27.3.5 Sample Practical Question

The following is a sample practical question from 2012.
Substance V is a simple salt which contains one cation and one anion. Carry our the experiments described below. Record carefully your observations and make appropriate inferences and hence identify the anion and cation present in sample V.
￼￼￼
(TABLE)

**Conclusion**

1. The cation in sample V is:
2. The anion in sample V is:
3. The chemical formula of V is:
4. The name of compound V is:

**Discussion**

This particular example was to identify calcium carbonate; however, the above procedure follows the same format commonly used for other unknown salts. The only differences may be the specific solutions used in some of the steps.

At times, the procedure may not be explicit or indicated whatsoever and the student is required to write the detailed procedure in addition to the observations and inferences.

Emphasize to students that in addition to the qualitative analysis procedure they have to do only one confirmatory test for cations and one for anions.

##27.4 Chemical Kinetics and Equilibrium

###27.4.1 Theory

Compared to the other two NECTA chemistry practicals - Acid/Base Titration (i.e. Volumetric Analysis) and Qualitative Analysis - Chemical Kinetics has few alternative chemicals that can be used.

The chemical reaction in the NECTA exam is a precipitation of sulphur.

(equation)

The preparation and procedure for Chemical Kinetics is very simple.

###27.4.2 Preparation

1. Prepare solution of Sodium Thiosulphate by placing appropriate mass of crystals (in accordance with desired concentration) in a water bottle and shake vigorously until completely dissolved.
2. Prepare acid solution using hydrochloric acid (preferable) or sulphuric acid. (Remember for calculations that sulphuric acid will have two H$$^+$$ ions to hydrochloric acid’s one.)
3. Take a white piece of paper and draw a large black X on it

###27.4.3 Procedure
The procedure can change depending on whether the variable of the reaction rate is temperature or concentration. The procedure outlined here will be focused on concentration. The sample NECTA question below is based on having a temperature variable.

1. Place in a beaker or glass the amount of sodium thiosulphate solution and clean water prescribed in the table.
2. Ready the stopwatch and pour in the prescribed amount of hydrochloric acid, starting the stopwatch as you do so.
3. Gently swirl the glass and place it over the paper with an X on it.
4. When the solution has precipitated enough sulphur to the point where the X is no longer visible, stop the stopwatch. Record the time.
5. The resulting solution is already neutralized, and can be merely diluted and disposed of as is. (Unless you wish to obtain the salt/sulphur mixture by evaporation for another day.)
6. Thoroughly rinse the glass and repeat, changing the amounts of sodium thiosulphate and water as prescribed by the problem.
7. Graph the results to calculated the desired variables

This practical is consistent and easy to practice, but it requires sodium thiosulphate which can be expensive and hard to get a hold of. (As mentioned, the hydrochloric acid can be replaced with sulphuric (battery) acid.) An alternative reaction to demonstrate chemical kinetics (that can be performed very cheaply) is the iodization of acetone, seen in the reaction below.

CH$$_3$$COCH$$_3$$(aq) + I$$_2$$ → CH$$_3$$COCH$$_2$$I (aq) + H$$^+$$(aq) + I$$^−$$(aq)

This reaction starts as a dark opaque solution and eventually proceeds to a colorless, transparent solution. It requires an acidic environment to occur (both hydrochloric or sulphuric acid suffice). The amount of time it takes to reach the point of colorlessness varies depending on the concentration of the acetone. It is easy to demonstrate the relationship between concentration and rate of reaction. (By varying the temperature or amount of acid catalyst the reaction visibly proceeds at differing rates.)

###27.4.4 Performing the Practical

**Materials**
7 beakers, 3 syringes, stopwatch

**Chemicals**
Nail polish remover, iodine tincture, sulphuric (battery) acid, water

**Preparation**

1. Prepare an iodine solution using iodine tincture. Solutions purchased in the local drugstores are often 0.2 M Iodine (with several other chemicals as well). Create a 0.02 M solution by adding 9 parts water to one part tincture. Put this solution in the first beaker.
2. Prepare an acidic acetone solution by mixing one part nail polish remover to one part 1 M Sulphuric acid solution. Put this solution in the second beaker. This is roughly a 5 M solution of acetone.
3. In the third beaker place clean water.

**Procedure**

1. Pour 8 mL of the acetone solution in a clean beaker. Clear a stopwatch and add 8 mL of iodine to the beaker, swirl the solution, and stop the watch when the solution becomes colorless.
2. Record the value and then perform the experiment again. This time start with 6 mL of acetone solution and 2 mL of clean water in a clean beaker. Clear a stopwatch and add 8 mL of iodine to the beaker, swirl the solution, and stop the watch when the solution becomes colorless.
3. Repeat the above steps with 4 mL of acetone solution, 4 mL of water, and 8 mL of iodine, and then finally one more time with 2 mL of acetone solution, 6 mL of water and 8 mL of iodine.

| Test | Volume of Iodine Solution (mL) | Volume of Acetone Solution (mL) | Volume of Water (mL) | Molarity of Acetone Solution | Time for Color to Disappear | Reciprocal of Time |
| ---- | ---- | ----| ---- | ---- | ---- | ----------|
| 1 | 8 | 8 | 0 | 5 M | | |
| 2 | 8 | 6 | 2 | 3.75 M | | |
| 3 | 8 | 4 | 4 | 2.5 M | | |
| 4 | 8 | 2 | 6 | 1.25 M | | |

This data can then be used to plot a graph of concentration of acetone against the rate of reaction.

Notes

- Patience is required for the tests using lower concentrations as they can take over 4 minute to complete.
- Concentrations may vary depending on where the tincture and remover are purchased.
- The endpoint of this reaction can be somewhat ambiguous depending on the color of the nail polish remover. Criteria for determining the endpoint may vary.
- It should be noted that if the nail polish remover is already a specific color it will affect the final color of the solution. Some solutions may never become fully colorless.
- The reaction used in the NECTA exams goes from transparent to opaque while this alternative goes from opaque to transparent. Make sure students understand this difference.
- The reaction used in NECTA exams is a neutralization reaction so there is little that needs to be done to process the waste. This alternative reaction is very acidic when finished so be prepared to neutralize it before disposal.

###27.4.5 Sample Practical Question

The following is a sample practical question from 2012. Your are provided with the following materials:

- **ZO:** A solution of 0.13 M Na$$_2$$S$$_2$$O$$_3$$ (sodium thiosulphate)
- **UU:** A solution of 2M HCl
- Thermometer
- Heat source/burner
- Stopwatch

Procedure:

1. Place 500 cm$$^3$$ beaker, which is half-filled with water, on the heat source as a water bath.
2. Measure 10 cm$$^3$$ of **ZO** and 10 cm$$^3$$ of **UU** into two separate test tubes.
3. Put the two test tubes containing **ZO** and **UU** solutions into a water bath.
4. When the solutions attain a temperature of 60°C, remove the test tubes from the water bath and pour both solutions into 100 cm$$^3$$ empty beaker and immediately start the stop watch.
5. Place the beaker with the contents on top of a piece of paper marked **X**.
6. Note the time taken for the mark **X** to disappear.
7. Repeat step (i) to (vi) at temperature 70°C, 80°C and 90°C.
8. Record your results as in Table 1.

**Table 1**

|**Experiment** | **Temperature** | **Time** |
| ------------- | --------------- | -------- |
| 1 | 60 | |
| 2 | 70 | |
| 3 | 80 | |
| 4 | 90 | |

**Questions:**

1. Write a balanced chemical equation for reaction between **UU** and **ZO**.
2. What is the product which causes the solution to cloud the letter **X**?
3. Plot a graph of temperature against time (s).
4. What conclusion can you draw from you graph?

**Discussion**

This particular example was to investigate how temperature affects the rate of a chemical reaction. Other experiments for chemical kinetics involve concentration, surface area, and a catalyst; however, the most common problem statement is how concentration affects the rate of reaction.

For all scenarios, make sure to tell students to start the stopwatch immediately after combining the reactants.
chapter28
#Chapter 28: Physics Practicals

##28.1 Introduction to Physics Practicals

###28.1.1 Format

The format of the Physics practical exam was revised in 2011 to keep up with the 2007 updated syllabus. As such, there will be no further Alternative to Practical exams, pending approval from the Ministry of Education. Prior to 2011, there were 3 questions in the Physics practical, and students chose 2 questions to do.

As of now, the Physics practical has 2 questions and students must answer both. Question 1 comes from Mechanics, and Question 2 can come from Heat, Light, Waves or Electricity. Each question is worth 25 marks, and students have 2 and a half hours to complete the exam.

**Physics 1 Theory Format**

The theory portion of the Physics exam comprises 100 marks, while the practical carries 50 marks. A student’s final grade for Physics is thus found by taking her total marks from both exams out of 150.

The theory exam for Physics contains 3 sections. Section A has 3 questions and is worth 30 marks - Question 1 is 10 multiple choice, Question 2 is 10 matching, and Question 3 is 10 fill-in-the-blank. Section B has 6 long answer questions and is worth 60 marks. Section C has 2 questions regarding the use of apparatus and simple technological appliances in everyday life, though students must answer only 1 of these questions. It is worth 10 marks.

*Note* This information is current as of the time of publication of this manual. Updated information may be obtained by contacting the Ministry of Education.

###28.1.2 Notes for Teachers

There is only one set of advance instructions for the Physics practical. Advance instructions are given to teachers at least one month before the date of the exam. Unlike Biology and Chemistry, there are no 24 hour advance instructions given for Physics. The advance instructions will state exactly how many and which apparatus candidates should be provided with for each practical question.

The Physics practical is different from the Chemistry and Biology practicals in that the exams feature a greater variety of questions. That means we need to teach it all, even if the teacher before you never found his or her way into the classroom and you realize at the end of Form Four that the students still have not studied the Form Two syllabus. If you are teaching all forms, do not wait to start practicals until later forms; always do a practical when the corresponding topic comes up. In addition, train the students well in the general principles of collecting data, graphing data, and writing up experimental points. These skills are required in every physics practical, and carry most of the points.

The practical section of the exam is a third of a student’s total score, and fully half of that is graphing and labeling experimental data. Though the practical is varied, a student does not necessarily need a deep understanding of the concept in question. If they are familiar with the apparatus and the process of drawing and interpreting a graph, the practical should be quite simple. Whenever possible, allow the students to play and experiment with the apparatus, whether it is a metre bridge, mirror, pendulum, etc. If they have done each of these experiments several times, they will be confident in their ability.

The more familiar your students are with these techniques, the better they will do. Perform these practicals as often as possible: when the topic comes up, when preparing for the mock and NECTA exams, and any time you can get them to come in for an evening session or a weekend. They will make many, many mistakes the first couple of times through but that is exactly what you want as they will learn from their mistakes and remember them.

###28.1.3 Common Practicals
**Mathematics** A brief overview of some of the mathematical and graphing skills required to perform many of the common physics practicals

**Mechanics**

- **Hooke’s Law (Form 1)** use a spring and various masses to determine either the spring constant or the value of an unknown mass graphically
- **Simple Pendulum (Form 2)** find the acceleration due to gravity using a pendulum and stop- watch
- **Principle of Moments (Form 2)** verify the Principle of Moments using masses and a ruler on a knife edge, or calculate the mass of a metre rule

**Light**

- **Plane Mirror (Reflection) (Form 1)** generally 3 varieties: find image distance, verify the Laws of Reflection, or find the number of images produced by two plane mirrors placed at different relative angles
- **Rectangular Prism (Refraction) (Form 3)** find the refractive index or critical angle for a glass block by varying the angle of incidence and measuring the corresponding angles of refraction

**Electricity**

- **Potentiometers** find the drop in potential along a length of resistance wire
- **Metre Bridges** determine the value of an unknown resistor using a known resistor and a galvanometer to find a point of equal potential along a resistance wire
- **Ohm’s Law (Form 2)** verify Ohm’s Law or determine the internal resistance of a cell

*Note* These are the most common practicals, but they are not necessarily the only practicals that can occur on a NECTA exam. Physics practical questions can come from a variety of topics which may not yet have been used in older past papers. Be sure to regularly check the most recent Physics Past Papers to get a good idea of the types of questions to expect.

###28.1.4 Recent Practicals
Given below is an attempt to characterize the Physics practical questions from recent years according to category and topic. Each question is given with its corresponding topic on top and the objective, or what is to be solved for, on bottom. This information can be used to try and find trends in what exam writers like to test students on, and what topics are most likely to occur on an exam. However, nothing should be assumed to be a guarantee, and students should be well-prepared in all practicals so that they can take on any question they may face on the NECTA.

Note a few things from the table below:

- Beginning in 2011, the format of the exam changed to consist of only 2 problems, one of which must come from Mechanics.
- The exam committee has a tendency to repeat some problems. For example, the Mechanics questions from 2004 and 2011 are nearly identical, as are the Light questions from 2004 and 2008.

| **Year** | **Mechanics** | **Light** | **Electricity** |
|----------|---------------|-----------|-----------------|
| 2013 | | | |
| 2012 | | | |
| 2011 | | | |
| 2010 | | | |
| 2009 | | | |
| 2008 | | | |
| 2007 | | | |
| 2006 | | | |
| 2005 | | | |
| 2004 | | | |

##28.2 Mathematics

No physics experiment is complete without a healthy dose of graphing and formulas. As math is typically the worst subject for most students, it is often upon the physics teacher to drive home the understanding of how to draw and interpret graphs, as well as how to apply formulas to those graphs. It comes down to a few simple things: correctly setting up a graph (scales, units, labels, etc.), plotting points from a table of data, and fitting a best-fit line. After this, the students need to find the slope of this line and its y-intercept.

###28.2.1 Graphing

Most of the graphs will be linear, meaning the slope is constant, so we apply the standard equation for a line

$$y = mx + c$$

where *y* represents the vertical axis, *x* represents the horizontal axis, *m* is the slope of the line and *c* is the point on the vertical axis where the line crosses. Almost every practical will make use of this equation, so be sure that your students understand it inside and out. It often helps to do repetitive practice using just the mathematical symbols before introducing physics concepts. Note that very rarely non-linear graphs appear, e.g. cooling curves in heat practicals. In this case students will not have to find a mathematical relationship, just describe and explain the trend in the data.

###28.2.2 Formulas

Now comes the physics; all the practicals will involve an equation that can be rewritten in this linear form. The exam question will dictate which variable is independent (*x*) and which is dependent (*y*). It is up to the student to simply rewrite the formula with each variable on its respective side and then infer what *m*, the slope, and *c*, the *y*-intercept, must be.
The most common formulas used for mechanics, light and electricity are as follows:

| Hooke's Law | |
| -------------------- | --------------- |
| Period of a Pendulum | |
| Principle of Moments | |
| Snell's Law | |
| Ohm's Law | |
| Resistance of a Wire | |
| Wheatstone Bridge | |

In each case, one quantity will be changed (independent) and another will be measured (dependent) over the course of the experiment. The student will therefore need to rearrange the equation so that the dependent variable is the subject in the form

$$y = mx + c$$

**Example Problem - Snell’s Law**

As an example, in an experiment to measure the index of refraction of a glass block, a student will be measuring angles of incidence and refraction. This means we need to use Snell’s Law

$$n_1 \times \sin{i} = n_2 \times \sin{r}$$

The question will typically ask students to plot a graph of their measurements, with sin *i* on the *y*-axis and sin *r* on the *x*-axis, or vice-versa. To rewrite Snell’s law in the form of *y = mx + c* is simple; we get

$$\sin{i} = \sin{r} \times \frac{n_2}{n_1}$$

and we can see that the value corresponding to *m* is the ratio

$$\frac{n_1}{n_2}\$$

and *c* must be zero.

Since we are trying to find *n$$_2$$* (the refractive index of glass), and we know *n$$_1$$* is 1.0, we simply measure the slope and solve to find *n$$_2$$*.

The approach itself is relatively simple, but students will need lots of practice with graphing, rewriting equations in linear form, and determining what corresponds to *m* and *c* in each case. The same approach is used to find the quantities in each of the equations above.

A complete list of each equation in its most commonly found *y = mx + c* form, along with its corresponding dependent variable *y*, independent variable *x*, slope *m* and *y*-intercept *c*, is given below. Note that these equations are not always used in the given form on practicals. It is up to the student to determine how each equation must be analyzed during an exam. The variables used and methods of graphing change from year to year, and so the following table should by no means be memorized or assumed to be applicable for a given problem.

| **Name of Law** | **Equation** | **$$y=mx+c$$** | **$$y$$**| **$$x$$**| **$$m$$** | **$$c$$** |
| ---------------- | ------------ | ------------ | ---------- | ------------- | ------------- | ------------- |
| Period of a Pendulum | | | | | | |
| Principle of Moments | | | | | | |
| Snell's Law | | | | | | |
| Ohm's Law | | | | | | |
| Resistance of a Wire | | | | | | |
| Wheatstone Bridge | | | | | | |

###28.2.3 Units
Be sure to always stress the importance of units when performing practicals, especially in graphing. Using the wrong units can lead to inaccurate interpretations of graphical data. For example, when using Hooke’s Law, a spring constant is typically given in units of $$^{\text{N}}/_{\text{m}}$$. But if a student is graphing mass (in g) against extension (in cm), then the slope will be in units of $$^{\text{g}}/_{\text{cm}}$$. In order to get units of $$^{\text{N}}/_{\text{m}}$$, one would have to first convert the slope into units of $$^{\text{kg}}/_{\text{m}}$$ and then multiply by the acceleration of gravity (g = 10 $$^{\text{N}}/_{\text{kg}}$$). A problem may or may not ask for specific units in its answers, but regardless, students should always be conscious of what units they are using when making calculations.

Thinking of units can also help students to understand a problem they are struggling with. If they can remember that slope is change in *y* over change in *x*, then they may be able to deduce the meaning of the slope of a graph by looking at the units of the *y* and *x* axes. The most important part of any experiment, though, is following directions. If a student can follow directions, which usually are clearly provided by the exam, and can graph data, they can easily perform any experiment. If anything, the practical exam is a test in a student’s ability to follow instructions.

##28.3 Mechanics
The mechanics section is mandatory on every exam and typically falls into three categories: Hooke’s Law, Simple Pendulum, and Principle of Moments. However, other topics are possible, as evidenced by a question on Archimedes’ Principle on the 2012 exam. These experiments use the following materials:

- Metre Rules If unavailable, go to a local fundi to mass produce them.
- Masses See Sources of Laboratory Equipment for local varieties.
- Springs Can be bought at lab stores, or can use substitutes such as rubber bands or strips of elastic from a tailor.
- Retort Stands May be available or can be made using a filled 1.5 L water bottle with a bamboo stick taped at the top and extending to one side.
- Eureka Can Cut off the bottom of a 500 mL water bottle and cut a slit at the top that can be folded downward to make a curved spout.

###28.3.1 Hooke’s Law
This is the most common practical, usually involving a spring but sometimes a rubber band or piece of string. This experiment can be tricky simply because NECTA likes to switch it up every year; try to give your students as much practice with different variations. It is likely that NECTA will require a spring of known spring constant, and you will need known masses. Either can be bought at a laboratory supply store in town, but it is possible to make your own. The practical is simple to perform, but there are some common mistakes: be sure the students understand that the extension is the change in length, not the ultimate length shown on the ruler. Also, do not confuse mass and weight, as is common.
An example question from the 2007 NECTA is shown below. After reviewing the topic with your students, let them try this on their own. You will need to repeat it several times before they are comfortable, using different springs and masses each time.

**Sample Practical Question**
The aim of this experiment is to determine the mass of a given object “B”, and the constant of the spring provided.

![2007-1-alt.png](images/2007-1-alt.png)

1. Set up the apparatus as shown in the figure with zero mark of the metre-rule at the top of the rule and record the scale reading by the pointer, S$$_0$$.
2. Place the object “B” and standard weight (mass) W equal to 20 g in the pan and record the new pointer reading S$$_1$$. Calculate the extension, e = S$$_1$$-S$$_0$$ in cm.
3. Repeat the procedure in (ii) above with W = 40 g, 60 g, 80 g and 100 g.

(a) Record your results in tabular form as shown below: Table of Results:

| S$$_O$$= | | | |
| --------- | ------------ | ---------------------------- | ---------------------------------- |
| Mass (kg) | Force, F (N) | Pointer reading S$$_1$$ (cm) | Extension = S$$_1$$ - S$$_O$$ (cm) |
| 0 | | | |
| 0.02 | | | |
| 0.04 | | | |
| 0.06 | | | |
| 0.08 | | | |
| 0.10 | | | |

(b) Plot graph of Force F (vertical axis) against extension e (horizontal axis).

(c) Use your graph to evaluate

1. mass of B
2. spring constant, K, given that force, extension, constant and weight of B are related as follows:

$$F = Ke - B$$

**Discussion**
This practical has two parts: the first is to find the spring constant *k*, the second is to find the mass of an unknown object *B*. By looking at the equation above, we can see that *F* is the dependent variable, *e* is the independent variable, *K* is the slope and *- B* is the intercept. When the graph is drawn, *K* and *B* can be found easily. Note that the intercept on the graph will be negative.

The procedure is simply to start from a certain point on the metre rule (it does not need to be a specific number) and to add masses one at a time, measuring the distance from your starting point to the new position. This distance is called the extension, e. Be sure that you are not simply reading the metre rule, but are measuring the distance from the starting point.

###28.3.2 Simple Pendulum
With some practice, this experiment should be simple for anyone to perform. The trick comes with the math and graphing (again, an example is shown below). The materials can all be local (string, stones, ruler) except for the stopwatches (for which you should consult the materials section).
The practical usually has one objective: to find the acceleration due to gravity, *g*. We know that the mass of a pendulum and its angle of deflection (for small angles) do not affect its period. Therefore we vary only the length *L* of the pendulum and measure its period, as shown in the following example question.

**Sample Practical Question**
The aim of this experiment is to determine the magnitude of the acceleration due to gravity, *g*. Proceed as follows:

![pendulum.png](images/pendulum.png)

1. Make a simple pendulum by suspending a weight on a string 10 cm long from a retort stand.
2. Allow the pendulum to swing for twenty oscillations, using a stopwatch to record the time. Repeat this procedure for pendulum lengths of 20 cm, 30 cm, 40 cm, and 50 cm.
3. Record your results in tabular form as shown below

| Pendulum Length (m) | Time for 20 oscillations (s) | Period (equation) | (equation) |
| ------------------- | ---------------------------- | ----------------- | ---------- |
| 0.1 | | | |
| 0.2 | | | |
| 0.3 | | | |
| 0.4 | | | |
| 0.5 | | | |

4. Plot a graph of T$$^2$$ (vertical axis) against Pendulum Length (horizontal axis).
5. Calculate the slope of the graph.
6. Use the slope to calculate the value of *g*.
7. What are possible sources of error in this experiment?

**Discussion**
The period of a pendulum can be calculated using

$$T = 2\pi\sqrt{\frac{l}{g}}$$

where *l* is the length of the pendulum, *T* is the period, and *g* is acceleration due to gravity. By squaring both sides, we get a much easier equation to graph:

$$T^2 = 4\pi\frac{l}{g}$$

In this equation we see that T$$^2$$ is the dependent variable (*y*-axis) and *l* is the independent variable (*x*-axis), so the slope must be

$$\mathrm{slope} = \frac{4\pi}{g}$$

When the graph is complete, the value of *g* can be calculated easily.
Many students are confused by the difference between the time for many oscillations and the period, which is the time for one oscillation. Be sure that they can change between the two easily.

Note that pendulum practicals do not always require students to find *g*. Sometimes they are just required to find the relationship between *l* and *t*. Again, it is essential that students read and understand the examination question, rather than memorize past solutions, and that they have lots of practice in collecting, organizing, and graphing data from a variety of experiments.

###28.3.3 Principle of Moments
This experiment is used to verify the Principle of Moments, or equilibrium, by balancing a meter rule on a knife-edge with masses at various distances. For this experiment, students need a solid understanding of the Center of Gravity, the Moment of a force, and equilibrium. Questions can range from finding the mass of an object to asking for the mass of the metre rule. They are all variations on the same practical: using the condition of equilibrium to find mass.

The following example is from the 2011 NECTA exam and asks students to find the mass of a battery using the Principle of Moments. Following is a brief explanation of the alternative practical of finding the mass of a metre rule.

####Finding the Mass of an Object

**Sample Practical Question**
The aim of this experiment is to determine the mass of a given dry cell size “AA”. Proceed as follows:

1. Locate and note the centre of gravity *C* of the metre rule by balancing it on the knife edge.
2. Suspend the 50 g mass at length ‘*a*’ cm on one side of the metre rule and the 20 g mass together with the dry cell at length ‘*b*’ cm on the other side of the metre rule. Fix the 50 g mass at length 30 cm from the fulcrum and adjust the position of the 20 g mass together with the dry cell until the metre rule balances horizontally. Read and record the values of *a* and *b* as *a$$O$$* and *b*$$0$$* respectively.
3. Draw the diagram for this experiment.
4. By fixing a = 5 cm from fulcrum *C*, find its corresponding length *b*.
5. Repeat the procedure in number 4 above for *a* = 10 cm, 15 cm, 20 cm and 25 cm. Tabulate your results.
6. Draw a graph of ‘*a*’ against ‘*b*’ and calculate its slope *G*.
7. Calculate *X* from the equation $$50 = \cfrac{b_0}{a_0}(20 + X)$$
8. Comment on the value of $$cfrac{b_0}{a_0}$$
9. State the principle governing this experiment.

**Discussion**

This practical utilizes the Principle of Moments to find the mass of a “AA” battery. Initially, a known mass of 50 g is balanced with the (battery + 20 g mass) system. Note that ‘*a*’ and ‘*b*’ are measured from the fulcrum and so students should be careful not to just read the cm mark on the ruler where each object is suspended.
Also note that students are required to actually find the centre of gravity *C* of the ruler rather than assuming it to be the 50 cm mark. This measured value of *C* is to be used as the starting point for all future measurements of *a* and *b*.

In part 7, students should recognize the equation $$50 = \cfrac{b_0}{a_0}(20 + X)$$ as coming from the Principle of Moments. Starting with

$$F_{\mathrm{clockwise}} \times d_{\mathrm{clockwise}} = F_{\mathrm{anticlockwise}} \times d_{\mathrm{anticlockwise}}$$

we get

$$(mg)_{\mathrm{clockwise}} \times d_{\mathrm{clockwise}} = (mg)_{\mathrm{anticlockwise}} \times d_{\mathrm{anticlockwise}}$$

or

$$(50 \text{g})(g) \times a_0 \text{ cm} = (20 \text{g} + X \text{g})(g) \times b_0 \text{ cm}$$

where $$cfrac{b_0}/{a_0}$$ is the ratio of the lever arm distances for the two weights being used. If the mass of the battery X is less than 30 g, this ratio should be greater than 1, but if the mass is greater than 30 g, the ratio should be less than 1.

**Finding the Mass of a Metre Rule**
This question is less frequently seen on NECTA exams as compared to finding an unknown mass. How- ever, it utilizes the same principles of equilibrium and balancing moments, and therefore is a useful alternative practical to ensure that students understand the concept rather than memorizing solutions to one version of the problem.
The mass of a uniform solid object, like a metre rule, is assumed to be at the center of the object. In the case of the metre rule, we can say that the center of mass is at the 50 cm mark, directly in the center. If we want it to be in equilibrium, the moments on either side of a pivot must be equal, or

$$\mathrm{\text{Clockwise moment}} = \mathrm{\text{Anticlockwise moment}}$$

To find the mass of the metre rule itself, we begin by placing a known mass at one point on the metre rule. We then move the pivot to one side or another until the metre rule is perfectly balanced in equilibrium. As shown in the diagram below, the pivot will not be at the 50 cm mark.

![meter-rule.png](images/meter-rule.png)

If the metre rule is in equilibrium, we know that the moments must be equal, or that

$$F_{\mathrm{clockwise}} \times d_{\mathrm{clockwise}} = F_{\mathrm{anticlockwise}} \times d_{\mathrm{anticlockwise}}$$

In this case, the anticlockwise force is the weight of the object, and the distance is that from the pivot to the object. The clockwise force is the weight of the metre rule, and the distance is that from the 50 cm mark (center of mass) to the pivot. Therefore our equation is:

$$W_{\mathrm{rule}} \times d_{\mathrm{rule}} = W_{\mathrm{object}} \times d_{\mathrm{object}}$$

Because the weight of the object is known, and the two distances can be measured, we can easily calculate the mass and therefore the weight of the metre rule:

$$W_{\mathrm{rule}} = \frac{W_{\mathrm{object}} \times d_{\mathrm{object}}}{d_{\mathrm{rule}}}$$

From this we can calculate mass of the metre rule using *F = mg*.

##28.4 Light

The light practical typically involves plane mirrors or glass blocks (rectangular prisms). Presumably you will have already done these practicals with the students in Form three (refraction) and Form 1 (plane mirrors), but a little practice will make the theory and execution clear, especially if they can work in groups. The materials you will need are as follows:

**Cork Board** Use cardboard for this, about 0.5 to 0.75 cm thick.
**Optical pins** Use sewing pins or syringe needles. If using syringe needs, be sure to crimp the ends so students do not prick themselves.
**Protractors** These are cheap and students are supposed to have them anyway. Small ones come in local mathematical sets.
**Glass Block / Rectangular Prism** A simple rectangular piece of 6 mm glass, about 8 cm by 12 cm, will work.
**Plane Mirror** You can buy mirror glass in town in small sections for 200/= or less; it should be available in villages through the local craftsmen if they work on windows. Alternately, you can smoke one side of a piece of glass to make the other side like a mirror.

###28.4.1 Plane Mirror (Reflection)
These are not as common as the rectangular prism, but they come in a variety of questions:

- Placing pins in front of a mirror at different distances and finding the distance of the image.
- Verifying the Law of Reflection at plane mirrors.
- Placing two mirrors at different relative angles to find the number of images produced.

These are not overly complicated, but you should definitely practice with your students creating images in mirrors – they are not as accustomed to playing with mirrors as you might be.
Given below is an example practical from the 2006 NECTA exam which asks students to find a relationship between object distance and image distance in a plane mirror.

**Sample Practical Question**
Set up the experiment as shown in the diagram below using plane mirror, soft board, three pins and a white sheet of paper.

![2006-2-alt.png](images/2006-2-alt.png)

Fix a white sheet of paper on the soft board. Draw a line across the width at about the middle of the white sheet (**MP**). Draw line **ONI** perpendicular to **MP**.
Fix optical pin **O** to make **ON** = **U** = 3 cm. By using plasticine or otherwise, fix plane mirror along portion of **MP** with **O** in front of the mirror. With convenient position of eye, **E**, look into the mirror and fix optical pins **A** and **B** to be in line with image, **I**, of pin **O**.
Measure and record **NI = V**. Repeat procedure for **U** = 6 cm, 9 cm and 12 cm.

(a) Tabulate your results as follows:

| U (cm) | 3 | 6 | 9 | 12 |
| ------ | ---- | ---- | ---- | ---- |
| V (cm) | | | | |

(b) Plot graph of **U** against **V**.
(c) Calculate slope, *m* of the graph to the nearest whole number.
(d) State relationship between **U** and **V**.
(e) Write equation connecting **U** and **V** using numerical value of *m* with symbols **U** and **V**.
(f) From your equation give position of the image when object is touching the face of the mirror.

**Discussion**

For a plane mirror, object distance and image distance are equal. That is, **U** and **V** should be approximately equal values for this practical. Note that students should extend line **ONI** far behind the mirror since they don’t know where exactly the image **I** is going to be. The location of **I** is found at the intersection of the extended line **ONI** and the extended line **AB** connecting the two optical pins.

From the graph, the slope should be found to be 1 after rounding to the nearest whole number. From this, we can see that **U = V**. When the object is touching the face of the mirror, the object distance **U** is 0, and so the image distance **V** will also be 0.

###28.4.2 Rectangular Prism (Refraction)
Students will be asked to find the refractive index and/or critical angle of the glass block by varying the angle of incidence i and measuring the corresponding angles of refraction r as described in the Mathematics section earlier. They will do this by placing two pins in front of the prism, which together form a ‘ray’ (the light ray), and then placing two more pins on the other side of the prism so that, when observed through the prism from either side, the four pins line up exactly. By drawing the lines that the pins make on the paper, the refracted ray inside the prism can be easily traced, and the refracted angle measured. An example question from the 2007 NECTA is given below.

**Sample Practical Question**
The aim of this experiment is to find the refractive index of a glass block. Proceed as follows:
Place the given glass block in the middle of the drawing paper on the drawing board. Draw lines along the upper and lower edges of the glass block. Remove the glass block and extend the lines you have drawn. Represent the ends of these line segments as SS$$_1$$ and TT$$_1$$. Draw the normal NN$$_1$$ to the parallel lines SS$$_1$$ and TT$$1$$ as shown in the figure below:

![2007-2a-alt.png](images/2007-2a-alt.png)

Draw five evenly spaced lines from O to represent incident rays at different angles of incidence (10°, 20°, 30°, 40°, and 50° from the normal). Replace the glass block carefully between SS$$_1$$ and TT$$_1$$. Stick two pins P$$_1$$ and P$$_2$$ as shown in the figure as far apart as possible along one of the lines drawn to represent an incident ray. Locate an emergent ray by looking through the block and stick pins P$$_3$$ and P$$_4$$ exactly in line with images I$$_1$$ and I$$_2$$ of pins P$$_1$$ and P$$_2$$. Draw the emergent ray and repeat the procedure for all the incident rays you have drawn. Finally draw in the corresponding refracted rays.

![2007-2b-alt.png](images/2007-2b-alt.png)

(a) Record the angles of incidence *i* and the measured corresponding angles of refraction *r* in a table. (Your table of results should include the values of sin *i* and sin *r*.)
(b) Plot the graph of sin *i* (vertical axis) against sin *r* (horizontal axis).
(c) Determine the slope of the graph.
(d) What is the refractive index of the glass block used?
(e) Mention any sources of errors in this experiment.

**Discussion**
In this experiment, pins are used to simulate a ray of light. If all of the pins are aligned as you look through the block, they act as a single ray. It takes practice to be able to align the pins while looking through the block, so practice often with your students.

Light slows down as it enters a denser medium, so in order to minimize the time required to pass through that medium, it changes direction until it moves back into its original medium. In this case, light is moving from air into glass and then back into air, so its direction changes while inside the glass, then returns to its original direction when passing back into air. This effect is called refraction and it depends on the nature of the media, in this case air and glass. Snell’s law gives us the relationship between the nature of the media and the resulting angles of incidence and refraction:

$$n_1 \times \sin{i} = n_2 \times \sin{r}$$

In this experiment, the incident angle *i* is being changed and the refracted angle *r* is being measured. The refractive index of medium 1 (air) is known as 1.0, so we can use these three to find the refractive index of medium 2 (glass). On the graph, sin *i* is the dependent variable and sin *r* is the independent variable, so the equation becomes

$$\sin{i} = \sin{r} \frac{n_2}{n_1}$$

In this case the slope must be $$frac{n_2}{n_1}$$

The refractive index of medium air is simply 1.0, so the slope is the refractive index of medium 2.
This practical is one of the easiest to perform with students because it does not require much preparation. Syringe needles should be readily available and glass blocks are cheap, so it is possible to have every students try this themselves many times before taking the exam.

**Finding the Critical Angle**
Some questions may ask students to find the critical angle of a glass block in addition to its refractive index. The relationship between critical angle, *C*, and refractive index, *n*, for a particular medium is given by

$$n = \frac{1}{\sin{C}}$$

or

$$\frac{\sin{i}}{\sin{r}} = \frac{1}{\sin{C}}$$

Thus a graph of sin i against sin r can be used to find the critical angle. However, take care to note that we must first take the reciprocal of the slope, i.e.

$$\sin{C} = \frac{1}{n}$$

This gives us sin*C*, so to get *C* by itself, we need to use mathematical tables. Turn to the page for Natural Sines and search the table for the 4-figure value you obtained above by taking $${1}\over{\text{slope}}$$ corresponding row gives the angle in degrees and the column gives the additional minutes of the angle.

For example, say we plot our graph of sin *i* (*y*-axis) against sin *r* (*x*-axis) and we calculate the slope to be 1.43. This is the refractive index of the glass block (since we can remember that glass has a refractive index of 1.5, we can do a quick mental check to make sure this makes sense). Then sinC = $${1}\over{1.43}$$ = 0.6993. From the mathematical tables, we get $$C = 44^\circ 22'$$ [6993 falls between 6984 (18′) and 6997 (24′), so we use the Mean Differences table to add on 4′ giving a total of 22′].

Note that the method of finding *n* and *C* changes if we are instead told to graph sin *r* (*y*-axis) against sin *i* (*x*-axis). Be sure to practice both versions with students to ensure their understanding.

##28.5 Electricity
This is by far the least attempted practical on the exam, but not because it is difficult. The electricity practical, if properly set up, is one of the easiest to perform. It can appear in many different forms but will typically involve a simple circuit and some kind of variable resistor in order to measure current or EMF for different resistances. The materials you will probably need are as follows:

**Connecting Wires** Use speaker wire; it is cheap and available in most villages and towns.
**Voltmeters, Ammeters, Galvanometers** This is unavoidable; you can get full digital multimeters in town for about 10,000/=, galvanometers can be found in any lab store or can be made using a compass and insulated copper wire.
**Batteries** Two to four D-size batteries should easily be enough for these experiments. Try to avoid Tiger brand if possible. Panasonic is highly superior in quality for roughly the same price.
**Resistance Wire** These are used to make small resistors for the metre bridge or potentiometer. The most common type of wire to use is nichrome, which can be found in a hardware or lab store. Steel will also work, though it is less resistant and therefore harder to measure.
**Metre Bridges** See the activity that describes the construction of a metre bridge and potentiometer. It is best to make both together as the construction is almost identical and both are used frequently.
**Variable Resistor (Rheostat)** This is optional as it is typically only used to set a level that can be easily read by the voltmeter. However, if you are using a multimeter, you can simply change the magnitude setting on the multimeter to account for unusually low or high resistances.
**Soldering Iron** Not required, but may be a good investment for making reliable battery connections. Using electrical tape can lead to inconsistencies. Check large towns.

###28.5.1 Potentiometers
This experiment is very simple but requires the correct materials, namely the meter bridge/potentiometer described above. A complete circuit is created with a switch (optional), power source, variable resistor and 1 m of bare resistance wire, all in series.

The potentiometer itself is simple to construct; all preparation is done by the teacher, so the student simply follows the instructions as shown in the following example from the 2007 NECTA.

**Sample Practical Question**
The aim of this experiment is to determine the potential fall along a uniform resistance wire carrying a steady current. Proceed as follows:

![2007-3-alt.png](images/2007-3-alt.png)

Connect up the circuit as shown in the figure. Adjust the rheostat so that when the sliding contact *J* is near *B* and the key is closed the voltmeter *V* indicates an almost full scale of deflection. Do not alter the rheostat again.
Close key *K* and make contact with *J*, so that *AJ = 10 cm*. Record the potential difference *V* volts between *A* and *J* as registered on the voltmeter.

Repeat this procedure for *AJ = 20 cm, 30 cm, 50 cm,* and *70 cm.*

(a) Tabulate your results for the values of *AJ* and *V*.
(b) Plot a graph of *V* (vertical axis) against *AJ* (horizontal axis).
(c) Calculate the slope of the graph.
(d) What is your comment on the slope?
(e) State any precautions on the experiment.

**Discussion**
This is a simple test of the relationship between the length of a wire and its resistance, which we know is

$$R=\frac{\rho l}{A}$$

Where *l* is the length of the wire, *ρ* is the resistivity of the wire, and *A* is the cross-sectional area of the wire. We expect that as the length of wire increases, its potential difference will also increase. This is because the resistance (and therefore potential difference) of a wire is directly related to its length. The voltmeter in this experiment is measuring just the potential difference over the length of wire (10 cm, 20 cm, etc.), so if we use Ohm’s Law to say that *V = IR*, we can write:

$$V = \frac{I \rho l}{A}$$

In this experiment, *I*, *ρ* and *A* are all constant, so the slope is

$$\mathrm{slope} = \frac{I \rho}{A}$$

Though it is not asked for directly in the question, we can find the resistivity, *ρ*, by measuring *I* with an ammeter/galvanometer and A with vernier calipers or a micrometer screw gauge.

###28.5.2 Metre Bridges
A metre bridge resembles a potentiometer, except that it uses a galvanometer to measure the difference in current between two points on the circuit, hence the name “bridge.” The same materials can be used as with the potentiometer, though it is best to use small coils of resistance wire for the small resistors (between 3Ω and 20Ω is a good resistance). A galvanometer can be made easily if one is not available.

![meter-bridge-2.png](images/meter-bridge-2.png)

Resistors R$$_1$$ and R$$_2$$ have different resistances, but they should be somehow similar so that one resistor does not take all of the current (this will make it difficult to measure the length to the galvanometer). About 5Ω and 10Ω, for example, would work well.
However, for the sake of the practical, one resistor should not be known; the objective of the practical is to find the unknown resistance. The long wire along the bottom edge is a metre of nichrome wire or other resistance wire. One terminal of the galvanometer is connected between the two resistors, and the other terminal is connected to a flying wire (or jockey) that is free to move along the length of the nichrome wire.
The practical instructs you to move the galvanometer’s flying wire back and forth along the nichrome wire until it reads zero. At this point, we know that no current is passing through the galvanometer, so the potential difference across it is zero. This means that the current flowing through R$$_1$$ is the same as that current flowing through R$$_2$$, and the current flowing through the nichrome wire is constant. From this we can conclude that

$$\frac{R_1}{L_1} = \frac{R_2}{L_2}$$

or that the ratio of the two resistors is equal to the ratio of distances from the flying wire to either end of the nichrome wire. The resistance of one resistor (say, R$$1$$) is known and the lengths L$$_1$$ and L$$_2$$ can be measured from the flying wire to either side of the nichrome wire. Using the ratio above, we can easily calculate the unknown resistance R$$_2$$. An example is given below from the 2006 NECTA exam.

**Sample Practical Question**

You are required to determine the unknown resistance labeled *X* using a metre bridge circuit. Connect your circuit as shown below, where *R* is a resistance box, *G* is a galvanometer, *J* is a jockey and others are common circuit components.

![2006-3-alt.png](images/2006-3-alt.png)

**Procedure:**

With *R = 1 Ω,* obtain a balance point on a metre bridge wire *AB* using a jockey *J*. Note the length *l* in centimetres. Repeat the experiment with *R* equal to *2 Ω, 4 Ω, 7 Ω* and *10 Ω.*

Tabulate your results for *R*, *l* and $$^1/l$$.

(a)
1. Plot a graph of R (vertical axis) against $$^1/l$$ (horizontal axis).
2. Determine the slope *S* of your graph.
3. Using your graph, find the value of R for which $$^1/l$$ = 0.02.

(b) Read and record the intercept R$$_0$$ on the vertical axis.
(c) Given that,

$$R = \cfrac{100 \text{X}}{l} - \text{X}$$

Use the equation and your graph to determine the value of *X*.

(d) Comment on your results in (a)(iii), (b) and (c) above.

**Discussion**

The procedure for this question is similar to most other wheatstone bridge problems: vary a known resistor and see how it affects the relative lengths in resistance wire required to balance the potential difference and give no current through the galvanometer. It may not be obvious at first, however, where the equation $$R = \frac{100 \text{X}}{l}$$ comes from.

Starting with the balancing ratio for a wheatstone bridge,

$$\frac{R_1}{L_1} = \frac{R_2}{L_2}$$

we can solve for the unknown resistor

$$R_2 = R_1\left(\frac{L_2}{L_1}\right)$$

Recall that L$$_1$$ and L$$_2$$ are the corresponding lengths from either end of the metre rule to the jockey (in cm), and so taking them together, we get

$$L_1 + L_2 = 100$$

Dividing both sides by L$$_1$$ gives

$$1 + \frac{L_2}{L_1} = \frac{100}{L_1}$$
￼￼
Then solving for $$\frac{L_2}{L_1}$$

$$\frac{L_2}{L_1} = \frac{100}{L_1} - 1$$

Now substitute this into our previous equation for R$$_2$$:

$$R_2 = R_1\left(\frac{100}{L_1} - 1\right)$$

Replacing R$$_2$$ with *R*, L$$_1$$ with *l* and R$$_1$$ with X for this problem and distributing gives

$$R = \frac{100 \text{X}}{l} - \text{X}$$

From this equation, we have dependent variable *R*, independent variable $$\frac{1}{l}$$, slope 100X and *y*-intercept −X. So we can obtain the value of the unknown resistor X either by using the y-intercept (note the resistance is a positive value) or by taking the slope divided by 100.

###28.5.3 Ohm’s Law

The practical may give any kind of experiment to use or verify Ohm’s Law in a simple circuit. Finding the e.m.f. and internal resistance of cells appears frequently. Students should be very familiar with the law, as well as the factors that determine resistance in a wire and the effect of internal resistance of a cell on a circuit. Given below is an example problem taken from the 2011 NECTA exam.

**Sample Practical Question**
You are provided with an ammeter, *A*, resistance box, *R*, dry cell, *D*, a key, *K* and connecting wires. Proceed as follows:

1. Connect the circuit in series.
2. Put *R = 1 Ω* and quickly read the value of current *I* on the ammeter.
3. Repeat procedure (b) above for *R = 2 Ω, 3 Ω, 4 Ω* and *5 Ω*. Record your results in a tabular form.
4. Draw the circuit diagram for this experiment.
5. Plot the graph of *R* against $$\cfrac{1}{I}$$.
6. Determine the slope of the graph.
7. If the graph obeys the equation $$R = I − r$$, then
(i) suggest how *E* and *r* may be evaluated from your graph.
(ii) compute *E*.
(iii) compute *r*.
8. State one source of error and suggest one way of minimizing it.
9. Suggest the aim of this experiment.

**Discussion**

To see where the equation $$R=\cfrac{E}{I}-r$$ comes from, first start with Ohm’s Law, $$V = IR$$. Accounting for the internal resistance of the cell, *r*, this becomes

$$V = I(R + r)$$

To solve for resistance, we divide both sides by I, which gives

$$(R + r) = \frac{V}{I}$$

From this we can see that, using E as e.m.f. for this problem,

$$R=\cfrac{E}{I}-r$$

In this form, the equation resembles the classic *y = mx + c*, where *R* is the dependent variable, $$\frac{1}{I}$$ the independent variable, *E* is the slope, and *r* is the *y*-intercept (note the internal resistance is a positive value).
chapter29
#Chapter 29: Hosting Science Events

There are many great ways to promote math and science education through engaging activities for students and teachers alike. These can be done regularly through extracurricular clubs, but can also be organized together as part of a larger Science Day event or multi-day Math and Science Conference. What follow are some general tips and suggestions for hosting some of these various activities.

##29.1 Box of Fun

The Box of Fun can be used as a teacher training exercise or as a student challenge.

- Gather an assortment of everyday materials (see Sources of Laboratory Equipment, p. 208) and arrange them randomly across a table or in a large box.
- Ask participants to use the materials given to demonstrate some topic or principle in a subject of their choice (Biology, Chemistry, Physics or Math).
- You may choose to put participants into groups and designate a specific subject for each group.
- After at least 30 minutes, have groups come up to present their idea.
- Additionally, you may ask groups to fill out an activity template (see Shika Express companion manuals) to document their ideas.

The Box of Fun is intended to foster participants’ creativity and encourage them to see science in the world around them. Rather than thinking first of a topic and then deciding what materials are needed to show it, this activity encourages teachers to first look around and see what is available to them, and then to think about how those things might be used to demonstrate some concept in science.

CAUTION: It may not be wise to assign specific topics to participants, as this can limit their creativity and may lead them to “destroy science” by pretending a local substitute gives the same result as a traditional lab material when it really doesn’t (e.g. pretending food colour is iodine because they have similar appearances).

The purpose of using locally available materials is that they help to connect students to their everyday environment while still achieving the same results as expensive lab equipment. However, if a local material does not give the same result, it should NOT be substituted merely for the sake of using local materials.

##29.2 Shika Express Gallery Walk
This activity can be used to share ideas of science demonstrations among students and/or teachers.

- Choose 4-6 activities or demonstrations for each subject (see Shika Express companion manuals for Biology, Chemistry, Physics and Math).
- Prepare the demonstrations and arrange them across a set of tables, 1-2 tables per subject.
- Spread the tables out evenly around a large empty room (e.g. dining hall).
- Divide participants into equal groups based on the number of subjects being presented.
- Have groups rotate among the different subject tables so that they are able to observe all demon- strations for each subject (approx. 15-20 mins per subject).
- Following the rotations, give 20-30 mins to allow participants to return to a demonstration of their choice for further investigation or to construct it themselves.

Suggestions:

- You may wish to have 1-2 student or teacher leaders for each subject to help explain the demon- strations during the rotations.
- Make copies of activity write-ups from the Shika Express manuals for each demonstration that participants can read as they walk around.
- Allow participants to perform the demonstrations themselves as much as possible, and then ask them to explain what they see.

##29.3 Science Fair
Science Fair projects provide a great opportunity for students to apply their knowledge and investigate their interests in science.

- Have interested students form groups of 2-3.
- Groups select a project idea based on a shared interest or question/problem to address. Encourage students to think about what problems or issues are faced in their own communities.
- Review the steps of The Scientific Procedure (see example activities on p. 152). Have students identify the problem and form a hypothesis for their project.
- Allow several weeks for groups to work on their projects (provide additional books or computer resources if available).
- When completed, allow students to set up and explain the various projects around the school for all students to see.
- Encourage students to apply to participate in the national Young Scientists Tanzania (YST) com- petition (www.youngscientists.co.tz) in Dar es Salaam.

##29.4 Science Competitions
Perform individual competitions or many strung together over the course of a day or weekend. For more, see the section on Science Competitions (p. 145).

##29.5 Science Day
Engage the entire school (or multiple schools) by combining several activities into a Science Day event.

- Invite a guest speaker to speak on career opportunities in math and science (e.g. accountant, engineer, doctor, nurse, carpenter, mechanic, store owner, etc.)
- Explain applications of math and science in all walks of life (e.g. farming, buying/selling, health/disease, transport, weather, drinking water, football, etc.)
- Incorporate Science Competitions - elect 1 or 2 teams from each Form to compete, with the rest of the school as an audience.
- Incorporate Box of Fun and Gallery Walk activities. It may be helpful to do the Gallery Walk first to provide examples to participants.
- Encourage girls’ empowerment wherever possible.
- Give out a survey to gauge students’ perception of science.

A Science Day event may not guarantee immediate improvements in test scores, but it shows students that the school and its teachers are not willing to give up on math and science, and neither should they! Continued promotion of math and science will help to change students’ perception of the subjects that they may initially write off as being too difficult. Excitement and interest is the first step in changing that perception.

##29.6 Math and Science Conferences
Gather students from several nearby schools to hold a special week-long Math and Science Conference in a nearby town or at a host school.In addition to those ideas presented for a Science Day event,

- Incorporate HIV/AIDS and malaria into science demonstrations/activities (see Shika Express companion manuals for ideas).
- Have students prepare Science Fair projects over the course of the conference and present on the final day.
- Give award certificates for participation and prizes for individual/team competitions.

Math and Science Conferences encourage leadership among their participants. Students attending such events are likely to be good ambassadors of science, sharing what they have done and learned with fellow students back at school. Those students may then try to improve their own performance in those subjects so that they can attend a similar event later on.

##29.7 Teacher Trainings
Conferences can also be directed towards improving teacher performance by conducting the suggested activities at a nearby Teacher’s College.
chapter30
#Chapter 30: Science Competitions

Students, just like nearly all other people, enjoy competing against one another. Likening math and science-related activities to the competition of a football league can be a wonderful motivator for students. Given below are some suggestions for utilizing competitions while teaching about science.

- Combine students of various abilities together on a team. This will allow the bright students to develop leadership skills and help to bring up the slow learners.
- Limit teams to 3-5 students. Balance the number of boys and girls on a team, or choose to have all-boys teams compete against all-girls teams.
- Allow students to pick their own team name, and possibly draw a team flag if time allows.
- Create a standings board for the competition. For example:

| | Egg Drop | Jenga Jengo | Raft Rally | Drop Zone | Bridge Challenge | TOTAL |
| ----------------- | -------- | ----------- | ---------- | --------- | ---------------- | ----- |
| Big Stars | | | | | | |
| Chelsea | | | | | | |
| Simba | | | | | | |
| Arsenal | | | | | | |
| Kings | | | | | | |
| Manchester United | | | | | | |

- Have teams present and explain their designs to the audience where applicable. Allow other students to ask questions/provide criticisms.
- Follow up each activity with a short lesson about a concept illustrated by the competition (e.g. Archimedes’ Principle for Raft Rally; see competition write-ups for more).
- Ask students how they would revise their designs or make improvements if they could do the activity again.
- Explain how the activities can be applied to solve real-life problems (e.g. Jenga Jengo for Civil Engineers).

##30.1 Egg Drop

**Time:** 1 hour

###How It Works:
Students must build a device to transport an egg through a given drop distance without cracking.

###What You Need (per team)

- Plain paper (4 sheets)
- Plastic Bag (2)
- String (1 meter)
- Balloons or, if unavailable, condoms (4)
- Straws (10)
- Toothpicks (10)
- Tongue depressors (4)
- Rubber bands (4)
- Index cards (4)
- Paper clips (10)
- Toilet paper (1 roll)
- 500 mL bottle (1)
- Newspaper (1 sheet)
- Egg (1)

###Rules

- 45 minute time limit for construction.
- Devices dropped from a height of 3-5 metres.
- Teams can use only the materials given, but do not need to use everything.
- Egg is placed at time of testing. It must be possible to place and remove egg freely without altering the device.
- Once egg is placed, no further adjustments may be made. This means the egg cannot have any kind of “seat belt” or strap fastened after placing the egg.

###Points

Egg Survives - 50 pts. Egg Cracks - 0 pts

- Scissors for community use

###Notes

- Do not give eggs to teams until time of testing.
- If possible, increase drop height for surviving eggs and give bonus 25 pts for each additional successful drop.

###Science Applications

**Air Resistance (Physics Form I):** Air resistance provides a frictional force which opposes the ob- ject’s motion as gravity attracts it towards the centre of the earth. This upward force reduces the speed of the object as it falls, allowing it to land more softly and protect the egg. Thus, we want to maximize the air resistance on the object (e.g. by using a large parachute).

**Pressure (Physics Form I):** The force of impact on the device when it hits the ground can be reduced by increasing the surface area which contacts the ground. Constructing a wide base (e.g. using balloons) reduces the impact on the egg and thus helps to protect it.

###Taking It Further
Did students utilize the parachute concept? If not, show a brief example. How does this help to protect the egg? Is it better to have a large parachute or a small one?

## 30.2 Jenga Jengo

**Time:** 30-45 minutes

###How It Works
Students must build the tallest structure possible, using only paper and tape, as quickly as possible and while ensuring good stability.

###What You Need (per team)

- Plain paper (25 sheets)

###Rules
- 20-minute time limit for construction.
- Cannot use tape roll as weight inside structure.
- Stability tested by waving book at structure (Wind Test).

###Points
1st to finish – 50 pts
Tallest structure – 50 pts
Passes Wind Test – 50 pts

- Tape Measure
- Stopwatch
- Book / waving device

###Notes
- All structures that pass the Wind Test are awarded 50 pts.
- Alternate Wind Test: place structures outside on a windy day. Those standing after 1 minute pass.

###Science Applications

**Centre of Gravity (Physics Form I):** Civil engineers construct buildings with a low centre of gravity, making them less likely to fall over due to wind forces. To maintain stable equilibrium, a building should have a wide base with a large mass, while the top of the building should have a small area and less mass.

###Taking It Further
- Show students pictures of buildings and structures from around the world after the competition. Did the students’ structures resemble any of them?
- Try variations, giving students index cards, straws or matches instead of plain paper.

##30.3 Raft Rally

**Time:** 30-45 minutes

###How It Works
Students must build a raft using only aluminum foil that can support the heaviest load before sinking.

###What You Need (per team)
- Aluminum foil – 20 cm × 20 cm sheet
- Straws - 4 (optional)

###Rules
- 10-minute time limit for construction.
- Replacement sheet may be given in case of rips/tears, at a 20 pt deduction.

###Points
1st Place – 100 pts
2nd Place – 75 pts
3rd Place – 50 pts
4th Place – 25 pts
Others – 0 pts

- Large container or bucket (clear if possible) filled with water
- Nails (× 200) / Bottle caps (× 200) / Other small weights for testing

###Notes
- As raft approaches the point of sinking, add weights more slowly.
- Raft is finished when water begins to enter, and total number of weights is recorded.

###Science Applications

**Archimedes’ Principle (Physics Form I):** *Archimedes’ Principle* states that

Upthrust = Weight of displaced fluid

Here, we want to maximize the force of upthrust to avoid sinking. So that means maximize the Weight of the displaced water: Weight = mass x acceleration due to gravity, or

$$W=mg$$

Gravity is a constant , but mass depends on two things: density $$\rho$$ and volume $$V$$.

We know that $$\rho = {m}\over{V}$$, so that means $$m = \rho V$$.

The density of the water is constant, so the only thing we can change is the Volume of water displaced. Thus to get the most upthrust and prevent sinking, we need to displace a large volume of water, i.e. build a raft with a large base.

###Taking It Further
- Ask students how they would revise their designs if they could do it again.
- Try variations, giving students straws, toothpicks, tongue depressors or index cards.

##30.4 Drop Zone

**Time:** 30-45 minutes

###How It Works
Students must build a parachute using limited materials to carry a paper clip passenger as close as possible to a target, while maximizing hang time.

###What You Need (per team)
- Paper clip (1)
- Plastic bag (2)
- String (1 metre)
- Plain paper (2 sheets)
- Scorecard (see example below)

###Rules
- 10-15 minute time limit for construction.
- Parachutes dropped from a height of 3-5 metres.
- Average hang time and distance from target taken over 3 trials for each team.

###Points
Hang Time (Longest): 1st – 50 pts, 2nd – 35 pts, 3rd – 20 pts, 4th - 5 pts, Others - 0 pts
Distance (Shortest): 1st – 50 pts, 2nd – 35 pts, 3rd – 20 pts, 4th - 5 pts, Others - 0 pts

- Tape measure
- Flipchart target
- Stopwatch

###Notes
- Scorecard:

| **Team** | **Trial 1** | **Trial 2** | **Trial 3** | **Average** | **Points** |
| -------- | ----------- | ----------- | ----------- | ----------- | ---------- |
| **Hang Time (s)** | | | | | |
| **Distance from Target (cm)** | | | | | |

- Measure distance from paper clip to centre of target.

###Science Applications

Air Resistance (Physics Form I): Air resistance provides an upward force on the parachute, which acts against the force of gravity and causes the object to fall more slowly. The larger the surface area of the parachute, the more slowly it will fall.

###Taking It Further
- Students may not be familiar with parachutes. Prepare a simple example to explain the concept and function.
- Ask students questions: Why does the parachute slow the object down? To maximize hang time, do we want a very large or very small parachute? Would a parachute work on the moon?
- Drop parachute side-by-side with a paper clip having no parachute. Which one made it safely?

##10.5 Bridge Challenge

**Time:** 1 hour 30 minutes

###How It Works
Students must build a bridge that can support the most weight, while using a limited budget of Science Shillings to purchase construction materials.

###What You Need (per team)
- Straws (20)
- Bamboo skewers (20)
- Bamboo stick (fimbo) (1)
- Tongue depressors (10)
- Toothpicks (2 small cans)
- String (3 meters)
- Office glue (1 tube/jar)
- Rubber bands (10)
- Pencils (2)
- Index cards (10)
- Duct tape (1 roll, for everyone)
- Ruler (1)
- Scissors (1)
- Science Shillings (20)

###Rules
- Approximately 45 minute time limit for construction.
- Teams begin with only a ruler, scissors and 20 Science Shillings. These items may NOT be used in construction of bridge.
- All building materials must be purchased from a science shop. Suggested prices are as follows:

Straws (bundle of 10) 1/=
Skewers (bundle of 10) 2/=
Fimbo 3/=
Tongue depressors (× 5) 2/=
Toothpicks (2 small cans) 1/=
String (1 metre) 1/=
Office glue (1 tube) 3/=
Rubber bands (× 5) 1/=
Pencil 1/=
Index cards (× 5) 1/=
Duct tape (30 cm) 1/=

- Bridges will be loaded by placing rocks or other weights into a small bucket that must rest on top of the bridge.
- Bridge must span a 30 cm gap between two chairs / tables.
- 1 student from each team must be designated as team accountant. Only this student may purchase items from the shop.

###Points
(Based on number of rocks / weights placed before bridge fails)
1st – 100pts, 2nd – 75pts, 3rd – 50pts, 4th – 25pts, Others – 0pts
**BONUS:** 5 pts per Science Shilling remaining after construction

- Small bucket
- Several large rocks (20) / other large weights
- 2 chairs / tables 30 cm apart

###Notes
- Science shop table
- Extra bamboo sticks (fimbos) – 2-4 • Index Cards for price signs – 12
- Student accountants may only purchase 1 of each item at a time. They must allow other students to make purchases before buying another of that item.
- At some point 15-20 minutes into construction time, shopkeeper may announce a newly received shipment of bamboo sticks (fimbos). However, due to demand, the price has increased to 5 /=.
- Bridge has failed when it either collapses / breaks or when the bucket can no longer be balanced on top of it. Record largest number of weights successfully added.

###Taking It Further
- Ask students to present their bridges and describe how they decided to manage their money. What materials did they purchase and why?
- What would they do differently if they could start again?
- Which team finished with the most total points after the BONUS? Did the most money spent result in the strongest bridge? Who was the most efficient with their money?
- Show students pictures of bridges from around the world after the competition. Did the students’ bridges resemble any of them?
chapter31
#Chapter 31: The Scientific Procedure

The following activities can be used as a method of introducing students to the scientific method. Rather than just performing the activities, first identify the question or problem with the students, then have them form a hypothesis for each step of the experiment. Students should record observations and data accordingly and use them to draw a conclusion about the activity.

Prepare an activity sheet for each student or have them copy it into their notebooks before performing the activities. Set up stations for the various activities and have students rotate among them in small groups.
After performing several of the experiments, ask students to come up with their own. Ask them to think about problems they face in their daily lives. Encourage interested students to turn their ideas into a science fair project to display for the school or community.

##31.1 Biology

###Hand Washing

**Materials:** Soap, water, bottle, basin/bucket, chalk, charcoal, food colour, stopwatch
**Setup:** Prepare a large amount of soapy water. Grind the chalk and charcoal into separate powders.

**Problem:** How long should we wash our hands?

| **Material** | **Hypothesis (in seconds)** | **Experimental Result** |
| ------------ | --------------------------- | ----------------------- |
| Chalk powder | | |
| Charcoal powder | | |
| Food color | | |

**Hypothesis:** Predict how much time it will take to completely clean your hands and record in the table.
**Procedure:** Start a stopwatch and have a student or teacher slowly pour soapy water over a basin while the student washes his or her hands. Stop the clock when the student’s hands are completely clean.
**Observations:** Record the time taken to completely wash your hands in the table.
**Questions:**

1. Why is it important to wash our hands?
2. When do we need to wash our hands?

**Theory:** Washing our hands with soap and water helps to kill harmful bacteria that can cause us to become sick if allowed into our bodies. It is very important to wash our hands before eating and after using the bathroom.

###Lung Capacity

![lung-capacitypng.txt](images/lung-capacitypng.txt)

**Materials**1.5 L bottle, basin, water, plastic tubes/straws, soap, marker, ruler
**Setup:** Make a scale on the bottle using a marker and ruler (e.g. 100 mL increments). Prepare a soap solution for washing the tubes/straws Problem: How much air can your lungs hold?

| **Breath** | **Hypothesis (Volume of air in mL)** | **Experimental Result** |
| ---------- | ------------------------------------ | ----------------------- |
| Normal breath | | |
| Full breath | | |
| After holding breath for 10 seconds | | |

**Hypothesis:** Record the volume of air that you think the lungs can hold for each case in the table.
**Procedure:** Fill a basin with water. Fill a 1.5 L bottle with water and invert it in the basin so that the mouth of the bottle is underneath the water. Place one end of the tube/straw inside the bottle under water. For each breath, blow into the tube to displace the water.
**Observations:** Note the reading on the scale before and after blowing into the tube and record the *difference* to give the amount of water displaced.
**Questions:**

1. Which breath produces the largest amount of air? Which give the smallest amount?
2. How long can you hold your breath?

**Hypothesis:** I can hold my breath for ￼________ seconds.

**Experimental Result:** I can hold my breath for ￼____________￼ seconds.

**Theory:** When we breath in air, our bodies use the oxygen and produce carbon dioxide in a process called *respiration.* Oxygen is transported in our blood throughout our bodies. When we hold our breath, oxygen is not circulated throughout our bodies and we begin to feel lightheaded.

##30.2 Chemistry

###Acids and Bases

![acids-bases-sci-meth.jpg](images/acids-bases-sci-meth.jpg)

**Materials:** Bottles, bottle caps, water, vinegar, lemons, baking soda, soda, soap, antacid tablets, rosella leaves, straws/syringes
**Setup:** Prepare solutions for each of the items above in separate bottles. Prepare indicator by placing rosella leaves in hot water.

**Problem:** What differences can we observe among acids and bases?

| **Solutions** | **Hypothesis (Which is different?)** | **Experimental Result** |
| ------------- | ------------------------------------ | ----------------------- |
| Vinegar, lemon, baking soda | | |
| Vinegar, baking soda, soap | | |
| Baking soda, antacid, soda | | |
| Soda, soap, vinegar | | |

**Hypothesis:** For each set of solutions, which one will reveal a colour different from the others? Record your predictions in the table.
**Procedure:** Place small amounts of 3 different solutions in separate bottle caps according to the table. Add a few drops of rosella indicator to each.
**Observations:** Record observations of colour change under Experimental Result in the table.
**Questions:**

1. Which solutions have similar properties?
2. Which solutions are acids? What colour do they show?
3. Which solutions are bases? What colour do they show?

**Theory:** Coloured leaves such as rosella act as indicators for identifying acids and bases. Adding rosella indicator reveals a red colour for acids and a blue colour for bases. Students do not need to understand the differences between acids and bases in order to observe their different behaviours. Locally available examples of acids include sour milk, citrus fruits and soda. Local bases include ammonia, toothpaste and detergent.

###Mixing Acids and Bases

![mixing-acid-base.jpg](images/mixing-acid-base.jpg)

**Problem:** WHat happens when acids and bases are mixed together?

| **Solutions to Mix** | **Hypothesis (what color?)** | **Experimental Result** |
| -------------------- | ---------------------------- | ----------------------- |
| Mix vinegar and lemon | | |
| Mix baking soda and soap | | |
| Mix vinegar and baking soda | | |

**Hypothesis:** Predict any colour changes or observations when pairs of solutions are mixed together. Record in the table.
**Procedure:** Mix small amounts of solutions together according to the table. Observations: Record observations (colour changes, etc.) in the table.
**Questions:**

1. What happens when an acid is mixed with an acid?
2. What happens when a base is mixed with a base?
3. What happens when an acid is mixed with a base?

**Theory:** Mixing acids with acids and bases with bases may cause the colour of the solution to turn darker or lighter depending on the solutions used. Mixing an acid with a base should reveal a colourless solution and produce carbon dioxide gas. You may need to vary the amounts of acid and base to get a colourless solution depending on their concentrations.

##30.3 Physics

###Complete the Circuit

![screen-shot-2014-09-10-at-21946-pm.png](images/screen-shot-2014-09-10-at-21946-pm.png)

**Materials:** Dry cell, speaker wire, bulb/ammeter, cardboard, various objects, e.g. rubber band, nail, paper, aluminum foil, toothpick, pen, scissors, bottle cap, coin, balloon, chalk
**Setup:** Connect a dry cell and bulb in series using speaker wire and attach to a sheet of cardboard. Leave two wires free and pin to the cardboard to act as a switch.

**Problem:** Which objects will light a bulb?

| **Object** | **Hypothesis (Light or No Light)** | **Experimental Result** |
| --------------- | ---------------------------------- | ----------------------- |
| Copper wire | | |
| Pen | | |
| Aluminum Foil | | |
| Paper | | |
| Nail | | |
| Toothpick | | |
| Bottle cap | | |
| Chalk | | |
| Scissors (blade) | | |
| Scissors (handle) | | |

**Hypothesis:** Predict which materials will cause the bulb to light when placed across the switch. Record predictions in the table.
**Procedure:** Test each object by placing it across the free wires to close the circuit. Observations: Record the result for each item in the table.
**Questions:**

1. Which materials caused the bulb to light?
2. These objects are made from what kind or materials?
3. What other objects in the room can you find to test? Will they light the bulb?

**Theory:** Conductors are materials which easily allow electrons to flow through them. Insulators are materials which do not easily allow the the flow of electrons. Examples of good conductors are most metals, water and the human body. Examples of good insulators are rubber, wood and plastic.

###Density Tower

![density-tower-sci-meth.png](images/density-tower-sci-meth.png)

**Materials:** Syringes, bottles, water, cooking oil, kerosene, spirit, honey, glycerine, tape, scissors
**Setup:** Prepare a test tube rack by cutting a bottle and filling it with dirt. Remove the plungers from
the syringes and seal them with tape, super glue, or by melting to opening closed.

**Problem:** Which liquids are more dense than others?

| **Liquid** | **Hypothesis (Position, 1 = bottom)** | **Experimental Results** |
| ---------- | ------------------------------------- | ------------------------ |
| Water | | |
| Cooking oil | | |
| Kerosene | | |
| Spirit | | |
| Honey | | |
| Glycerine | | |

**Hypothesis:** Predict the order in which the liquids will settle from the bottom of the syringe. Assign 1 to the bottom liquid, 2 to the one above it, and so on.
**Procedure:** Pour a small amount of each liquid into a syringe, observing after each addition.
**Observations:** After adding all liquids, record the order in which they rest, starting with 1 at the bottom.
**Questions:**

1. Which liquid finished at the bottom?
2. Which liquid finished at the top?
3. Which liquid has the greatest density?
4. Which liquid has the lowest density?
5. What happens if you place a small object (e.g. paper clip, eraser, paper) in the tower?

**Theory:** *Density* is a property of different materials and liquids. It is a ratio of its mass to its volume. Dense liquids sink to the bottom, while less dense liquids rise to the top. A small object placed in the tower will settle in the liquid which is nearest its own density.

###Sinkers and Floaters
**Materials:** Basin of water, various objects, e.g. nail, paper clip, paper, aluminum foil, soda cap, matchbox, pen cap, toothpick, balloons, flour
**Setup:** Have a controlled environment to test sinking and floating objects

**Problem:** Which objects sink or float when placed in water?

| **Object** | Hypothesis (Sink or Float?)** | **Experimental Results** |
| ---------- | ----------------------------- | ------------------------ |
| Nail | | |
| Paper clip | | |
| Pen cap | | |
| Soda cap (dropped) | | |
| Soda cap (placed carefully) | | |
| Toothpick | | |
| Paper | | |
| Aluminum foil | | |
| Matchbox | | |
| Balloon (empty) | | |
| Balloon (filled with flour) | | |
| Balloon (filled with water) | | |
| Balloon (filled with air) | | |

**Hypothesis:** Predict whether each object will sink or float when placed in the basin of water. Record in the table.
**Procedure:** Place each object in the water. First place them very carefully, then drop them in.
**Observations:** Record the results in the table.
**Questions:**

1. What factors affect whether an object sinks or floats?
2. How do large objects such as boats float?

**Theory:** *Flotation* depends on several things. A bottle cap placed carefully on the surface of the water will float, but when pushed under, will sink. A sheet of aluminum foil will float while a sheet of the same size which is folded several times will sink. A balloon filled with flour sinks, one filled with water just floats, and one filled with air floats above the surface.

If an object’s *total density* is greater than that of water, it sinks, but if less than water, it floats. Air has a density less than water, so when air is trapped in objects such as bottle caps or balloons, they float because their total density is less than water. When air is removed (folded aluminum foil) or replaced by water (bottle cap), the total density of the object is just the density of the material. A matchbox pushed under water rises back to the surface because its density is less than that of water.

Boats are able to float despite being built from dense materials because of the large volume of water they displace and the large amount of air inside the boat. A boat with a larger surface area displaces a larger volume of water and thus can carry a larger load before sinking.

Follow up this activity with the *Raft Rally* science competition.

###Mixing Colours

**Materials:** Various food colours, syringes, bottle, scissors, tape, paper
**Setup:** Prepare a test tube rack by cutting a bottle and filling it with dirt. Remove the plungers from the syringes and seal them with tape, super glue, or by melting to opening closed.
**Problem:** What happens when we mix different colours?

| **Colors to Mix** | **Hypothesis (What result color?)** | **Experimental Results** |
| ----------------- | ----------------------------------- | ------------------------ |
| Red and green | | |
| Yellow and blue | | |
| Red and yellow | | |
| All colors | | |

**Hypothesis:** Predict which colour will result when the two colours given are mixed together. Record it in the table.
**Procedure:** Use syringes to remove small amounts of each colour and place on a sheet of paper. Be sure to lay down plenty of paper so that the colours do not bleed through onto the table!
**Observations:** Record the resulting colour mixture in the table.
**Questions:**

1. How can you make orange from other colours?
2. What colour do you get by mixing all of the colours together? 3. What are some uses of coloured dyes?

**Theory:** Red, green and blue are primary colours of light. Other colours are made by different combi- nations of these primary colours. Coloured dyes are used for many applications, including clothes, paper and printing pictures.