EXTENDED EXPERIMENTAL INVESTIGATIONS: Chemistry

One caution in all of this: merely measuring the concentration of the various components of a selection of wines (ethanol, pH, acidity and so on) may not make a good EEI. Manipulation of variables gives students a better chance of demonstrating all aspects of the assessment criteria.


The first little grape buds are seen in August.	A rose bush is planted at the end of each row of vines as an indicator of infection	Fermentation tank at Sirromet

Oxidation of wine to vinegar (1)
Once a bottle of wine is opened and the air gets in it starts to go off (even if it is re-stoppered). It may not be noticeable for a couple of days but it get more and more acidic as the ethanol is oxidized to ethanoic acid (acetic acid) by acetic acid bacteria (called Acetobacter). This may be okay if you want to make wine vinegar but not so good if you want the wine to be drinkable. The question is: what factors affect the ability of the bacteria to oxidize the ethanol in wine. Acetic acid bacteria are aerobic microorganisms and thus will not grow in anaerobic (without air) conditions. You may guess temperature, but what about acidity (from the natural tartaric acid) or preservatives (perhaps SO₂ or benzoic acid), or even the amount of ethanol in the wine (spirits, such as Vodka, with 45% ethanol concentrations don't turn to acetic acid). It would seem that a good EEI could be developed from oxidizing wine (perhaps by a controlled aeration of wine samples with an aquarium pump ($10) and measuring the acidity after a given time (eg aerate each sample for 24 hrs at the same bubble rate). The independent variables (IV) could be one or more of:
    1. temperature; the optimum growth temperature for Acetobacter is 25-30ºC, with no growth observed at 40ºC. Weak growth was observed even as low as 10ºC, but none at 8ºC.
    2. acidity or pH (perhaps add some tartaric acid to have a range of starting pHs (the optimum pH for the growth of acetic acid bacteria is 5.5-6.3, however, these bacteria can survive at the low pH values of between 3.0 and 4.0 found in wine. A pH of 3.3 and lower is inhibitory to most bacteria in wine, but not to acetic acid bacteria. Maybe try pHs of 2, 3, 4, 5, 6.
    3. ethanol concentration. Ethanol is a good carbon source for acetic acid bacteria, but is also inhibiting at concentrations that are too high. One report I saw said that in wine containing 5% ethanol, only 58% of the Acetobacter was active and that this was reduced to only 13% in wine containing 10% ethanol. At 15.5% it seems all Acetobacter activity is inhibited (stopped). You could take some wine and add ethanol to it as the independent variable. The problem is - where do you get ethanol from? Your school may not have a licence to buy 100% (absolute) ethanol, or even the 95% azeotropic mixture (with water) so you may have to distill your own from shop-bought wine. From the density of the distillate (use an SG bottle) you can calculate how much to add to some fresh wine.
   4. additives. Sulfur dioxide should prevent the growth of acetic acid bacteria in wine and is sometimes used commercially for this purpose. SO₂ in wine consists of the free form (molecular SO₂, bisulphate and sulphite ions) and a bonded form. At normal wine pH only about 5% of the free SO₂ occurs in the molecular form (which is the most active anti-microbial form) and the other 95% as bisulfate and sulfite ions. Concentrations of up to 20 mg/L of free SO₂ will kill the bacteria. The simplest thing to do is to use powdered sodium metabisulfite which is available from home-brew shops (and from many health food stores as a anti-bacterial bottle wash for bottling fruit and so on). Try 0 to 25 mg/L free SO₂.

I have attached a great paper "The occurrence, control and esoteric effect of acetic acid bacteria in winemaking" by W.J. Du Toit, and I. S. Pretorius from the Department of Viticulture and Oenology, Institute for Wine Biotechnology, Stellenbosch University, South Africa. It was published in Annals of Microbiology, 52, 155-179 (2002). Click here to download.

The total acidity can be measured by titrating against sodium hydroxide. Look up a table to get the best indicator (weak acid, strong base). The photos below shows Jennifer and Rachel - Yr 11 Chem students from Moreton Bay College - turning wine into vinegar. Be warned though that leaving a sample of wine open to the air for a couple of days may make it taste sour (acetic acid) but this may not be sufficient change in acidity to be picked up by an acid-base titration. More aeration that this may be necessary; eg, keeping it stirred, or having a large surface area. The other problem in measuring acidity is when you have red wine as the pigments disguise the indicator colour change. Chemistry teacher Gareth Whittaker from Mary MacKillop College, Brisbane, advises the use of a pH probe rather than an acid-base indicator; and titrating to an end-point of pH 8.2 (this is called a potentiometric titration to distinguish it from a direct titration which uses an indicator). The method he uses (O24A-FD) comes from the Association of Analytical Communities, AOAC. Click here to download a copy.


Only the finest cask wine is used.	Sample in the refrigerator. Not ethe big surface area.	Jen calibrates the pH meter.

Oxidation of Wine 2 - rate of reaction
In the above EEI, it was suggested that the acidity of the wine be measured after a set time (one day, or one week etc). I said that you could have acidity as the dependent variable (DV, as measured by titration), and one of - temperature, pH, %ethanol or SO₂ as the independent variable (IV). However, another approach is to use "time" as the independent variable. In this case you would set up an experiment where air was blown through a sample of wine and the acidity measured at regular intervals (eg every hour for six hours; or every day for 5 days). You would then plot acidity on the y-axis and time elapsed on the x-axis. The rate of reaction would be the slope of the line at a particular time. Does the rate vary over the whole time period? Does the rate vary as the acidity increases (is there a relationship)? Perhaps you could compare red and white wine. Does the red anthocyanin in the red wine act as an antioxidant as some people believe? What a fabulous EEI, and you'd even have some wine vinegar for your fish and chips afterwards.

Heat of combustion
High school experiments on heats of combustion usually involve burning a candle or alcohol and trapping the heat in a beaker of water. The errors are usually massive and chimneys etc are used to try to trap the heat – with little success. How could the accuracy be improved? You could explore ways and provide a theoretical reason for your trials. Alternatively, you could compare the accuracy of DH_C values of methanol, ethanol, propenol; or even of the three C₄H₉OH isomers. Why might the accuracy be different? What does this tell you about intramolecular bonding? Are there any correlations with BPt?

Heat of combustion of mixtures
The E10 blend of fuel for cars consists of a mixture of petrol with about 10% ethanol added. It is designed to reduce the consumption of non-renewable fuels such as petrol. A good research question for an EEI would be to ask if the resultant heat of combustion of a mixture of fuels is related simply to the proportion of the fuels in the mixture and their DH_c values (for example if Fuel A has a DH_c of 1000 kJ/mol and Fuel B has a DH_c of 2000 kJ/mol does a 50:50 mixture of the two have a DH_c of 1500 kJ/mol. Perhaps there are intermolecular interactions (eg H-bonding effects) between the components. Is there an effect with some mixtures (such as alcohols - with their H-bonding possibilities) that are not apparent with non-polar alkanes? A good suggestion from one teacher is to investigate the DH_c of mixtures of ethanol and butan-1-ol in varying ratios (0:100, 20:80, ...80:20, 100:0). One caution: don't let the errors (which are quite extensive if you use a spirit burner as shown in the photo above) make you think there is a trend when there is not. You will need to control all other variables very carefully to keep all errors as constant as possible.

Heat of combustion of isomers and polyols
Following on from the ideas above, it would be interesting to investigate any trends in the heat of combustion of fuels depending on factors such as their shape. For example, if you took three isomers of C₄H₉OH (butan-1-ol, butan-2-ol and 2-methyl-propan-2-ol) - all commonly available at school - would their DH_c values vary and why? Do they have different boiling points - and what does this tell you about intermolecular forces. Is this related to DH_c and why? Another experiment might be to compare an alkane with it's partially oxidized cousin (eg hexane and hexanol). Does the presence of oxygen make it more combustible? What about comparing a simple alcohol with its diol (two -OH groups), for example butan-1-ol with butan-2,3-diol. The possibilities are endless. If your school doesn't have the chemical you need you may get some help from a university.


H-bonding is greater in n-propanol but does this affect the heat of combusion?	BP 97ºC	BP 82ºC

Corrosion
Corrosion happens all around us - our cars rust, bridges and other steel structures fail, and we spend billions of dollars each year in replacement and maintenance costs as a result. There are a number of methods used to minimize or prevent corrosion, which include alloying, metallic coating, organic coating, use of inhibitors, and anodic or cathodic protection. Corrosion is one of the more popular topics in Queensland schools for an EEI as we have a warm, humid climate and the bulk of the population lives along the coast. Iron is the most abundant metal on earth and has been a boon to the building industry since the Iron Age. However, as it is susceptible to corrosion or rusting, structures made of iron, such as bridges and ships, need to be regularly monitored for rusting. If not, the damage caused by rusting can be very expensive to fix and, perhaps, hazardous. This is a particular problem in the shipping industry where the moist, salty conditions are ideal for accelerating the rusting process. For shipping and many other uses, iron is converted to one of its alloys, carbon steel, to make it stronger and less susceptible to corrosion. Salinity is only one of many factors that will contribute to the nature and extent of iron and steel corrosion observed at shipwrecks. Others factors, such as the concentration of dissolved oxygen, pH, temperature of the water, among others, may have a significant bearing on the corrosion of a particular wreck. That is why shipwrecks at different ocean depths and latitudes may vary in the nature and the extent of corrosion.

Good EEIs often are in the context of shipwrecks. If so, it is not enough to merely put steel nails in different solutions and look at the loss of iron. You should be looking at the environments that ships can be found in and considering how you can simulate corrosion on a speeded up scale. Also important is what you will use for the metal: steel may be okay – but what alloy is it? Is pure iron any use – with no carbon to act as active sites for corrosion? Teachers who have been to Australian Corrosion Association conferences say that their website has useful information. As far as the best way to measure the amount of rusting, you may like to contemplate the advice given by Chemistry teacher Daniel Bischa from Pioneer State High School, Mackay, Queensland. Download his comments here.

pH and photosynthesis

Oxygen is evolved during photosynthesis but the conditions for maximum reaction rate are intriguing. It can be affected by many things, including: sunlight - its intensity and wavelength, temperature, CO₂ and O₂ availability, water (which closes stomata and restricts CO₂), and any factor that influences the production of chlorophyll, enzymes, or the energy carriers ATP and NADPH, such as pH and Mg²⁺ availability. You could test the effect of pH and temperature. It sure won't be linear but how well your prediction (hypothesis) and results agree will be interesting. There are a lot of variables to control and complex biochemical reactions to examine.

Tinny taste of fruit in tin cans
If you look inside an opened can of fruit you will notice that the can appears to have a bare metal surface. The surface is tin which has been electroplated over thin steel sheet, hence "tin can". Sometimes there is an almost invisible clear lacquer film, sometimes Bisphenol-A (BPA) - a possible carcinogen. Tin is a fairly reactive metal and if you leave an open can of fruit for a couple of days the fruit tastes "tinny" as the tin is oxidized by the air in the presence of food acids. The tin acts as the anode (Sn → Sn²⁺ + 2e^-) and the underlying steel acts as a cathode (2e^-+ 2H⁺ → H₂). The steel does not corrode as it is protected by the tin and because the area of iron exposed through tiny pits in the tin is small, the reaction is slow and said to be under cathodic control. If nitrates are present in the food, they will cause rapid detinning by two reactions at the cathode: NO₃^- + 2e^- +2H⁺ → NO₂^- + 4 H₂O (slow); and NO₂^- + 6e^- +8H⁺ → NH₄⁺+ 2H₂O (fast).

A tin/iron cell

Research by P. W. Board at CSIRO published in Food Technology in Australia in 1973 showed that the rate of detinning was dependent on pH and concentration of nitrate ions. The major source of high nitrate fruit at the Golden Circle Cannery in Brisbane was papaw (papaya) and if this was used in canned fruit salad and the nitrate level was high, a can with a plastic lacquer on the inside had to be used. Low nitrate pawpaw needed no such lacquer and so was preferred as it made costs cheaper. Pawpaw farmers received more money for their fruit if it was "low nitrate". Here's a good EEI: make your own tin/iron cell. You could cut up a "tin" can into strips and stand the strips up in a beaker of dilute food acid but maybe a better way would be to construct a cell like the one shown to the left. You could vary the amount of nitrate and pH to investigate rates of detinning.

How to test for tin irons? The simplest way is to do a "iodometric" (redox) titration using a standardised solution of iodate and iodide ions as the titrant. When iodate ions (IO₃^-) are added to an acidic solution containing iodide ions (I^-), an oxidation-reduction reaction occurs IO₃^- + 6H⁺ + 5e^- → ½ I₂ + 3H₂O while the iodide ions are oxidised to form iodine 2 I^- → I₂ + 2e^-. Combining these half-equations demonstrates the reaction between iodate and iodide 2 IO₃^- + 10 I^- + 12 H⁺ → 6 I₂ + 6 H₂O. It is the iodine formed by this reaction that oxidises the Sn²⁺ to Sn⁴⁺ acid as the iodine is reduced to iodide ions. A starch indicator is no use in this technique as the low pH destroys it's action; you need to add an organic solvent to see the iodine clearly. I have attached a method to download that will work although it is a bit more difficult than a regular titration. Colorimetric methods are difficult as the chemicals are hard to get and the methods complex: download one here. The best bet if you have access to a professional lab is to have them do atomic absorption analysis for you. Your best bet may be to not use a tin can but make up a cell with tin and iron electrodes in an electrolyte of acid and nitrate. Good luck! A copy of the relevant chapter from a MSc dissertation can be downloaded by clicking here.


Cross section of a tin-coated steel can cut through with a pair of scissors. Don't use "the good scissors". The thicknesses of each layer is: oil 1 nm, oxide 1 nm, tin 1000 nm, alloy 100 nm, steel 0.2 mm.	Tin crystals are visible on the inside of this fruit can. Note that you can't tell if there is lacquer coating over the tin on the inside.


The lid and bottom of the can is coated inside with a white paint. The walls are not. Here some of the white paint has been sandpapered off the lid and drops of conc. HCl added. Bubbles show up only where the paint has been removed.	Use a drop of conc. HCl to test if there is a coating. You can see bubbles of H₂ where the lacquer has been removed but none on the lacquered part.	A drop of conc. HCl on the outside of the can shows that a lacquer is present (no bubbles).

Stability of Vitamin C in solution
Vitamin C is sensitive to heat, light and oxygen and the degree of sensitivity depends on the pH of the solution (more stable at lower pHs). In food it can be partly or completely destroyed by long storage or overcooking. By refrigeration the loss of Vitamin C in food can be substantially diminished. An interesting EEI would be to see how some of these factors really affect a Vitamin C solution. It may be a good idea to simulate fruit juice by making up an appropriate solution with added citric acid, some citrates, glucose/fructose and so on. Should you measure the concentration of the ascorbic acid with time (and graph) or just measure it each day or after a week or two weeks? What will you control? What will your independent variable be: sugar concentration, [H⁺], light, oxygen, temperature? If you intend to measure the concentration as a function of time elapsed you should read my caution below. The possibilities are endless but you'd need to back up your hypothesis with some justification from the literature. Your hypothesis could be in the following form: "That the concentration of ascorbic acid in solution decreases faster as the [manipulated variable] is increased/decreased".


Moreton Bay College girls prepare fresh orange juice. Well maybe not that fresh - I can see a mouldy orange at the back.

In a high school lab, the easiest way to measure ascorbic acid concentration is by titration. There are two common methods, both of which work well. The first is by DCPIP titration. DCPIP is 2,6-dichlorophenolindophenol and reacts with ascorbic acid in a 1:1 ratio. It is a blue dye that produces a nice pink end point but is quite expensive. A 1 g bottle costs about $45 (but you only need about 0.2g per litre) - available from Rowe Scientific, Brisbane Ph 3376 9411, email roweqld@rowe.com.au. See Vitamin C DCPIP titration to download the method. A second DCPIP method that has been trialled extensively appears in “The Laboratory - A Science Reference and Preparation Manual for Schools” by Barbara Dungey – available from Southern Biological (click to see book details). It was recommended by Karen Marchant of Home Hill State High School, Queensland. The method features a phosphoric/acetic acid extracting solution. I have adapted this method and added sample calculations. It can be downloaded by clicking the link: Vitamin C DCPIP Method 2. The one caution with DCPIP and the cause of so much misery amongst students is that DCPIP is not easy to dissolve; you need to leave it overnight and then decant or filter it the next day. The third method is by iodine titration and uses cheaper and more easily obtained chemicals. For a copy of this method (courtesy of Deb Smith, HOD Science, Centenary Heights SHS) click: Vitamin C Iodine titration. You could even compare the two methods.


The older they get the lower the concentration!	DCPIP goes in the burette	End point is a faint blue

Chlorine loss in a swimming pool - due to sunlight intensity
Home swimming pools are usually sanitized with chlorine-based compounds such as calcium hypochlorite, Ca(OCl)₂or sodium hypochlorite NaOCl, which produce the hypochlorite ion HClO^- when dissolved in the pool water. Chlorine in a pool can get consumed in many different ways, but the most common is from sunlight and aeration which convert chlorine in an oxidation state of +1 into chloride ion in an oxidation state of -1. Reports suggest that in strong sunlight, up to half of the HOCl is destroyed within 17 min. A good EEI would be to make up some pool water and add a measured amount of either calcium or sodium hypochlorite and measure the rate of consumption of free chlorine in pool water when exposed to sunlight. As a second IV you could look at the rate of loss at different pHs. The standard method for determining free chlorine is to measure the amount of oxidant by its ability to liberate iodine from acidified iodide solution. Take a chlorine-containing water sample, add an excess of KI solution to liberate free iodine which produces an indigo-blue colour formed with a fresh starch indicator. Find the amount of this iodine released by back titration with sodium thiosulfate. Click here to see a good method. The problem with the iodometric (iodine titration) method is that it takes a long time for students to collect data. Janet Grice suggests Doug De La Matter's Methyl Orange method. Her Yr 12 Pool Chemistry handout is also available. And I've attached an article from Chem Matters supplied by Janet Grice. As another IV you could look at amounts of aeration by bubbling air through it. Note the warning below!

Chlorine loss in swimming pool water - dependence on colour
Chlorine loss from pool water is known to be due to the action of sunlight (see text above). However, it is possible that the breakdown of chlorine is greater for different wavelengths of light than others. For example, does it breakdown as quickly under red light as under blue light? It would be an interesting EEI to see which colour/s have the greatest effect. You could make up some pool water with a known amount of chlorine (using Ca(OCl)₂ or NaOCl), place in a stoppered test-tube (why stoppered?) and wrap in a single layer of cellophane. You should be able to design the rest of the method yourself but you'd need several colours of cellophane and to measure the free Cl at several intervals of time (see experiment above for titration suggestions). Your problem will be to ensure the same intensity of light gets through to the solution (yellow may not absorb as much as blue for instance). The image below shows the wavelengths of light most absorbed by each type of cellophane; this is called their "l max" (lambda max), that is, the wavelength most absorbed. I did this on a spectrometer at Moreton Bay College but you could run them again if you can get access to a spectrometer. You would also need to know what % transmission (or absorbance) occurs for each colour; I didn't do that. As a second IV you could try thickness: one layer, two layers etc of cellophane to see if the response is linear. Have fun!

Chlorine loss in swimming pool water - the role of cyanuric acid stabilizer.
The biggest problem with chlorine as a sanitiser in swimming pools is that it breaks down and dissipates very easily under the sun’s radiation. This can be fixed by adding cyanuric acid. Cyanuric acid (1,3,5-triazine-2,4,6-triol) is used as a “stabilizer” for chlorine in swimming pools and stops it breaking down so quickly in sunlight. On a bright sunny day, nearly all of the chlorine in a pool can be lost in less than two hours unless a stabilizer (like cyanuric acid) is present. The addition of about 30 mg/L (ppm) cyanuric acid to swimming pool water reduces destruction of the free chlorine by sunlight. In the stabilization process, a portion of the chlorine residual is temporarily bonded to the cyanuric acid molecule which protects the chlorine from the destructive effects of sunlight. The nature of this bond is such that the chlorine continues to be released as long as a demand exists. The ideal level of cyanuric acid is 30-80 mg/L but no more than 100 mg/L (100 ppm) as a maximum. An interesting EEI would be to assess the ability of cyanuric acid to prevent the degradation of chlorine in water when exposed to sunlight. Perhaps you could add solutions of varying concentrations of cyanuric acid (eg 0 to 100 mg/L) to some water which has chlorine present (maybe 10 mg/L) and put it in the sun (or fluorescent light) for so many hours. At the end, you could measure the concentration of chlorine and see if there is a relationship between loss of Cl and concentration of cyanuric acid.

Secondly, you could take this EEI further (but this will be harder): perhaps of the cyanuric acid breaks down too as it tries to prevent chlorine loss. It is not supposed to but you could check. The test for cyanuric acid is a reaction with a melamine solution which forms a fine, insoluble, white precipitate (melamine cyanurate) that causes the water to cloud in proportion to the amount of cyanuric acid in it. For a school chemistry EEI you could buy a cyanuric acid test kit from a pool shop (about $30). The kits have a range of 20-100 mg/L but in 10 mg/L increments - which is not that accurate. There are two main types, one (called "the disapppearing dot") where you mix an equal volume of the test solution (see photo below, right) with your sample and add it to a graduated tube until you can no longer see a black dot on the bottom. The level of the liquid when this happens gives a reading on the side of the tube in mg/L of cyanuric acid. The second method is where you add your water to a tube (see photo below, centre) and add a melamine tablet and crush it. The solution will go cloudy and you raise the black dot on the bottom until you can just see it. A scale is on the lifter is graduated in mg/L. To be more accurate, you could prepare a set of standard cyanuric acid solutions and measure their turbidity in a spectrometer (l max = 420 nm) after the test solution or tablet is added from the kit. This will allow you to prepare a standard curve from which your experimental solutions can be compared. A method for this was published in the Water Research journal. Click here to download an extract.


Label reads "Active constituent 996g/kg cyanuric acid".	Cyanuric acid test kit. Tablet type. Click image to enlarge.	Test kit label. Solution type. Click image to enlarge.


You need to decide what sort of container to use. Glass (container B) and PET plastic (container C) absorb in the UV region but enough radiation should get through. The solution in an open container (A) gets the light without any absorption.	Graph of hypochlorite concentration vs time for four different amounts of cyanuric acid. Courtesy of Mitchell Oxley, Yr 12 EEI, Redlands College, Australia.

Chlorine loss in swimming pool water - due to urine and perspiration
Chlorine, used as a sanitiser in swimming pools, may be consumed in several ways: by bacteria, decomposition by sunlight and by oxidizing nitrogenous compounds such as ammonia (NH₃) and urea (NH₂)₂CO introduced into the water as components of perspiration, urine and other bodily excretions. This suggests a good EEI. You could investigate the effects on added urea on the hypochlorite ion levels in water over time. Normally pool water is kept at about 3 - 5 ppm chlorine but you could start with something a lot higher. You’ll need to find a method for measuring chlorine – best done for an EEI by titration as there is more to discuss (see above).

Effect of copper on the growth of algae (may be more suited to a student doing Biology).
The last thing you want in your swimming pool is algae – the green plant that grows on the walls and bottom of the pool. There are several ways to control it: keeping the sanitizer (chlorine) levels correct helps but often a copper-based algicide (algae killer) is used. The copper ion (Cu²⁺) is a very effective algicide to both kill and prevent algae formation. Swimming pool companies say that about 0.03 to 1.0 mg/L (0.03 to 1.0ppm) of free copper ion must be present to be effective and safe. The word "free" is used because “bound” copper (copper is tied up in an insoluble form) is not available to work as an as algicide. For non-biological systems (where no living plant or animal is present) a continuous level of 1.0 ppm is enough to assure effective algae control; more is superfluous and may damage surfaces and equipment. The toxicity of copper to algae has been the subject of a number of studies over the past 40 years because of its widespread use for the control of algae in natural waters. This suggests a good EEI. You could try growing algae in solutions with different copper ion concentrations from say 0 to 1 ppm. One problem you will have to sort out is how to measure the amount of algae in the samples. Perhaps it can be done using a spectrometer, or by measuring the depth at which you can just see a black cross appear/disappear (like in a simple nephelometer tube they use in Biology, or like the Secchi Disk method for turbidity in natural waters). Safety note: copper is a heavy metal ion and is considered hazardous. It is important that you become aware of the risks. Care should be used when handling this product.


Algae growing in nutrient (no copper).	Copper sulfate solutions are blue - but only at higher concentrations such as the 500 ppm of Cu²⁺ ions shown here. You can just see some blue at 50 ppm but not at lower amounts. In your experiment all solutions will be colourless.	In these test tubes, algae has grown faster in the left-hand tube than in the right. The problem is - how do you measure the amount?

An interesting study by Drs Jenny Stauber and Mark Florence from CSIRO’s, Division of Energy Chemistry, Lucas Heights Research Laboratories, Sydney, Australia found that copper ions depressed both cell division and photosynthesis in many species of algae notably the common freshwater green alga, “Chlorella” (Chlorella pyrenoidosa). Reference: J. Stauber and T. Florence, ‘Mechanism of toxicity of ionic copper and copper complexes to algae’, Marine Biology 94, 511-519 (1987).

In their experiment they maintained Chlorella pyrenoidosa in MBL medium on a 12 hour light: 12 hour dark cycle (Philips 40 W fluorescent tube, white, 6500 K - see photo below) at 21ºC. They found that a Cu²⁺ concentration of 7.9 x 10^-7M (5 x 10^-5 g/L, equal to 0.05 mg/L or 0.05 ppm) gave a 50% reduction in growth. Click here to see what the MBL medium consists of (this may be too complicated for high school EEI). One question you need to sort out is how to measure algae growth (perhaps measure the absorbance in a spectrometer). If you don't do Senior Chemistry you may need to brush up on your formulas for amounts and concentration. The copper sulfate your school lab has is most probably copper sulfate pentahydrate (CuSO₄•5H₂O). It has a molar mass of 249.5 g/mol. Copper itself has a molar mass of 63.5 g/mol. Thus, to make a 1000 mg/L Cu²⁺ solution (1000 ppm) you would have to weigh out 1000 x 249.5 ¸ 63.5 g of CuSO₄•5H₂O per litre of distilled water (3.929 g/L). Make sure you use distilled water as tap water will go cloudy. You can then do serial dilutions (1:10) to reduce this to 100, 10, 1, 0.1 ppm Cu²⁺, and from there you can make the solutions you want.

A fluorescent tube (watch the spelling). The 6500K is an indication of the whiteness of the colour. It is said to be the "colour temperature" and is measured in the unit "kelvin" (K). It doesn't mean it reaches that temperature though (= 6227ºC); it is just the colour given off by an object at that temperature. Temperatures over 5000K are called cool colors (blueish white), while lower color temperatures (2700–3000 K) are called warm colors (yellowish white through red).

Note: this EEI suggestion (above) may be more suited for Biology students. Click here to go to the Biology EEI webpage where it is discussed, plus other algae and copper suggestions.

Effect of availability of copper on the toxicity of a copper algicide (may be more suited to a student doing Biology).
When you look at algicides in a pool shop, they most likely will have two types of copper-based solutions for sale: a 5% CuSO₄ solution (5 g copper sulfate pentahydrate, CuSO4•5H₂O per 100mL) and one that is an organic complex of copper – usually called “chelated” copper (about 7% Cu). These copper complexes ( such as copper alkanolamine complex) are said to be more toxic to algae than ionic copper because both the metal and the ligand (organic part of the molecule) are introduced into the cell. For an EEI you could compare the algicide ability of both forms at the same concentration of Cu²⁺. This may be more suited to a Biology EEI, or at least if you study Biology as well.


Ionic copper (40g/L copper sulfate pentahydrate). Click image to see closeup of the label. Click here to see label on back of bottle. Price $25/L.	Chelated copper is a better algicide as it keeps working for several weeks whereas copper sulfate is only good for a day or so. Pool water usually contains a high concentration of carbonate ions, so the copper ions in CuSO₄ will react quickly with the carbonate ions and form an insoluble precipitate of copper carbonate.	You can also get non-copper based algicides. The one above ($40) contains benzalkonium chloride a well-known disinfectant used in Dettol. Click image to see close up of label.

Precipitation of copper carbonate in swimming pools
To kill algae in a swimming pool either ionic copper (in the form of copper sulfate) or chelated copper can be used (see above). Pool manuals and pool chemical suppliers say that the problem with using the ionic form - copper sulfate pentahydrate, CuSO₄•5H₂O - as the algicide is that it doesn't last long in pool water. Pool water has carbonate ions (CO₃^2-) present from the addition of sodium carbonate or sodium bicarbonate as a buffer against pH changes. The carbonate ions react with the added copper sulfate to form a precipitate of copper carbonate: Cu²⁺(aq) + CO₃^2-(aq)→ CuCO₃(s). So it would appear that any copper ions added to pool water would immediately be precipitated as the carbonate and thus not available to kill algae. But pool chemical suppliers say that the copper ions work for several hours which is enough time to bust open the algae's cells and kill them. This suggests a fascinating EEI. You could look at the rate of precipitation of copper carbonate in aqueous solution (pool water). The solubility product (K_SP) of copper carbonate at 25ºC is 1.4 x 10^-10. The equation K_SP = [Cu²⁺(aq)] [CO₃^2-(aq)] = 1.4 x 10^-10 means that if the ionic product of [Cu²⁺(aq)] and [CO₃^2-(aq)] is greater than 1.4 x 10^-10precipitation will occur. Pool water typically has a carbonate ion concentration of 100 ppm (mg/L) expressed as CaCO₃. Using relative molar masses this means the actual carbonate ion concentration is 60/100x100 = 60 mg/L or 0.0001 M. You do the maths! It is also recommended that the copper ion concentration in pools be about 0.05 ppm (mg/L). This is about 0.0001 M. The ionic product is 1 x 10^-8 which is greater than the K_SP so a precipitate should form. However if both the copper and carbonate concentrations are 0.00001 M each the Trial Product is 1 x 10^-10 which is less than K_SP so no precipitate should form. You could try various (equal) concentrations of Cu²⁺and CO₃^2- and examine the turbidity of the resulting solutions using a spectrometer. I would suggest a wavelength of 400 nm (although Balch recommends 560 nm for turbidity (R. T. Balch, Measurement of Turbidity with a Spectrophotometer, Ind & Eng Chem Anal Ed V 3, no. 2, p124-5). I got higher absorbances at 320 nm (UV) and 820 nm (near IR) but these may have been artefacts of the instrument (the plastic in the cuvettes absorbs strongly in the UV). Most importantly, you could see how the turbidity varies with time (perhaps every hour or every day) as the pool chemists suggest. If you don't have a spectrometer you could look at settling rates of visible precipitates (> 0.0025 M solutions) as a function of concentration, temperature or pH.


In a clear pool you can easily see the bottom (photo above). When suspended solids are present it looks cloudy.	Cloudy pool water can be caused by a number of things: live algae give it a cloudy green tinge. If the algae is dead and bleached it can look cloudy blue (photo, right).	But cloud can be caused by other things: suspended calcium carbonate from incorrect pH, or suspended copper carbonate if a copper sulfate algecide is added (as in the above photo).


Mixing of copper sulfate (already in the test tube) and sodium carbonate solution (from the pipette) precipitates (light blue) copper carbonate. The last image looks like the pool water above right. Source of Photos: Journal of Chem Ed.

A mixture of equal volumes 0.1 M CuSO₄ and 0.1M NaCO₃ gives a brilliantly coloured precipitate which settles (this much in 20 minutes).

When equal volumes of CuSO₄ and NaCO₃solutions are mixed a precipitate will form if the ionic product is less than K_SP. The concentration of the two solutions before mixing in the spectrometer cuvettes shown above are: (1-3) 0.001M, 0.0001M, 0.00001. Cuvettes 4 and 5 are the distilled water blanks. The precipitates (in cuvettes 1-3) are not visible to the naked eye but gave absorbances of 0.180, 0.067 and 0.000 at 320 nm.


A spectrometer is the most accurate way to detect the small amounts of precipitation at low concentrations. This Moreton Bay College student is using a donated Unicam SP1700.	The absorbance of the 0.001M solutions was 0.180 at 320 nm. Notice how the more concentrated solutions in the test tube rack on top of the spectrophotometer have already settled.

If a spectrometer or turbidimeter is not available a nephelometer will suffice (though not as accurate). The units for a nephelometer are nephelometric turbidity units (NTU) and the relationship between NTU and absorbance (A) is approximately: NTU = 0.19 + 926 x A (at 750nm). A reading of 80 NTU corresponds to an absorbance of 0.086 at 750 nm.


The black wavy lines on the bottom of nephelometer tube can be clearly seen when it is empty.	With distilled water present the wavy lines are still easy to see.	With a cloudy carbonate precipitate added the wavy lines are harder to see.	Looking through the top of the nephelometer tube the wavy lines can just be seen. We can record the turbidity as 80 NTU.

Getting PAM to clarify dirty water
You may have seen the ad for the World Vision charity where the little African girl is holding a plastic bottle full of dirty water for drinking. Communities like hers benefit from clean drinking water and one way to achieve this is through sanitising (with chlorine) and clarifying - using a “coagulant” that causes the suspended particles to coagulate (come together) and settle to the bottom ("flocculation") and are big enough to filter out. Two coagulants/flocculants commonly used for water are alum (aluminium sulfate Al(SO₄)₃•nH₂O, where "n" is usually 14 or 18). This is an inorganic flocculant and is discussed in the next suggested EEI further down this page. The other type of flocculant is the organic polymer type and the most common of these are called cationic polyelectrolytes. Two examples are polymer polyacrylamide (PAM) and poly di-methyl di-alloyl ammonium chloride (PolyDADMAC). The process is simple: the coagulant is added to water mixture and is then slowly stirred in a process known as flocculation. This water churning induces particles to collide and clump together into larger lumps, or “flocs.” The coagulant works by creating a chemical reaction and eliminating the charges (negative or positive) that cause particles to repel each other. The process requires chemical knowledge of source water characteristics to ensure that an effective coagulant mix is employed. Improper coagulants make these treatment methods ineffective. These polyelectrolytes are not only used for drinking water; they are used in industry for such applications as clarifying paper mill wastes and dewatering primary and secondary activated sludges. This suggests a good EEI.


Every year in Ethiopia, about 250000 children die because of waterborne illness and sanitation-related issues.	PowerFloc and Polysheen Plus are both cationic polyacrylamides.	You can get the flocculant polyacrylamide in either a cationic, anionic or non-ionic form. This one from a pool shop is cationic.

You could make up some “dirty” water by adding clay – not too much, maybe 1 g per 5 litres – and adding an organic polymer flocculant such as PolyDADMAC or PAM – and giving it a good stir. I'd suggest using some terra cotta clay from the Art Department of your school as it is a nice red brown colour and the floc easy to see. Don't let them give you "paper clay" as it has ground up paper that stuffs things up. You’d need at least five different amounts and probably duplicates or triplicates (trials) of each. Your problem is also to control the variables: how long to stir for, how fast to stir, how long to allow settling, what to measure (height of floc, turbidity of “supernatant” liquid (clear liquid above the floc). Other variables to try: temperature, pH, salinity. Instead of clay, you could use CaCO₃, BaSO₄ or limewater Ca(OH)₂. What might be a good Research Question? It is not much good just saying "do organic polymer flocculants clarify dirty water?" - because you know they do. Perhaps "Is PAM or PolyDADMAC better at clarifying muddy water?" At some stage you need to develop an hypothesis. Perhaps there is an optimum amount of flocculant that can be used (too much or too little doesn't work at well). You'd also need some theory to support your hypothesis. Where to get the flocculants? The cationic PolyDADMAC is available as Focus Brand "Water Polish" or many other proprietary names. Polyacrylamide (PAM) is also available from pool shops with brand names like Aquatic Element's Aquatic Clear Advantage, Premium Quality's Ultimate Clarifier, Bioguard's Polysheen Plus, or PowerFloc. Note: don't get mislead by the internet: polyacrylamide is also available from the gardening section of a hardware shop as “water retaining crystals” – brand names like Hortico (see below); and surprisingly also available from shops that sell disposable nappies (eg Huggies, Pampers) which have the polyacrylamide as the water absorbent (rip one open). However, these forms of polyacrylamide are NOT suitable as they are not activated and don't work - I've tried it. Stick to the pool shop product. One last warning: I had terrible trouble getting PAM to coagulate muddy terra cotta water, whereas PolyDADMAC worked quickly. Why is that? Maybe the charges on the clay ions are not neutralised by PAM. Hmmm, a good EEI.


Ultimate Clarifier is an example of a cationic polyelectrolyte known as polyacrylamide (PAM). It costs about $24 per litre but sometimes pool shops have bulk quantities that they can let you buy a few mLs of. It doesn't work well with clay turbidity. If you can read the label, you need 300 mL per 30000 litres of pool water: that's 1:100000 dilution but for muddy water in a lab you could try a more concentrated brew. In the photo opposite I added 1 drop to the 25 mL (1:500).			Water Polish is the brand name of one type of poly di-methyl di-alloyl ammonium chloride (PolyDADMAC) and this works well with clay. The photo above compares PolyDADMAC (left) with PAM (right) 10 minutes after being added to muddy (terra cotta) water.

Clarification of water with alum
Cloudy water for domestic water supplies is commonly treated with alum (Al(SO₄)₃•18H₂O). The name 'alum' is a bit confusing as there is also a double sulfate of potassium and aluminium with the formula KAl(SO₄)₂·12H₂O) commonly called 'alum'. The 'alum' used as a coagulant is the first one: aluminium sulfate - but be aware that there are several types of aluminium sulfate, each with different amounts of water of crystallization. The most common in schools is the one with ·18H₂O, sometime called octadecahydrate (M_r = 666.4). The other common one is ·14H₂O. The reason I mention this is because they will have different molar masses and this will be important when weighing it out. Alum acts as a coagulant, which binds together very fine suspended particles into larger particles that can be removed by settling and filtration. In this way, objectionable color and turbidity (cloudiness), as well as the aluminum itself, can be removed from the drinking water. By the addition of a small amount of alum to water, it can be filtered through ordinary paper without difficulty, and yields a brilliantly clear filtrate, in which there is no trace of suspended matter. If it believed that alum not only clarifies a water, but also removes disease germs and ptomaines, so its use is of incalculable value to society. A good EEI would be to make up a sample of water with suspended clayey matter and then filter it through the best filter paper you have at school. To the (still) cloudy filtrate you could add alum solution (about 20 to 1000 mg/L) to see if it settles the clay and enables you to filter the solids out (weighed filter paper). Here are some ideas for your hypothesis: try different amounts of alum (there is an optimum amount - too much alum will actually impede the coagulation/flocculation process. Try different acidity/alkalinity as the pH is a very important parameter in water treatment, especially for effective coagulation. Each coagulant has a narrow optimum operating pH range. For example, alum tends to work best at a dosed-water pH of 5.8-6.5). Aluminium sulfate should be readily available at school - if not then go to the pool shop. Remember to include the •18H₂O (or whatever) in the formula when working out the molar mass. A great EEI with great social importance.


Commercially available "alum". Sometimes the label says "aluminium sulfate" sometimes - like the Focus pack - it says nothing and you have to consult a MSDS on-line.		Premium Quality brand "Floc-out" is aluminium sulfate. The enlargement of the label says 400 grams per 10000 litre pool.


Alum added to muddy (terra cotta) water. Initially (left), after 10 minutes (centre) and after 1 day (right).			Contains a solution of alum. The concentration (actually density) is measured with a hydrometer.

Reaction Rate and Surface Area
Controlling reaction rates is one of the great challenges facing scientists and engineers in modern day life. You put food in the refrigerator to slow down decay, and you use hot water to wash up as fat reacts faster with detergent in hot water. Temperature is one way - but so is the control of surface area: fine sugar dissolves faster than coarse sugar; sawdust burns quicker than a lump of the same wood; drugs are made with different particle sizes to control the speed of release into the blood stream. This seems like the basis of some great EEIs. In Year 8 you probably did a science experiment with Alka Seltzer tablets to see what factors affected how fast the tablets dissolved. You would have looked at temperature (tried hot and cold water), and surface area (whole tablets versus crushed up ones). However, for surface area you would not have done it quantitatively (numerically, by calculating the surface area using a ruler). This suggests a great EEI but if you plan to do it with Alka Seltzer tablets the chances of getting an "A" will be greatly limited by how well you can control the variables. If you were to try it as an EEI you could try break up tablets into 2, 4, 6, 8 pieces and measure how long they take to dissolve. You could measure the sides of the chunks with a Vernier calliper and calculate the surface areas. However, the reaction time for a whole tablet is about 52 s, and for one broken into 8 pieces - about 50 s. A completely crushed tablet - a powder whose surface area is immeasurable - takes about 23 s. So that's why the EEI will be off to a bad start. Lastly, as the tablet dissolves and the chemicals react, its surface area decreases so that factor is no longer controlled. As well, as it dissolves it breaks up into small pieces so you have another problem, and it is hard to see when it is all dissolved as there are bits of gunk floating around. Lastly, sometimes the tablet floats on top of the bubbles on the surface of the water and the tablet can't get to the water.


Think of the tons of these that must get used in school experiments.	Work out the total surface area.	Cut them up and measure all pieces accurately to calculate the initial surface area.


Get the stopper in fast. Could always use chilled water (in all) to slow things down. But how will you stop splashing over the top? And how do you cope with the decreasing surface area?	So far this one has produced 0.061g of CO₂ gas. Tare the flask and tablet (to zero) before you mix them.

The solution may be to try a different surface area experiment (see below):

Reaction of marble and acid
You may have also tried this using marble chips and calcium carbonate powder. But how to measure the surface area? My suggestion is to get some marble tiles from a tile shop and cut them into strips with a masonry blade on an angle grinder. If you have five strips you can break one in half, one in quarters and so on. Using a Vernier calliper it will be easy to measure surface area. The method id quite straightforward after that. Nevertheless, give some thought to how you will control temperature, and how much acid you will need, and what concentration so that it is not the limiting reagent. Lastly, what will you measure for reaction rate: time taken for the chips to dissolve, or a flask on a balance with recordings taken every minute...and so on. A great technique for an EEI would be to pour off the acid after a given time and titrate it with standardized NaOH solution to see how much acid (and hence marble) was used.

Cut a marble tile with a masonry blade into several long strips
say 10 cm x 1 cm x 1 cm Make measurements with a Vernier calliper

Break the strips into 2, 4, 6 or 8 pieces.
I used a cold chisel to cut this one into 8. Measure all the pieces. Let them react with acid.

Zinc and acid
This is a lot simpler as you can cut zinc sheet with scissors. You'll still need Vernier callipers (American spelling: caliper) to measure the dimensions, so the error will be larger than with the marble but it may be a lot simpler to make sense of. Hmmm, what concentration acid will I need? Will it heat up?

Displacement Reaction
This has the potential to be a great (not just a good) EEI. More reactive metals will displace less reactive ones from solution. If you've done the Redox unit in Chemistry you will be aware that a reactive metal like zinc, when placed in an aqueous solution of a salt of a less reactive metal (eg Cu as CuSO₄ solution) a reaction will occur. The zinc will dissolve to form Zn²⁺_(aq)ions, and the Cu²⁺_(aq) ions in the CuSO₄ solution will accept electrons from the zinc to become copper metal. The solution starts off as a bright blue colour due to the presence of Cu²⁺_(aq) ions but as these are consumed the solution gets less and less blue. If you have access to a spectrometer then this would be easy to measure. You could make up a solution of known concentration (recall that CuSO₄ is actually in the pentahydrate form CuSO₄ · 5H₂O when working out molar masses). Just measure the absorbance of it (the lambda max for the copper solution is 740 nm) on the spectrophotometer and from a graph (assuming the Beer-Lambert Law is still working from 1852) and if you know the absorbance you can work out the concentration of the copper ions. Your teacher may suggest that you prepare a range of standard solutions of Cu²⁺ to produce your own calibration curve. Now, using a zinc strip and ones cut into halves, quarters and so on (all measured with Vernier callipers) you can place them in identical copper sulfate solutions and measure the change in blueness after say an hour or a day. What a great EEI. I wish I was young again - I'd do it.

After 1 minute the two solutions have different Cu²⁺ concentrations shown by the different intensities of blue colour. This means more Zn has reacted in one than the other. You can see the brown copper metal that has been displaced out of solution by the zinc. You don't have to use zinc - any more reactive metal than copper would work. You'll need one whose ions are colourless. There are lots of combinations you could try (eg Mg + Cu²⁺).

For copper ions, set the wavelength to 740 nm. Different metal ions have different values for lambda max. The filtered Cu²⁺ solutions can be analysed in a spectrophotometer. My thanks to Moreton Bay Boys College for access to their instrument.

Corrosion of Copper by Sulfuric Acid
It may seem surprising but there are almost no journal articles by chemistry researchers on the effect of surface area on reaction rate - in industry or academia. Those that do relate to the area of catalysts rather than the main reactants (but that does suggest another EEI topic). The most recent paper as a stimulus for a high school chemistry EEI is one by industrial chemists Glenn Damon and Ray Cross from the Michigan College of Mining and Technology, Houghton, Michigan published in Industrial and Engineering Chemistry journal V28 (2) in February 1936. They reacted sulfuric acid with small squares of copper placed 2 cm under the liquid surface. However, to manipulate the surface area variable they varied the surface area of the solution exposed to the atmosphere. You could prepare a small circular piece of polystyrene foam (with a hole cut in the middle) and float it on the surface of the acid. This will give limited access of oxygen to the solution and hence limit the corrosion of the copper. It is a neat experiment and may give you a few ideas. Click here to download it.

Cetyl alcohol and water evaporation losses
Billion of litres of water normally lost each year through evaporation from the nation’s waterways – including reservoirs, lakes and dams. Evaporation from Australian water bodies ranged from 1.3 m to 1.9 m per year (Brisbane is 1.6 m per year), with an average evaporation rate of 0.5 litres per day per square metre, or 5000 litres per day for every one hectare of open water. For a water body covering about 100 hectares of open water – a medium sized reservoir or dam – approx 190 million litres of water (or 75 Olympic-sized swimming pools) is lost every year through evaporation – this is equivalent to the annual consumption of 380 typical Australian households. However, recent trials in Australian conditions by several council/municipal water managers and commercial cotton farms confirmed evaporation savings of about 30% using various long-chain alcohols. These alcohols such as cetyl alcohol – also known as hexadecanol, CH₃(CH₂)₁₅OH – develop an invisible film (or monolayer) on the water surface, creating a barrier that limits the escape of water vapour. Chem-Supply in Australia have cetyl alcohol (Code: CL044-500G) for about A$39 per 500 gram bottle Lab Reagent (LR) grade (plus $27.50 for 3-5 day delivery) but a commercial grade is also available elsewhere (but in big quantities). This suggests a good EEI. You could look at the evaporation of water from an open container, with and without a monolayer of long-chain alcohol. The independent variable could be the amount of alcohol (per sq metre) or the thickness of the film, and the dependent variable could be the amount of water evaporated. To get reasonable evaporation rates you really need to use an electric fan blowing across the top of the water (for at least 24 hours). How you measure the change in water level is up to you (by mass, by height). You'd get even better results if the ambient temperature was warm (eg in a fume cupboard with the heating lamps on). Research being done by Ian Craig, Erik Schmidt and Michael Scobie from the National Centre for Engineering in Agriculture (NCEA), University of Southern Queensland (USQ) into the use of these monolayers can be downloaded here. I based the idea above on an article "Alternative methods for the reduction of evaporation: practical exercises for the science classroom" by Peter Schouten and colleagues from Griffith University's School of Engineering, Gold Coast, Australia. Peter has allowed me to make it available for download here. It was published in Physics Education (2012, V47, No 2, p 202-210)


WaterSavr, the cetyl alcohol based chemical monolayer, floating on the surface of water in an experiment.	WaterSavr, the cetyl alcohol based chemical monolayer, has just been distributed over the surface of this dam at Toowoomba, Queensland. The hose shown is part of a new process being trialled by scientists at the University of Southern Queensland.	Pure hexadecanol (cetyl alcohol) monolayer on the surface of water.

Heating up gases
You would have seen how gases expand when they are heated. Your teacher may have heated a flask with a balloon on the top to show it expanding; you may have seen a balloon shrink when dipped in liquid nitrogen at -198ºC; and it is the principle behind how hot air balloons work. In class you would have called the law describing the relationship between temperature and volume Charles's Law or perhaps Amonton's Law (V µ T when T is in kelvin and P and n are kept constant). There could be a great EEI in revisiting this relationship. There is no point in just verifying it as this has been done a million times. What you want to do is to extend the investigation of this law to look at the impact of changing variables and to consider allowing for errors. The diagram below shows a setup that may be useful. It really just show the connection of two things: a flask with a sidearm (maybe a Büchner flask) and a graduated glass syringe. The exact positioning is something you should determine. Glass syringes are precision-made with low friction between the plunger and the barrel (unlike plastic ones that have high friction). Your should have some in the chem lab and if not they are reasonably cheap (about $50 for a 100 mL one). You need to introduce a gas (eg CO₂) into the flask and surround the flask with water in a beaker on a hotplate. As it slowly heats (I mean slowly, maybe 20ºC to 80ºC over 40 minutes) the gas expands and the syringe is pushed out. With the syringe on it's side there is no need to worry about the weight of the plunger. You could compare gases - oxygen, nitrogen, hydrogen for example. But how to get samples of these gases? You may have cylinders but you could produce H₂ and CO₂ by reaction (or let some dry ice sublimate); let some liquid nitrogen evaporate (or remove oxygen from air). And why not propane (BBQ gas) or butane (cigarette lighter fluid)? Remember that balloon gas is not just helium - it has 3% air mixed in with it. The main point is that the law holds for ideal gases but at atmospheric pressure and room temperature they won't be that ideal. And is the deviation from ideality dependent on the molar mass of the gas, or whether it is polar or non-polar, and where on earth do you get a polar gas from (HCl is too dangerous)? What range of temperatures will you use (consider liquid nitrogen, dry ice). What value will they give you for absolute zero when the V/T graph is extrapolated? How do you draw the line of best fit (is least-squares the best, does it give you the most accurate value for absolute zero?). And what is the volume of the gas in the apparatus? And what is the best way to measure temperature (of the gas as in the diagram, or of the water surrounding it)? Perhaps the temperature of the gas in the flask is the water temperature and the temperature of the gas in the syringe that of the surrounding air (work out a weighted average). And how do you control atmospheric pressure (do you have a barometer, or perhaps get the data from the meteorological bureau website). What a fabulous EEI. I must put this on the Physics EEI webpage as well.

Aging (fermenting) orange juice
Here's a comment off a health food blog from a guy called Vincent: "I was too lazy to wash out a 2 L carton of Tropicana orange juice after dinner last night. I go to wash it out today and the carton was bulging quite noticeably. Those crazy orange juice fermenting bacteria work fast! The carton let out a nice puff of air when I opened it up and it tastes so sour." What has happened here? Orange juice has a lot of natural sugars in it. Bacteria love it if you let them get in. The refrigerator only slows growth of bacteria, it doesn't kill them. These bacteria aren't necessarily the kind that make you sick, but they will start to grow and will begin to break down the orange juice. It will start to ferment-if you taste it it will be bubbly and will taste sour from the build up of acids - possibly acetic acid from the alcohol. Is there an EEI in this? There certainly is and it needs careful consideration about controlling variables and you need to think about what acids are present besides citric and ascorbic. What variable might you manipulate? The total acidity can be measured by titration with sodium hydroxide.

Year 12 students from Moreton Bay College - Bianka and Cassie plan to measure the titratable acidity of orange juice.

Cheesemaking.
Here's one from Gary Turner at St Mary’s Catholic College South Burnett. Most major newspapers have a life-style section in which appear columns about cheeses and wines. Australia has several small cheese-making plants in which hand-craft is as important as technology. Cheese-making is a promising industry within the local region. A closely related product, amenable to student-investigation is the making of sour-cream, which is commonly used in several fast-foods of interest to teenagers. First: You will be following a ‘standard’ procedure for making a simple cheese (e.g. ricotta) or sour cream to give you the background skills and chemistry involved in making a cheese, and to explore the factors involved. (This can be done as a group). Second: You are then to select another cheese that interests you and individually make this cheese and explore the factors that affect the result (e.g taste and texture and hardness). This section of the work will also require you to define which factors you can reasonably test in a school-laboratory, and which variables in the production that you can vary. Third, you are to compare your cheese to a similar commercially available cheese and report on the differences and likely causes of that difference. (The factors that you can compare will be those that you have defined in the second section above). A copy of this cheese EEI is available for download here. A useful video from the ABC TV Landline program maybe worth watching. It shows Yr 12 science students from Sandgate Sate High School (Queensland) making cheese under the guidance of master cheesemaker Russell Smith and Chemistry teacher Alison Turner. The link to the video is here. Another Landline report shows the students entering their cheeses into the Royal National Association show ("The Ekka"). See "Lateline Masterclass" here. Remember - making cheese does not make an EEI.


Cheesemaker Russell Smith instructing students.	Stills taken from You Tube video.

Ginger Beer Fest
Ginger beer is made traditionally by the yeast fermentation of a mix of sugar, water and ginger. It is rarely produced commercially but often home brewed. The beverage produced industrially is generally not brewed (fermented), but carbonated with pressurized carbon dioxide. It is really just a soft drink, sweetened with sugar or artificial sweeteners. However, there are some manufacturers who still brew it the old way: in Queensland, Bundaberg Brewery produces an excellent brewed ginger beer. It is cloudy and if you hold the bottle up to the light and you'll see it's full of ginger pieces. A good EEI would be to brew your own at home using one of the many recipes available on the internet. At St Mary’s Catholic College South Burnett, chemistry teacher Dr Gary Turner suggests this for his Year 12 EEI: First: You will be following a ‘standard’ procedure for making a simple beer (e.g. two-day ginger-beer) to give you the background skills and chemistry involved in making a beer, and to explore the factors involved. Second: You are then to select another recipe(s) that interests you and individually make the beers and explore the factors that affect the result (e.g taste and alcohol-content). This section of the work will also require you to define which factors you can reasonably test in a school-laboratory, and which variables in the production that you can vary. A copy of the EEI task sheet is available for download here.

Lactic acid and the fermentation of milk
Lactic acid forms in milk due to the action of fungi and bacteria acting on the lactose sugar. The most important lactic acid producing bacteria is Lactobacillus. The presence of lactic acid, produced during the lactic acid fermentation is responsible for the sour taste and for the improved microbiological stability and safety of the food. A good EEI might be to investigate the factors influencing the rate of formation of lactic acid upon the addition of some starter bacteria (eg plain yoghurt). I won’t say what they are but a couple of the following are suspects: heat, amount of bacteria added, light, access to air, shape of container, sugar concentration, initial pH, amount of fat (normal, low fat, skim), degree of agitation, and so on. The acidity in milk is sometimes measured by titration with a 0.1 M NaOH solution, and indicates the consumption of NaOH necessary to shift the pH-value from 6.6 (corresponding to fresh milk) to a pH-value of 8.2 - 8.4 (phenolphthalein end point). People sometimes wrongly assume that the titratable acidity is due to lactic acid - an organic acid with the formula CH₃-CHOH-COOH. However, fresh milk contains practically no lactic acid and the consumption of NaOH is used to change the pH-value of the following components: carbon dioxide, citrates, casein, albumin and phosphates which gives the appearance of a lactic acid concentration of about 0.13% The determination of "acidity" in fresh milk by means of titration is therefore more a measure of the buffer action of milk than anything else. If you try to calculate the theoretical pH of milk based on the titratable acidity (using the K_a for lactic acid), you will get stupid results - like a pH of 2.5 for milk.

In an EEI, it is likely that you want to talk about the “developed acidity”, which is the result of bacterial activity producing lactic acid during milk collection, transportation, and processing. In order to avoid the uncertainties about the degree of titratable acidity or developed acidity, it is necessary to use a different method for determining lactic acid. A rapid colorimetric method for the quantitative estimation of lactic acid in milk is available but way beyond the facilities of a high-school lab. The only way out of this conundrum is to measure “titratable acidity” (rather than calling it "lactic acid concentration") but acknowledge the errors and subtract the initial “acidity” from the subsequent values obtained during the experiment. Be careful if you intend to measure titratable acidity as a function of time eg "time elapsed" (rather than just as a function of some manipulated variable (such as temperature). See the note that follows.

Note about identifying variables: "time elapsed" can be a controlled variable or independent variable (or both) in this experiment (and others that involve collecting data over a period of time).

CASE 1: In the fermentation experiment you may, for example, choose to have the "temperature" as the independent (manipulated) variable (say 0°C,10°C, 20°C, 30°C...) and "titratable acidity" as the dependent variable. If these are measured just once, say after 1 week, then "time" is a controlled variable (along with initial pH, sunlight, sugar concentration, aeration, exposed surface area etc). You could prepare a graph where you plot "titratable acidity" (y-axis) and temperature (x-axis) and there will be one line.

CASE 2: However, "time" can be an independent variable as well. You use the "temperature" as the independent variable but if you measure the dependent variable (titratable acidity) every week at 0, 1, 2 and 3 weeks then you really have two experiments in one. There are two independent variables: "time" and "temperature" but they can be examined separately. A plot of titratable acidity (y-axis) vs time (x-axis) would show 4 lines (if you used 4 different temperatures). This would be most valuable as it would show you the fermentation rate at each temperature. You could prepare another graph where you plot titratable acidity (y-axis) and temperature (x-axis) to get 4 lines (one for each weekly measurement including the titratable acidity at t=0). This would be harder for you to visualise and interpret however. The two graphs together could be analysed "... to identify relationships between patterns, trends..." IP3 (VHA) and "analysis and evaluation of complex scientific interrelationships" (EC1, VHA). The two graphs provide stronger evidence for inter-relationships than either graph alone.

Discharge of a lead accumulator car battery.
A car battery is also known as a lead accumulator or lead-acid battery as it consists of lead, lead oxide and lead sulfate with an electrolyte of sulfuric acid. During discharge the following reactions occur:
Anode Reaction: Pb(s) + HSO₄^-(aq) → PbSO₄(s) + H⁺(aq) + 2e⁻
Cathode Reaction: PbO2(s) + HSO₄^-(aq) + 3H⁺(aq) + 2e− → PbSO₄(s) + 2H₂O(l)
There is an increase in the concentration of H⁺(aq) ions during this discharge and this can be monitored by titration with a base such as sodium hydroxide. To discharge the battery rapidly but steadily the students Olivia and Kayla at Moreton Bay College used 12V car lightbulbs across the terminals. They asked - is the change of [H⁺] proportional to the duration of discharge? Perhaps the rate of discharge as well as the duration important. Should they monitor voltage and current as well? You decide. My thanks to their teacher Mrs Cathy King for welcoming me into her lab.


Battery is off as they standardize the sodium hydroxide.	Battery is discharging through two 12 V lamps in parallel.

Thermal stability of sodium bicarbonate (why my cake didn't rise).
Sodium hydrogen carbonate NaHCO₃ also known as sodium bicarbonate or “bicarb of soda” - is an important component in various pharmaceutical drugs (tablets, capsules, syrups) and cooking (scones, cakes). Because of its widespread use, the stability of sodium bicarbonate in solid state, both as a raw material and as a formulation component, is of high interest to the pharmaceutical and food technology scientists. When sodium bicarbonate is stored as a powder, it degrades over time to carbon dioxide and sodium carbonate after absorption of moisture at lower temperature, or degrades directly to carbon dioxide and sodium carbonate without absorption of moisture at elevated temperature (Shefter et al., 1975 - see below). Therefore, it is critical to maintain appropriate temperature and relative humidity during the storage of the raw material and finished product as well as during manufacturing. Sometimes you want it to breakdown - but only when you're ready: cooks rely on the breakdown of “bicarb” at high temperatures into carbon dioxide to make cakes and scones rise. A good EEI would be to examine the conditions for the thermal breakdown of NaHCO₃. You could do this several ways: one would be to take some samples of NaHCO₃ and hold them at various temperatures from room temperature to the maximum temperature of your oven and titrate a solution of the sample against HCl after a fixed time (eg 1 hour).

Because we have two substances in the solid mixture (NaHCO₃and Na₂CO₃) there will be two analytes in the solution of the mixture, namely HCO₃^- and CO₃^2-. The problem is: how do you tell how much of each there is. Here are the reactions so you can see what goes on in neutralisation:

Carbonate titration:
The Na₂CO₃ present in solution will react with the acid as the titration proceeds until it is all converted to NaHCO₃ (aq):

Na₂CO₃(aq) + HCl (aq) → NaHCO₃ (aq) + NaCl (aq)

Bicarbonate titration:
As you add more acid, the bicarbonate that was present initially and the bicarbonate produced by the first titration will react:

NaHCO₃(aq) + HCl (aq) → NaCl (aq) + CO₂(g) + H₂O (l)

What we need is two different indicators, one to indicate the endpoint for the reaction between H⁺ and CO₃^2- and the other to indicate the endpoint for the reaction between H⁺ and HCO₃^-. The first is phenolphthalein and the second is bromocresol green. To calculate the amount of bicarbonate in the mixture, you subtract the amount of carbonate from the total amount of bicarbonate. The titration is quite complex because of the dissolved CO₂ generated which you have to boil off. See Oliver Seeley's How to Titrate Carbonates webpage for some great hints and cautions.

ALTERNATIVES:

1. Alternatively you could titrate samples taken over a range of times (10, 20, 30…etc minutes). You’d have different graphs for the two approaches but what a great EEI. Is commercially available sodium bicarbonate different to the analytical reagent from the chem lab? Be careful if you buy "baking powder" rather than "baking soda"; baking powder has starch and sodium acid pyrophosphate added to give more gas. I have attached two scientific papers that may provide some ideas: one is Effect of relative humidity and temperature on moisture sorption and stability of sodium bicarbonate powder by Kuu, Chilamkurti & Chen (1998) from the International Journal of Pharmaceutics (1998), and another A Kinetic Study of Sodium Bicarbonate by Schefter from the journal Drug Development Communications (1975). They are heavy going but this is Year 12 Chemistry after all.


Baking soda and baking powder are two different things.	Heat the bicarb in a lab oven.

2. Another alternative is to investigate the effect of heat on a solution of sodium bicarbonate.

● CCA Treated Timber Leaching
The most common wood preservative in Australia is chromated copper arsenate, or CCA which produces a greenish colour in the wood (see photos below). Concerns sometimes arise over the use of treated lumber in vegetable beds. In the USA, CCA preservative was phased out in 2003, for virtually all residential uses, including fencing, decking, children's play equipment and raised garden beds. Two other products, ACZA (ammoniacal copper zinc arsenate) and ACQ (ammoniacal copper quat) have replaced CCA and in the USA may be used for raised bed construction. We are told that CCA, ACZA, and ACQ may be safely used to construct vegetable beds. But who would believe that. The Queensland Environmental Protection Authority ensures that when CCA fences or posts are burnt by bushfires in national parks the ash is buried as a contaminant to keep away from the public. If you have access to a spectrophotometer, you could investigate the amount of leaching from CCA pine by taking some CCA pine shavings (caution - gloves) and immersing them in water. You could vary conditions such as pH, amount of stirring, time in contact. As usual, a risk assessment will need to be submitted before you start.


CCA treated garden bed	CCA fence in a picnic ground

Anodising Titanium
Titanium is an amazing metal. It is strong, light and corrosion resistant. It can be alloyed with many metals to increase its range of applications for industrial, aerospace, recreational, and emerging markets. Its behaviour when anodised is remarkable. Anodizing titanium produces an oxide coating which generates an array of different colours, making it appealing for art, costume and body piercing jewellery and architecture (eg Guggenheim Museum). The color is an interference effect much like that in a soap bubble. The anodised colour depends on the voltage (see chart below). You could investigate the relationship between colour and voltage using different electrolytes. The big problem is getting 100 volts. Connecting a heap of 9v batteries in series might do the trick.

Iron filings in fortified cereal

A healthy adult needs about 18 mg of iron each day. Dietary iron is found in large amounts in organ meats such as liver, kidney, and heart. It is also present naturally in egg yolks, some vegetables, and shellfish. In these foods, iron is typically present as Fe (III) ions. Our body absorbs iron in the small intestine in the form of Fe (III), which then is reduced to Fe(II). Under normal conditions, our body absorbs only 5-15% of the iron in the food that we eat. Cereals are fortified with food grade iron filings as a food supplement. This iron is metallic iron (Fe). In the stomach the metallic iron is oxidized and eventually absorbed through the small intestine. You can see the iron if you pass a magnet over a slurry of breakfast cereal. I used Sanitarium's Light 'n' Tasty but any will do (see my centre photo below using a 100mm macro lens). The question is - how to make an EEI out of this? You need to do more than measure and compare the amount of iron in breakfast cereals. A good EEI would be to investigate methods of extraction of the iron perhaps involving the use of a magnetic stirring bar before analysis. Could you dissolve the metal in acid? Is all the iron in the form of elemental iron (filings) or is there some natural iron compound present?

Polyurethane foam
Polyurethane is a synthetic polymer widely used in flexible foam seating, seals and gaskets, tyres, bearing bushes, adhesives and sealants. The type you may be familiar with from school (if you made a polyurethane foam mushroom) is called a 'rigid foam' and is used for insulation panels and surfboards. They are made from two monomers - isocyanate and polyol. In 1984 water was accidentally introduced into a reaction mix and the first foam was made. A good EEI could be to look at the conditions required to produce the different densities of rigid foam. You could use equal amounts of the monomers and try them with different temperatures or different amounts of stirring. You could even try adding more water to the polyol monomer. You could even try making a variable density foam by placing the reaction vessel (plastic cup) on a cold surface. The main thing is to explain why you'd want a particular density and hypothesise how it could be achieved. It will be heaps of fun.

Alcohol-water mixture: concentrations and the contraction of volume

When you mix ethanol and water together the final volume is less than the sum of the separate volumes you started with. This shrinkage is known as 'volume contraction' and is due to the strength of the hydrogen bond. Such a bond is strong in water but weaker in alcohols, however, when a mixture is made the dipole-dipole forces tend to make the alcohol-water clusters small. Technically, we could say "departures from Raoult's law are often found in liquid mixtures resulting in volume nonadditivity". In practice, this contraction can have vital consequences. Medical researcher know that alcohol absorption into the bloodstream and the resultant volume contraction can upset the plasma concentration of various biochemicals and lead to all sorts of complications. A good EEI would be to measure the volume contraction of various mixtures of ethanol and water 25:5, 50:50, 75:25 and so on to see how the percentage contraction varies (the point of maximum contraction could be found). And if it is true that the effect is due to H-bonding, should the contraction be different for alcohols exhibiting weaker or stronger dipole-dipole forces (eg the monohydroxy alcohols: methanol, 1- and 2-propanol, tert-butanol)? I wish I was doing an EEI - this one would be great. I'd be looking at the work of our famous friend Dmitri Mendeleev (of Periodic Table fame) who found that a 1:3 mixture gives the biggest contraction (see Analytical and Bioanalytical Chemistry, 2009, V 395 (1) 2009, p7-8).

Alcohol-water mixture: temperature and the contraction of volume
In the suggestion above, the investigation of volume contraction of ethanol water mixtures was suggested. Of equal interest would be the effect of heat which is known to affect the strength of the H-bond; so you could see how stable the % contraction was over a range of temperatures. Safety warning: alcohol water mixtures can burn even when the amount of alcohol is less than 50% - and especially at higher temperatures. As well, if the contraction effect is due to H-bonding, shouldn't the contraction be different for alcohols exhibiting weaker or stronger dipole-dipole forces (eg methanol, propan-1-ol, propan-2-ol and methyl propan-2-ol)? As a matter of interest, mix together CS₂ and ethyl acetate and you get volume expansion (but CS₂ is too dangerous for high school experiments).

Capillary action
Capillary action is the tendency of a liquid to rise in narrow tubes or to be drawn into small openings such as those between grains of a rock. Capillary action, also known as capillarity, is a result of the intermolecular attraction within the liquid and solid materials. A familiar example of capillary action is the tendency of a dry paper towel to absorb a liquid by drawing it into the narrow openings between the fibers. Some liquids exhibit more capillarity than others; for example, there is a big difference between water, salt water, ethanol and hexane. A good EEI would be to compare capillary action (between two microscope slides; see below) for polar and non-polar liquids, or non-polar ones of different density, or salty water vs distilled water, or as a function of temperature or capillary gap. You could also somehow use capillary tubes (see below). The possibilities are huge, but don't get too carried away.

Browning of apples

Apples turn brown when peeled and exposed to air. This discolouration is due to a process called enzymatic oxidation and is catalysed by the enzymes present in the apples. The enzyme polyphenol oxidase (phenolase), in contact with oxygen, catalyzes one step of the biochemical conversion of plant phenolic compounds to brown pigments known as melanins (brown, like a suntan). It occurs at warm temperatures when the pH of the plant material is between 5.0 and 7.0. Browning can be stopped! Vitamin C, being a highly reactive anti-oxidant reacts with the O₂ in the air, preventing/slowing down the enzymatic oxidation of the apples. Another way to reduce browning is to lower the pH in order to inactivate the enzyme. Ascorbic acid is used commercially to prevent enzymatic browning as it acts as both an acidulant and antioxidant. To make an EEI out of this you could test the browning when controlled volumes of acids of various [H+] are used, and then with ascorbic acid of known [H+] to see how much is due to the antioxidant property. Temperature could also be assessed. It is also said that Fe and Cu speed it up but Ca²⁺ slows it down. Question: how will you measure the browning? Remember this is chemistry not MasterChef.

Effect of catalyst concentration on reaction rate: enzymes
You will have read that catalysts are substances that speed up reactions but that they are only needed in small quantities. A great EEI would be to test this proposition and to see if there is a quantitative relationship between amount of catalyst and reaction rate. Maybe once you have added sufficient catalyst then adding extra makes no difference. A good catalyst for this experiment would be a biological catalyst - an enzyme. Enzymes, like other catalysts, catalyze reactions by lowering the activation energy necessary for a reaction to occur. The molecule that an enzyme acts on is called the substrate. The enzyme molecule is unchanged after the reaction, and it can continue to catalyze the same type of reaction over and over. The enzyme catalyze will speed up the breakdown of hydrogen peroxide (the substrate) into water and oxygen. Say you took 10 mL of 3% peroxide solution, added some water, and added different amounts of catalase to each, you may get different reaction rates. The catalase can be made up into a suspension with water and different amounts added dropwise (0, 5, 10, 20 etc) to the peroxide solution. What to measure? The simplest way to start making measurements is to stop the reaction by adding sulfuric acid as this destroys the enzyme's functioning and the reaction stops. As this is a chemistry EEI should look for standard chemical methods of analysis. No doubt you've learnt about titrations, so to see how much peroxide was used up you could titrate say 10 mL aliquots of the solution against a standard KMnO₄ solution in the burette. Another method may be to use a pressure sensor in the neck of the flask and hook it up to a datalogger. If you want to examine another variable, you could hold the amount of catalase constant and vary the temperature, or vary the pH (I said before that adding sulfuric acid would destroy the enzyme, but how sensitive is it to pH).


Add the catalase dropwise, but also work out how many drops equal 1 mL.	In test tubes, the faster rate will give a bigger froth.	Titrate with KMnO₄ solution to a brown end point.

Effect of catalyst concentration on reaction rate: copper catalyst
This is a tricky one but there is so much to explore. It is a bit more complicated because it involves reaction rates and redox theory - but the method should be quite simple. When you add hydrochloric acid to zinc it reacts at a certain rate; however, if a piece of copper touches the zinc then the zinc reacts much faster. Why? And does this means copper is a "catalyst" because it speeds up the reaction (doesn't it have to appear unchanged at the end - well is it)? An electrolytic cell is created where the more reactive metal (Zn) is the anode and Cu is the cathode, with the HCl acting as an electrolyte, so hydrogen ions are likely to be reduced to hydrogen gas at the copper electrode. Without the copper, there is no obvious electrolytic cell, so hydrogen gas would have to be produced by the direct reaction between zinc atoms and hydrogen ions at the surface of the metal and this is slower. A great EEI could be investigating the effect of varying the amount of copper in contact with the zinc. A great way to get copper in close contact with the zinc is to deposit copper metal directly on to the zinc by a displacement reaction. To do this you would dip the piece of zinc into a solution of copper sulfate and a coating is immediately deposited. If you dipped identical pieces of zinc strip into a copper sulfate soution - but each to a different depth (0 cm, 1 cm, 2 cm ....) and for the same time you would get different areas of coating. Then react the strips individually with hydrochloric acid and measure the rate by whatever method you like (change in mass due to lost hydrogen, amount of gas produced, temperature change, titration the final solution against NaOH).

Zinc is easy to cut - you don't need big muscles.


Dip the zinc strip into copper sulfate solution to the required depth.	With careful measurements you can calculate the % of the surface area that has the copper catalyst (cathode) deposited.

Deactivation of pineapple enzymes
If you've ever tried to make a jelly with pineapple or kiwifruit in it you may have been sorely disappointed. It may not set because the enzyme catalyst has played up. All living cells produce enzymes which catalyze metabolic reactions. The enzyme that you could investigate in an EEI is one that is produced in pineapple and hydrolyzes certain kinds of proteins called gelatins. Gelatin used in jelly is derived from skin, bones, and/or connective tissue of animals (vegetarians have to use agar type jellies). Gelatin proteins, when dissolved in hot water and allowed to cool, form a semi-solid or gel state; hence the name gelatin (or gelatine). Hydrolyze, here, refers to breaking up the protein polymer in such a way as to prevent its forming this gel state. The hydrolyzing enzyme from pineapple is denatured (destroyed) by heat; but not freezing - I don't think. Enzymes can also be denatured by changes in pH, detergents or radiation. You could take some pineapple and subject it to different heat treatments and see its effects on the gelatine. You have to make up a device and method for testing gelatine that allows replicable and meaningful testing. I recall using the Bloom Strength test at Golden Circle - it was the mass in grams required to press a 12.5 mm diameter plunger 4 mm into the gel. If you did it at home you could even eat the results.

pH of vinegar solutions

If you've completed an acids and bases unit you will be aware that strong acids and bases, like HCl and NaOH respectively, dissociate fully. However, weak acids and bases only partly dissociate and the equilibrium constant (K_a or K_b) gives a measure of this dissociation. The quantitative behaviour of acids and bases in solution can only be understood if their K_a (or pK_a) values are known. Such knowledge finds applications in many different areas of chemistry, biology, medicine, and geology. For example, many compounds used for medication are weak acids or bases, and a knowledge of the pK_a values can be used for estimating the extent to which a compound enters the blood stream. Acid dissociation constants are also essential in aquatic chemistry and chemical oceanography, where the acidity of water plays a fundamental role. In living organisms, acid-base homeostasis and enzyme kinetics are dependent on the pK_a values of the many acids and bases present in the cell and in the body. Here's a suggestion for an EEI: if you make up a solution of known concentration of say acetic acid CH₃COOH, and then measure it's pH, you can calculate the K_a using standard formulas. No doubt there will be an error but your EEI could be to investigate the source of this error and try ways to minimize it (is it the formula, the calibration of the pH meter, the dilutions, the temperature?). You could see if the error varies with starting concentration of the acid [HA] and also look at the effects of temperature. A comparison of several weak acids may also be revealing. How accurate is the pH meter for dilutions of strong acids? And then there are the bases. The possibilities are endless.

Water retention in disposable nappies
Today's state-of-the-art disposable nappy will absorb 15 times its weight in water. This phenomenal absorption capacity is due to the absorbent pad found in the core of the nappy. This pad is composed of two essential elements, a hydrophilic polymer and a fibrous material such as wood pulp. The polymer (eg polyacrylamide) is made of fine particles of an acrylic acid derivative, such as sodium acrylate, potassium acrylate, or an alkyl acrylate. These polymeric particles act as tiny sponges that retain many times their weight in water. An interesting EEI would be to measure the water absorption properties of the acrylate polymer using 'fake' urine (water, sodium chloride, urea, hydrochloric acid perhaps) in water in appropriate amounts. Does the nappy work equally well on individual solutions of the urine components (or are polar compounds different to non-polar ones)? How does temperature affect its properties? Where to get the polyacrylamide. You could rip open a nappy but a more controlled way would be to buy "water storage crystals" from the hardware shop. Get a Materials Safety Data Sheet (MSDS) for the one you buy to see what percentage of the crystals are polyacrylamide (should be over 90%).

One baby uses about 5000 nappies 500 g of Hortico crystals cost $14.81 at Bunnings. You can get 1 kg of Eden crystals for $20.99

Heat of reaction and E°
All chemical and biochemical reactions involve an energy change; e.g., chemical energy may be transferred as electrical, kinetic, light, sound, or (most often) to heat energy. Chemical to heat energy changes occur, for example, in displacement reactions such as: M(s) + Cu²⁺(aq) → M²⁺(aq) + Cu(s). Chemical to electrical energy changes occur, for example, in simple electrical cells; thus, a potential difference (V) is observed if the metal (M) is more or less reactive than copper. It would seem reasonable that the amount of heat evolved is directly related to the voltage of the cell. How true is this? Does it hold over a wide range of voltages, and is it concentration dependent? The photo (on the left) below may give you a start but how on earth will you measure the temperature change without heat loss? The photo on the right makes you glad you didn't have one of these on your lap.

Strength of plastic

If you’ve ever lifted a full plastic shopping bag you’ll know that some are stronger than others. Manufacturers make plastic objects with different strengths to suit different needs. It is not only thickness that is important, but the type of plastic, its density and amount of crosslinking. For example PE comes in several types: high-density polyethylene (HDPE), low-density polyethylene (LDPE), medium-density polyethylene (MDPE) and so on. You could imagine what UHDPE and VLDPE stand for. High-density polyethylene resin has a greater proportion of crystalline regions than low-density polyethylene. The size and size distribution of crystalline regions are determinants of the tensile strength of the end product. HDPE, with fewer branches than MDPE or LDPE, has a greater proportion of crystals, which results in greater density and greater strength. LDPE has a structure with both long and short molecular branches. With a lesser proportion of crystals than HDPE, it has greater flexibility but less strength. Why not make this an EEI? You could compare tensile strength with density and perhaps use temperature as a second IV. How do PP and PE compare if they have the same density? Hmmm! I’d have a look at the Australian Standard ASTM D 638 Test method for tensile strength of plastics. You don't need a fancy machine - you can do it at school.

Biodegradable Plastics
With the concern for plastics on the environment, many manufacturers now provide biodegradable plastics. For example, the plastic wrappers that many magazines and journals come in are biodegradable (see the photo below of the wrappers for Chemistry in Australia, and Australian Physics). It would be interesting to see how biodegradable these plastics are. Is it due to UV light, water, heat or what. I suspect UV light is a likely candidate but a perusal of a manufacturer's website should reveal the answer. Non-biodegradable bags are also available. You may have a UV light at school or else you can get a UV fluorescent bulb or tube from the hardware store for about $8. This would be an interesting EEI and the strength testing could be based on the suggestion above (or maybe you have a better idea). May take a few weeks to get a reasonable result so get started early.


Wrapper of Australian Physicist	Chemistry in Australia wrapper	Carry bag for frozen food.

Steam distillation of eucalyptus oil
Eucalyptus oil is used as component in pharmaceutical preparations to relieve the symptoms of influenza and colds, in products like cough sweets, lozenges, and inhalants. It has antibacterial effects on pathogenic bacteria in the respiratory tract. It used to be a big industry in Australia but has declined as cheaper imports have taken over. Nevertheless, eucalyptus oil, olive leaf oil and ti-tree oil are of vital importance to Australian industry - and society. One of the most disappointing laboratory experiments you can find is the steam distillation of these oils from leaves. They never work very well and you usually end up with a disappointing emulsion - not clear oil. For a good EEI you would need to do more than just extract some oil; you could have a go at improving the method by trialling different heating and collection methods, different aged leaves and so on; all carefully thought out and justified - not just trial-and-error. If you are stuck you could look at oranges or cloves. A trip to an olive leaf distillery would be a fun day out. The photos would look good in your report.

Soapmaking - the saponification of vegetable oil
A soap is the sodium or potassium salt of a long chain fatty acid. Soap making has been around for thousands of years and its manufacture is quite simple. However, there are many pitfalls because the chemistry involved is quite complex. A good EEI would be to make soaps from both sodium hydroxide and potassium hydroxide using a variety of saturated and unsaturated vegetable oils and to compare their properties with commercial soaps and detergents. Instructions for making soap can be found easily but you'd need to work out ways (and reasons) for changing the reactants and their quantities: that is, what problem are you trying to solve, and what is your hypothesis? To keep the investigation manageable, you would be wise to consider just two independent variables (perhaps type of hydroxide and saturation of the oil) and control the rest (salt, temperature, concentrations etc). The tests might involve suds formation in hard and soft water and ability to remove an oil spot. You could add some perfume and give the leftovers to mum for Mother's Day.

Breaking strain and crosslinking in polymers

Electrolysis of solutions
Electrolysis is commercially highly important in the separation of elements from naturally-occurring sources such as ores using an electrolytic cell. It involves the passage of an electric current through an ionic substance that is either molten or dissolved in a suitable solvent, resulting in chemical reactions at the electrodes and separation of materials. It is used in the production of metals such as aluminium, lithium, sodium, potassium and magnesium, and of non-metals such as chlorine. The electrolysis of water produces hydrogen and oxygen and that could make an interesting EEI. It is known that for water to be electrolysed, it has to have an ionic substance added such as sodium chloride. You could see how the efficiency of the electrolysis is affected by the voltage across the electrodes, and by the concentration of salt present. You'd need to relate your results to the E° value for the non-spontaneous reaction and what happens at voltages lower than that. You may look at the changing rate of generation of the gases as time passes or at the volume after a set time. It's up to you. Does the car ad below make sense? Is it chemically feasible?

Natural buffers

Our blood cannot tolerate a drastic shift in pH. It's a good thing, then, that human blood contains a buffer of carbonic acid, H₂CO₃, and sodium bicarbonate, NaHCO₃. This buffer regulates drastic shifts in the pH of our blood. If this buffer system was absent from our blood, the eating acidic or basic foods would cause the pH would swing too high (alkalosis) or too low (acidosis) and the result could be deadly. Another buffer system is that of a mixture of 0.1M Na₂HPO₄ and 0.1M NaH₂PO₄. As you add 0.1M NaOH or 0.1M HCl to the buffer solution and record its pH if will be noticeably different to that if you just used water instead of the buffer. Is the buffer any better if you used 0.5M solutions? What if you didn't have equal concentrations? Your EEI could be to locate a buffer system in nature (eg a lake) and test it out using natural environmental chemical changes (eg acid rain, increased greenhouse gases) and find if it has any limits.

Electroplating
Electroplating is a common industrial process. It is used to bestow some particular property on an object that it doesn't normally have, for example, abrasion and wear resistance, corrosion protection (galvanising, anodising), or aesthetic qualities (nickel or chrome plating). By applying an electric current, a layer of metal such as copper or nickel can be deposited onto a conductive object. In industry currents of about 500 A are common but in the laboratory a 12V power pack can suffice. A simple experiment that can form the basis of an EEI involves the use of a copper plate and a graphite rod as the cathode and anode, respectively. Nickel ion solution is used as the electrolyte. Under the influence of the battery, positively charged nickel ion can migrate to the cathode, pickup electrons and deposit on the surface of copper electrode; and there you have nickel plating. You could investigate the role of 'strike': initially, a special plating deposit called a "strike" may be used to form a very thin plating with high quality and good adherence to the substrate. This serves as a foundation for subsequent plating processes. A strike uses a high current density and a bath with a low ion concentration. The process is slow, so more efficient plating processes are used once the desired strike thickness is obtained. The striking method is also used in combination with the plating of different metals. Or you could investigate current density (amperage of the electroplating current divided by the surface area of the part) in this process strongly influences the deposition rate, plating adherence, and plating quality. The higher the current density, the faster the deposition rate will be, although you get poor adhesion. You may even produce some nice jewellery for mum.

Crosslinking in Slime
Slime is merely polyvinyl alcohol (PVA) that has been crosslinked by the addition of borax Na₂B₄O₇.10H₂O (sodium tetraborate). Various types of slime have been manufactured but the polymer polyvinyl alcohol is reasonably cheap and is readily available from suppliers because it is widely used as a thickener, stabiliser and binder in cosmetics, paper cloth, films, cements and mortars. Crosslinked PVA is used in hot or cold packs as they are not dangerous if the fluid leaks out. pH is critical in maintaining the crosslinks in slime. Too much acid will weaken the gel but this can be restored with the addition of alkali. A good EEI would be to test the resultant viscosity (you design the apparatus and procedure) as increasing amounts of borax is added (but you must hypothesise and theorise first); and/or to increase the [H⁺] by the addition of acid and then lowering it by the addition of NaOH. Another great EEI.

Electrorefining metals
Virtually all copper produced from ore receives an electrolytic treatment by electrorefining from impure anodes. In electrorefining, the anodes consist of unrefined impure metal, and as the current passes through the acidic electrolyte the anodes are corroded into the solution so that the electroplating process deposits refined pure metal onto the cathodes. In order to achieve high production rates, high current densities are desirable but an excessive current density causes at least two problems: increased impurity levels in the cathode deposit; and anode passivity occurs at current densities above 25-28 mA/cm². Hence, in industry, the current density is always low. An interesting experiment would be to set up an electrorefining cell for copper and find out the optimum current density and/or acid concentration.

Aspirin hydrolysis using a spectrometer
Aspirin is the common name for acetyl salicylic acid (ASA) and is an important drug on the market today. For example, the treatment of thromboembolism often requires the use of ASA. Aspirin is rapidly absorbed from aqueous solution and hydrolysis occurs during the absorption phase and first pass through the liver. It then is converted to salicylic acid (SA) in the blood, predominantly in the liver but also in blood cells, plasma, and kidneys. The hydrolysis of ASA to SA has been the subject of many investigations and can be studied in a high school laboratory if equipped with a visible spectrometer such as a Spectronic 20. The rate constant “k” for the reaction depends on pH, temperature, buffer concentration, and ionic strength. It can be followed by measuring spectrophotometrically the appearance of the complex of SA with ferric chloride, FeCl₃. The method can be found in The Journal of Chemical Education 2000, Vol 7, p 354 by L. Borer and E. Barry.

Conductivity of solutions
Electrical conductance is important in a variety of scientific contexts; e.g., nerve impulses, electroplating, electrical cells, and the extraction of metals by electrolytic reduction. You might expect the conductance of an aqueous ionic compound to be dependent on several independent variables, including the concentration of dissolved compound. A good EEI would be to examine this hypothesis: As the concentration (M) of sodium chloride increases, within the range ?? to ?? mol/L, the conductances (C) of the aqueous solutions increase in direct proportion. However, does the relationship hold for all concentrations, temperatures, electrode area, electrode separation and voltages?

Ion exchange resin
Ion-exchange resins are widely used in different separation, purification, and decontamination processes. The most common examples are water softening and water purification. Most recently, they can be used for biodiesel recovery. In many cases ion-exchange resins were introduced in such processes as a more flexible alternative to the use of natural or artificial zeolites. The resins are usually small plastic beads that contain ionic groups attached to a polymer-based resin. These ionic groups can be exchanged for similarly charged ions. There are many possibilities for an EEI here. Start with a cation exchange resin and plan an experiment to find the extent to which Na⁺ ions (from say NaCl solution) exchange with the hydrogen ions on the resin. You could Investigate the rate of exchange of ions by leaving the exchange resin in the sodium chloride solution for different periods of time (and plot graphs). Or you could investigate the effect of using different concentrations of sodium ions on the rate of exchange or the effect using cations such as potassium, calcium, aluminium, copper (II), and iron (II). Do they exchange to the same extent and at a similar rate?

Concrete hydration
The importance of concrete in modern society cannot be overestimated. Look around you and you will find concrete structures everywhere such as buildings, roads, bridges, and dams. There is no escaping the impact concrete makes on your everyday life. Concrete is prepared by mixing cement, water, and aggregate together to make a workable paste. It is molded or placed as desired, consolidated, and then left to harden. Adding gypsum, CaSO₄, to Portland cement prolongs the hardening. The most important compounds present in cement are: 3CaO•Al₂O₃, tricalcium aluminate; 3CaO•SiO₃, tricalcium silicate; 2CaO•SiO₃, dicalcium silicate; and CaO, calcium oxide. The 2CaO•SiO₃ reacts slowly with water to yield Ca(OH)₂ and H₂SiO₃. This reaction not only helps in holding the material together, but also makes the concrete less pervious to water. The hardening process is due in part to the hydration of the compounds present and is probably influenced by the crystallization of these hydrates. Concrete with too little water may be dry but is not fully reacted. The properties of such a concrete would be less than that of a wet concrete. You could make up thin slabs of concrete in a shallow trough with different amounts of water and test their breaking strain. What if you were unable to get fresh water - would seawater be just as good? If you try other additives, you have to say why you think they'd work (otherwise it's not chemistry - it's just backyard trial-and-error). The possibilities are endless.

Recycling an aluminium can
Recycling an aluminium can save 95% of the energy required to produce a new can from ore. Aluminum cans are easily recycled into new aluminum cans, but they can also be recycled into other useful aluminum products. A good EEI would be to convert aluminum cans into alum (potassium aluminum sulfate). Large amounts of alum is used by the paper industry as a filler in paper and secondly for drinking water purification. Merely converting can to alum is hardly the basis for an EEI. You’d need to apply some problem-solving and creative thinking. Perhaps you could look at ways of maximizing the yields and minimizing the input energy (heat) and chemical resources (KOH).

The health of our river
The Brisbane River and the waterways of the Moreton Bay catchment play a vital role in the economy, lifestyle and liveability of South-East Queensland. These waterways support the largest population of any catchment in the State and provide a nationally significant drinking supply. They also provide recreational and employment opportunities and are of cultural significance to the people of the region. But they are under enormous pressure from population growth. Scientific research indicates current levels of human impact on our waterways are unsustainable and our behaviour and practices must change if we are to halt and reverse the current decline in water quality. This context affords some great EEIs but you need to be careful that you don't just end up testing water samples and making some statements about water quality. If you plan to assess the health of the river you would need to state which tests you are using, why, the techniques, the sampling, the appropriateness of the tests if the water is saline, and so on. A good EEI would also be to ask “What is the effect of depth of water and temperature on dissolved oxygen as measured by using the Winkler technique?” or "What is the effect of salinity on chemical tests?" or “What is the more effective way of measuring salinity and what is the effect of tides on salinity?”.Many schools use this context for EEIs and its societal importance is obvious.

Strength of fired pottery clay

Pottery is one of the oldest human technologies and art-forms, and remains a major industry today. It is made by forming a clay body into objects of a required shape and heating them to high temperatures in a kiln to induce reactions that lead to permanent changes, including increasing their strength and hardening and setting their shape. Firing produces irreversible chemical changes in the body. As a rough guide, firing temperatures are in the range of about 1000 to 1400 °C. However, the way that ceramics mature in the kiln is influenced not only by the peak temperature achieved, but also by the duration of the period of firing. A good EEI (especially if you do Senior Art) might be to examine the hardness of the fired clay as a function of temperature; or as a function of time. If you were more adventurous you could look at different atmospheres within the kiln. One word of caution. This is a chemistry EEI and chemistry must be at its heart to distinguish it from applied technology or art. A pyrometric cone (see photo below) is a spike-shaped piece of clay used to measure temperature in a kiln when firing pottery. Cones have carefully calibrated melting points, indicated by their cone number. They are used to visually determine when a kiln has reached a desired temperature, by observing when a given cone in an observation port starts to droop. They are very attractive too.

· Polarisation of light in acidified sugar solution
Certain materials (sugar in this experiment) are optically active. When polarized light passes through an optically active material, its direction of polarization is rotated. The angle of rotation depends on the thickness of the material and the wavelength of the light. You could make up a solution of sugar (the disaccharide called sucrose) and hydrolyse it using dilute acid to form the monosaccharides glucose and fructose: C₁₂H₂₂O₁₁(sucrose) + H₂O + H⁺ => C₆H₁₂O₆(fructose) + C₆H₁₂O₆(glucose) + H⁺. The product is called invert sugar. As the reaction proceeds, the degree of polarisation changes and this can be observed using crossed polarisers either side of the solution placed on an OHP. Inverted syrups are sweeter than sucrose solutions and because there is glucose present in inverted sugar syrup it is substantially more hygroscopic (water retaining) than sucrose. This means that the syrup tends to keep products made with it moist for longer than when sucrose is used alone. It is likewise less prone to crystallisation and therefore valued especially by bakers. You could look at the effect of angle vs. concentration vs time; depth effects; acidity effects; temperature. If you are after a real challenge you could investigate if the reaction rate constant is dependant on acid concentration.

Carbon dioxide in soft drinks
Design an experiment to measure the solubility of CO₂ in lemonade as a function of temperature (by titration). Results obtained by this procedure are intended to indicate a trend in the solubility of the carbon dioxide as a function of temperature. In aqueous solution, carbon dioxide exists in many forms. First, it simply dissolves: CO₂(g) → CO₂(aq) Then, an equilibrium is established between the dissolved CO₂ and H₂CO₃, carbonic acid: CO₂(aq) + H₂O(l) ↔ H₂CO3(aq) The double-headed arrows mean that the reaction is reversible; that is, the products can react to form the reactants. The “species” are said to be in equilibrium if the rate of the forward reaction equals the rate of the reverse reaction. In this case there will be no change in their concentrations. Only about 1% of the dissolved CO₂ exists as H₂CO₃. Carbonic acid is a weak acid which dissociates in two steps. H₂CO₃ → H⁺ + HCO₃^¯ K_a1 = 4.2 x 10^-7 ; HCO₃^¯ → H⁺ + CO₃^2- K_a2 = 4.8 x 10^-11. When titrated, all the CO₂(aq) is reacted so a titration is a measure of total CO₂ content: H₂CO₃(aq) + 2NaOH(aq) → Na₂CO₃ + 2H₂O.


Carbon dioxide forms a weak acid in solution.	Titrate against 0.05 M sodium hydroxide	Use a phenolphthalein indicator	A faint pink color should persist for 30 seconds at the end point.

Acidification of seawater by carbon dioxide
Ocean acidification is a current and future problem for our ocean. The average coastal ocean pH is 7.9 in Queensland (Great Barrier Reef) Australia, but it is changing because of the addition of carbon dioxide to the atmosphere and the subsequent absorption of that carbon dioxide into the ocean. The pH of the Reef has decreased by 0.5 units in the last 60 years; it is becoming more acidic [Ref 1]. Some of the organisms at greatest risk include larva and shell-forming animals at the base of the food web that provide food for larger species. Organisms faced with the stress of ocean acidification can migrate, acclimate or go extinct. Additional stressors that increase the impact include temperature increase and habitat loss. The increase in ocean acidification is both an environmental and economic concern. This suggests a good EEI. Basically, you need to add chips of dry ice of different masses to water in a sealed container and measure the pH. Safety is of the utmost concern here as the dry ice can easily burn you. The big problem is how to control the controlled variables. You may also want to compare distilled water with salt water because the chemical components of the ocean acts as a buffer absorbing more carbon dioxide than freshwater can without a change in pH. Because the temperature of the ocean is also rising from global warming, the temperature variable would be worth considering.


Changes in pH at Arlington and Flinders Reefs, Australia, taken from "Evidence for ocean acidification in the Great Barrier Reef of Australia", Geochimica et Cosmochimica Acta, V73 (8), 15 April 2009, p2344, by Gangjian Wei, Malcolm McCulloch, Graham Mortimer, Wengfeng Deng, Luhua Xie.	Location of Arlington Reef on the Great Barrier Reef where pH was measured. Wouldn't it be great to be a student at a high school in Cairns?

Distillation of alcohol
Distillation is one of the oldest and still most common methods for both the purification and the identification of organic liquids. It is a physical process used to separate chemicals from a mixture by the difference in how easily they vaporize. Distillation relies on the fact that the vapor above a liquid mixture is richer in the more volatile component in the liquid, the composition being controlled by Raoult’s law. Not all mixtures of liquids obey Raoult’s law, such mixtures; called azeotropes, mimic the boiling behavior of pure liquids. These mixtures when present at specific concentrations usually distill at a constant boiling temperature and cannot be separated by distillation. Examples of such mixtures are 95% ethanol-5% water (bp 78.1 °C). I think you could make a successful EEI out of an experiment where you distill various ethanol/water combinations and measure the %ethanol in the distillate as a function of time or temperature of the vapour. You would need to consult vapour pressure diagrams.

Annealing
Metals are used for many different purposes. Two hundred years ago, the town blacksmith produced nails, hammers, wheel rims, knives, and horseshoes from the same basic metal. In some applications, a metal must be able to bend easily without breaking, whereas in other cases the metal must resist bending. Today metallurgists can produce these results by using different metals, alloying metals, and by heat treating metals. The substitution of a different metal or using a special alloy is often costly. Therefore heat treatment of a common metal is often the most cost efficient method of producing a metal that has the properties required in a specific application. Most metals respond to heat treatment, but the treatment temperatures are unique for different metals. A great EEI is to examine the effects of annealing, quenching, and tempering on metals. A steel bobby pin would be a useful starting point but you’d need to control the amount of heating and quenching and see how the properties vary with changes. Ask yourself “what type of treatment produces the hardest metal; and the strongest metal”? The male Blood Elf (below) from the World of Warcraft is carrying a quenching bucket. Nothing to do with chemistry however.

Anthocyanins in wine
Anthocyanin pigments are responsible for the attractive red to purple to blue colors of many fruits and vegetables including dark wine grapes. Interest in the anthocyanin content of foods has also intensified because of their possible health benefits. They may play a role in reduction of coronary heart disease, increased visual acuity, as well as antioxidant and anticancer properties. Anthocyanins are relatively unstable and often undergo degradative reactions during processing and storage. Measurement of total anthocyanin pigment content along with indices for the degradation of these pigments are very useful in assessing the color quality of these foods. There is a method used for determining anthocyanins in wine. It was developed by Fuleki and Francis and you’ll find it on the web. You'll also need a visible spectrophotometer (520 nm).

Corrosion of roofing steel
Corrugated galvanised iron is a building material composed of sheets of hot-dip galvanised mild steel, cold-rolled to produce a linear corrugated pattern in them. Galvanized metals prevent rust not only by protecting the metal from direct oxygen contact, but also by the electrochemistry of zinc. When iron rusts its oxidation state is increased as electrons are transferred away from the metal. Zinc acts as an electron donor in a slightly complex electrochemical reaction, thereby preventing the oxidation of the underlying metal. Nevertheless, rusting will be inevitable, especially if the local rainfall is at all acidic in nature. So for example, corrugated iron sheet roofing will start to degrade within a few years despite the protective action of the zinc coating. An EEI might be is to see how far the zinc protection extends over bare metal. Small scratches don’t rust but if the scratch is 10 mm wide will it? What if kept in a humid atmosphere? Is Lysaght zincalumeÒ better than ordinary galvanising?

Paper chromatography of leaves
Paper chromatography is an analytical chemistry technique for separating and identifying mixtures that are or can be coloured, especially pigments. This can also be used in secondary or primary colours in ink experiments. Most leaves are green due to chlorophyll. This substance is important in photosynthesis (the process by which plants make their food). You have probably done experiments where the different pigments present in a leaf are separated using paper chromatography. However, to make this a good EEI you need to take it further. Which is the optimum solvent (propanone, ethanol, hexane etc) and why (polar, non-polar, low viscosity, high BPt and so on)?

Diffusion of aqueous ions
Diffusion is the process by which molecules spread from areas of high concentration, to areas of low concentration. Diffusion occurs in liquids but more slowly than in gases because the particles are not as free to move about. When a crystal of Pb(NO₃)₂ and a crystal of KI are placed on opposites sides of a petri dish filled with water a yellow line of PbI₂ forms across the dish closer to one crystal than the other. This gives you an idea of the rate of diffusion of ions. You could repeat this with different combinations so long as they form a precipitate. Is it just the molar mass of the ion, or is it related to the charge, or perhaps something to do with electronegativity. Would anything happen with a non-polar solvent such as hexane, and if not, why not? A great EEI that will keep you entertained for weeks. There will be safety issues with heavy metal ions (eg Pb²⁺) so be warned.

Migration of ions
Ions, being charged, will migrate towards electrodes of opposite charge. For example, the migration of manganate ions (MnO₄^2-) can be observed if you cut a piece of filter paper slightly smaller than a microscope slide and moisten the filter paper with tap water. Then fasten the paper to the slide with crocodile clips and put a small crystal of potassium manganate (K₂MnO₄) in the centre of the paper. When you connect the clips to a power supply set at 12 V DC you should notice the migration of the coloured manganate ion towards the negative electrode. Occasionally permanganate (MnO₄^-) and manganate (MnO₄^2-) salts are confused, but they behave quite differently. How different is the speed of migration for larger ions, for ions of different charge. Is voltage related to migration speed. What a great EEI this is turning out to be.

Testing water hardness
Tap water in some parts of the country is very pure and is said to be ‘soft’. It easily makes a lather with soap. Water from other parts may contain various dissolved impurities and is described as ‘hard’ water. Temporary hardness may be removed by boiling, but permanent hardness survives the boiling process. You can measure water hardness by finding out the volume of a soap solution (of known concentration, eg 10 g of plain laundry soap per 100 mL of 80 % ethanol or metho) required to form a permanent lather with a known volume of the water to be tested (eg 5 mL) in a test tube. Actually, this is the standard Clarke's soap solution invented by Dr. Thomas Clarke, Professor of Chemistry at Aberdeen University, in 1843. A interesting EEI would be to see if the amount of soap needed is correlated with the concentration of various ions responsible for hardness (Ca, Mg). You could make up solutions with a range of concentrations of 'hardness' ions and see how much soap is needed to make a permanent lather (one that lasts for 30 seconds) is obtained when shaken. Try adding dropwise increments of the soap solution from a burette. 'Temporary' hard water can be made by using decanting a saturated solution of Ca(OH)₂; and permanent hard water can be made by using either 1 g CaSO₄.2H₂O or 1 g MgSO₄.7H₂O in 100 mL water. Permanent hard water contains Ca or Mg salts other than the hydrogen carbonates. Some tests you could do are: untreated deionized water (control), untreated tap water (real life); a comparison of untreated temporary hard water and untreated permanent hard water with boiled temporary hard water and boiled permanent hard water. You could investigate the effect of adding sodium carbonate crystals (washing soda) to temporary hard water; or the addition of adding sodium carbonate crystals (washing soda) to permanent hard water. Analysis of water hardness in major Australian cities by the Australian Water Association shows a range from very soft (Melbourne) to very hard (Adelaide). Total Hardness levels reported in various government reports are listed below:

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Cut a marble tile with a masonry blade into several long strips say 10 cm x 1 cm x 1 cm	Make measurements with a Vernier calliper


Break the strips into 2, 4, 6 or 8 pieces. I used a cold chisel to cut this one into 8. Measure all the pieces.	Let them react with acid.


After 1 minute the two solutions have different Cu²⁺ concentrations shown by the different intensities of blue colour. This means more Zn has reacted in one than the other.	You can see the brown copper metal that has been displaced out of solution by the zinc. You don't have to use zinc - any more reactive metal than copper would work. You'll need one whose ions are colourless. There are lots of combinations you could try (eg Mg + Cu²⁺).


For copper ions, set the wavelength to 740 nm. Different metal ions have different values for lambda max.	The filtered Cu²⁺ solutions can be analysed in a spectrophotometer. My thanks to Moreton Bay Boys College for access to their instrument.


One baby uses about 5000 nappies	500 g of Hortico crystals cost $14.81 at Bunnings.	You can get 1 kg of Eden crystals for $20.99