Please note: the suggestions below are just ideas for student experiments; they have been trialled but do not guarantee success. It is up to the teacher and student to see if they are practical for their school situation.
Electrochemical cells
The suggestions that follow are possible modifications of the mandatory practical: construct a galvanic cell using two metal/metal-ion half cells.
When investigating electrochemical cells there are two approaches that can be taken:1. Open circuit. This is where the cell is not being discharged. This is by far the most popular type for students to investigate as it is much easier to control the variables. In this situation, the only external device across the anode and cathode is a voltmeter or digital multimeter (DMM) used to meassure the cell potential (the voltage). The current through the external circuit will be extremely low as the electrical resistance of the meter is so high. Numerically, the resistance of the meter is typically 10 megohm (10 million ohms), so for a cell voltage of 1.0 V, the current is equal to 1/10,000,000 = 1 x 10-7 A (0.1 microamp). This means that only a small amount of the cell's chemical species are reacting.3 vs NH4 NO3 ), submerged surface area of the electrodes or salt bridge. As I said, these may have an impact with cells under load, but not with an open circuit cell. Don't bother with them unless your experimental design is based on a null result. This can be useful but is unlikely to provide data suitable for analysis and discussion at a high level.
2. Cell under load. This is where the cell is being discharged, usually by some external 'load' device placed across (between) the anode and cathode. This will be a device that allows a moderate current to flow through the external circuit. Examples are light bulbs (torch bulbs), or a fixed resistor (eg 100 ohm). In my experience, resistors are better than bulbs. For a 100 ohm resistor across a 1.0 V cell, a current of 0.01 A (10 mA) will flow. This means an appreciable chemical reaction is occuring.
Let's look at some possibilities for open circuit cells, and cells under load.
OPEN CIRCUIT CELLS
You are bound to have made a Daniell Cell while doing classwork on electrochemical cells. It is written Zn(s)/Zn2+ (aq) // Cu(s)/Cu2+ (aq) which shows that it is made up a zinc half cell (anode) and a copper half cell (cathode), joined by a salt bridge. The standard electrode potential is 1.10 V (measured at 25°C). The spontaneous reaction is: Zn(s) + Cu2+ (aq) → Zn2+ (aq) + Cu(s).
HINT 1. The relationship between voltage and temperature is given by the Nernst Equation Ecell = Eo cell – RT/nF x ln Q, where R is the universal gas constant and is equal to 8.31 J K-1 mol-1 , T is temperature in Kelvin, n is the number of electrons transferred for every ion (n = 2), F is Faraday's constant = 96485 C mol-1 , and Q is the activity constant = [Zn2+ ]/[Cu2+ ] for a Daniell Cell. The term ln Q stands for the natural log of Q, or loge Q. For equimolar solutions of zinc and copper, that is [Zn2+ ] = [Cu2+ ] the value of Q = [Zn2+ ]/[Cu2+ ] is 1) and because the value of ln(1) is zero, that whole term in the equation (RT/nF x ln Q) reduces to zero so there will be no effect of a temperature change.
NOTE (to Queensland students). The Nernst Equation is not listed as subject matter in the Queensland Chemistry syllabus
HINT 2. You will need to make the concentrations different (eg 0.3 M and 0.01 M). Depending on which one you make bigger, you will get an increase in voltage with temperature (when [Zn2+ ]<[Cu2+ ]), or a decrease in voltage with temperature (when [Zn2+ ]>[Cu2+ ]). Just check the equation.
HINT 3. The bigger you make the difference in the concentrations of Zn2+ and Cu2+ , the bigger will be the second term of the Nernst Equation and so the bigger will be the change in voltage with temperature change. Here's an example to show what I mean. If the Q ratio is small, eg 2: ([Zn2+ = 2.0M] and [Cu2+ = 1.0M], lnQ is 0.693 so a change in temperature has almost no measurable effect (at 278K (6 o C E = 1.0916 V, and at 333K (60o C, E =1.0900 V). That’s for the biggest temp range safely possible in the lab. A digital multimeter (DMM) will struggle to show a difference especially when the scale reading uncertainty for a DMM would be +/- 0.001. I find that keeping [Zn2+ ] large and making [Cu2+ ] small, I get more consistent results.
HINT 4. Be careful with solubilities of zinc compounds. Anhydrous zinc sulfate, ZnSO4 , has a solubility of about 58 g/100 mL at 20o C or about 3.5 mol/L. However, zinc sulfate heptahydrate ZnSO4 .7H2 O has a solubility of only 54 g/100 mL which is about 1.87 M. So be certain about which type you have. The label on the main container will state the molar mass. If it says MW 161.47 then its the anhydrous form. If is says MW 287.54 then its the heptahydrate form. Zinc nitrate on the other hand is much more soluble. It is usually supplied to schools as the hexahydrate, Zn(NO3 )2 .6H2 0, which has a molar mass of 297.5 g/mol, and has a solubility of at 184.3 g/100 mL (6 mol/L). When developing your method check the label for the exact formula and molar mass. The anhydrous form of zinc nitrate, Zn(NO3 )2 , is not readily available and it has a molar mass of 189.4 g/mol.
HINT 5. I suggest you keep the Cu2+ concentration as low as possible, but not too low as the conductivity depends on the concentration of ions. I've found that a [Cu2+ ] of 0.001 M or 0.0001 M works well. When [Zn2+ ] = 3.0 M, and [Cu2+ ] = 0.0001 M, the activity constant Q = [Zn2+ ]/[Cu2+ ] = 3.0/0.0001 = 30,000. This gives an ln Q = 10.3 so the impact of temperature on voltage should be evident when using school voltmeters or multimeters.
HINT 6. Keep the volumes of electrolytes in each half-cell as small as possible. You should not need to go past about 100 mL. Some schools do it on a microscale with less than 1 mL per half-cell.
HINT 7. When making up the copper solution from copper nitrate or copper sulfate, make the pH of the final solution in the range 3-4. At pH >4 the copper ions react with OH- ions and dissolved oxygen in the water to form the insoluble (precipitate) copper hydroxide Cu(OH2 . This coating on the surface of the copper electrode interferes with the functioning of the cell and you may end up with voltages lower than you expect.
HINT 8. If you construct a Daniell Cell and heat it up the voltage and temperature can be plotted to show their relationship (with the independent variable T (K) on the x-axis). You could plot T values in o C if you like. It should be linear with a negative slope (gradient) if you have construced the cell with [Zn2+ ]>[Cu2+ ]. Draw a line of best fit and state the equation for the line.
HINT 9. You could have Excel calculate the R2 (R squared) value. This is a measure of the 'goodness of fit' of the trendline to the data points. That is, a measure of the precision or spread of results around the trendline. A value of R2 > 0.98 would be considered excellent. It means that the formula has captured the data well. Another way of saying it is that the equation and trendline accurately represent 98% of the data points.
HINT 10. The equation could be rewritten as Ecell = Eo cell – kT where k is the constant (R/nF). This can be rearranged as Ecell = –kT + Eo cell which is now in the form y = mx + c. You should be able to take it from there. Here is a picture of my setup (I used 0.3 M Zn2+ and 0.01 M Cu2+ solutions of the electrolytes and a saturated KNO3 saltbridge):
The open cell voltage at 28°C was 959mV
At 50°C the voltage had dropped to 950mV.
When I plotted all the values I got a linear graph (well, almost) with a intercept of 1.059V and a slope of -0.0023 V/°C.
This is a great investigation for the errors come thick and fast. When I let it cool back down the voltage continued to drop. Now that is worth investigating. Perhaps the electrolyte concentrations were changing as the reaction took place (most likely), or perhaps the salt bridge was drying out from the heat. And I know how to fix that.
Here's the theoretical graph based on the Nernst equation using [Zn2+ ] = 0.3 M, [Cu2+ ] = 0.01 M. The intercept on the y-axis should be the Eo cell.
HINT 12.
Electrochemical cell - effect of temperature on a zinc/ferricyanide cell
Zn(s) | Zn2+ (aq, 0.001M) || Fe(CN)6 3- (aq, 0.005M), Fe(CN)6 4- (aq, 0.005M) | Pt
Instead of the platinum electrode, for high school you can use a carbon rod. The differences are small and not worth worrying about.
Here are my results. It was pretty rough and ready but it shows what to expect.
Set it up and measure the EMF over the range from say 0°C to 60°C. Be warned: potassium ferricyanide Fe(CN)6 3- and potassium ferrocyanide Fe(CN)6 4- are poisonous; use gloves when weighing them out. Make sure you do a good risk assessment and check with your teacher before starting. You may be refused permission.
2. Electrochemical cell - effect of concentration changes on a Daniell Cell
The relationship between voltage and concentration is given by the Nernst Equation Ecell = Eo cell - RT/nF x ln Q. The value of Q (activity constant) is equal to [products]/[reactants] (= [Zn2+ ]/[Cu2+ ]) . You will need to make the concentrations different (eg 0.3 M and 0.01M). As the concentration of the zinc ion is increased the lnQ gets bigger so the Ecell gets smaller. That would make a neat experiment. The solubility of zinc sulfate is about 60 g/L whereas a 1 M solution would require 160 g/L and that's why I suggested a maximum concentration of about 0.3 M. The graphs below are only theoretical. I'd say you would have fairly big error bars on the grap (if you use error bars). I've made a You Tube video on adding custom error bars to a graph so long as you have at least two data points (voltages) for each concentration. See Adding custom error bars to a graph https://youtu.be/mVS7YWENShs
We can linearise the graph by plotting cell voltage vs ln[Zn2+ ]. I've plotted the negative of the log to make the graph have positive values.
Remember, the graphs above are only theoretical and yours will be different. Explaining the difference is where you can get some marks.
Electrochemical cells - effect of the salt bridge
NO LOAD: By "open circuit" I mean without a load resistance across the metal
electrodes. I read on one physics help-page that if you add more salt bridge strips the voltage should creep up a bit as the internal resistance of the cell would be lowered. That would make sense if a current was flowing through a load resistor - but if it was open-circuit and the only load was that of the digital multimeter (very, very high and thus a tiny current) the internal resistance wouldn't matter.
So I tried it and as I added more salt-bridge strips (saturated KNO3 ) and the voltage actually went down. I got a beautiful negative linear relationship when I plotted it (V = -0.0154N + 1.034) with N being the number of strips and R2 = 0.9918. I could scarcely believe it. It implies that with no salt bridge the voltage would be 1.034V which is close to the open-circuit voltage. This would be a fabulous investigation if you were prepared to do some research. Send me your results.
With just the one salt-bridge strip the open-circuit voltage was 1.02V
With five salt-bridge strips the voltage was 0.96V
UNDER LOAD: What I did next was to try the same experiment but with the cell under load. That way the salt bridge would be the route for the internal migration of ions. Look at it this way: a cell has an internal ohmic-equivalent resistance. The salt bridge is a resistive path for the flow of ions. Ions, plus their hydration shells and possible coordination shells, have low drift velocities through a medium. The salt bridge can constrain the ion current of the cell and its material equilibrium rate. A poor salt bridge by length, diameter, concentration, and permeability will increasingly adversely affect a cell's current and voltage as discharge rate increases. Under a load resistor, the voltages for a single salt bridge was 0.70 V and for 5 bridges it was 0.92 V. So that seemed to work.
The total resistance of the circuit RT is equal to the sum of the resistances of the voltmeter itself (RV ), the solution (RS ), and the salt-bridge (RB ); that is RT = RV + RS + RB
We can omit RS as it is so small compared to the other two. Seeing the current is the same through each resistive load, we can multiply each term by the current (I) and the equation becomes VT = VV + VB or VV = -VB + VT . We also know that the resistance of 2 salt bridges is half that of 1 salt bridge, so the resistance of "n" salt bridges is RB x 1/n.
We thus end up with an equation: VV = -VB x 1/n + VT . This is in the form y = mx + c, so if you plot VV on the y-axis against 1/n on the x-axis, the slope should be VT .
Electrochemical cells - effect of changing the load resistance As mentioned above, as the cell's current increases, the salt bridge will have increasing trouble providing a route for the internal migration of ions. Another EEI worth trying is to keep everything constant (including the salt bridge) and change the value of the discharge resistor. A 3.3Ω resistor will allow a bigger current than a 100Ω one. There should be a proportional change in volatge as the resistance is decreased but it should start to go funny once the salt-bridge can't keep up. I'm going to try this.
Electrochemical cell and polarisation
However, as the cell discharges the concentration of ions at the electrodes changes. For example, in a Daniell Cell, the [Zn2+ ] increases at the anode, and [Cu2+ ] decreases at the cathode. This is called concentration polarization . It is observed when diffusion, migration and convection are insufficient to transport the reactant to or from electrode surface at its initial rate. If the cells are agitated (by a magnetic stirrer for instance) the polarization should be reduced and the voltage be closer to the initial voltage. Heat may do the same thing by increasing the convection currents, but heat also has other effects that will complicate matters (see above).
A 3.3Ω load resistor has been connected across the electrodes, and the voltmenter put in parallel to it to monitor its changing potential (voltage). At the start (0 s elapsed as shown on the stopwatch) the voltage was 4.1 mV and the current 0.17 mA
After 6 minutes, the voltage was down to 2.9 mV and the current 0.13 mA. When you calculate the power output of each cell (P=VI) it has gone from an initial amount of 0.7 µW to 0.38 µW.
Concentration cells
A concentration cell is an electrochemical cell made up of two half-cells with the same electrodes, but differing in electrolyte concentrations. As the cell as a whole strives to reach equilibrium, the more concentrated half cell is diluted and the low concentration half cell has its concentration increased via the internal transfer of ions, and the external transfer of electrons. Therefore, as the cell moves towards chemical equilibrium, a potential difference is created.
An example for this type of cell is where the cathode consists of Cu2+ (1.0 M)/Cu and the anode consists of Cu/Cu2+ (0.1 M). In this cell, the flow of electrons from the anode to the cathode is due to the oxidation of Cu metal at the anode and the reduction of Cu2+ ions at the cathode into metallic copper.
Electrolysis and temperature
Setup used at Our Lady's College, Brisbane. The cell was just a part of an conductivity kit (with the bulb removed and bypassed).
You can see the bubbles of hydrogen on the cathode. The electrolyte was 0.1M KOH and we got 0.4 V and 27 mA at 28°C. At 55°C the current was 40 mA. The hand belongs to Yr 11 Chemistry student - Emma - at Our Lady's College.
Heating the water to a higher temperature leads to a smaller voltage being required to achieve the same rate of hydrogen production (as measured by the electrical current). Alternatively, at higher temperatures, if the voltage is held constant then the current will increase and so will the rate of hydrogen output. So there is the basis for a good EEI. Make sure you pay attention to the controlled variables: electrode type and area, concentration of electrolyte, and so on. Because polarisation can affect the variables, it may be best to take your measurements quickly after the start. Note: rather than use an electrolytic cell, a similar investigation can be undertaken using electrochemical cells (as above). In that case the Nernst Equation applies and you can easily see how the absolute temperature is factored in.
Tinny taste of fruit in tin cans
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 H2 O (slow)
NO2 - + 6e- +8H+ → NH4 + + 2H2 O
(fast)
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 + 3H2 O 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 H2 O. 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 [*Vogel, 1961, p374] 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 [*Vogel, 1961, p 800]. 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 my MSc dissertation can be downloaded by clicking here .
*Vogel, AI 1966. A Textbook of Quantitative Inorganic Analysis, Longmans, London.
Cross section of a
tin-coated steel can cut through with a pair of scissors. Don't use Mum's "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).
Detinning of tin-plated steel cans
Tin is a moderately corrosion resistant material that is widely used in tinplate for food beverage. Tinplate is light gauge, low carbon steel, coated with commercially pure tin. Worldwide, more than 90 billion food packaging cans are used per year; therefore this application is by far the largest of its diverse industrial applications. Dissolution of metallic tin from a steel can can be either undesirable or desirable; and this suggests a great EEI whereby you could look at the factors affecting the rate of detinning of steel cans by tartaric acid solution.
Tartaric acid has a pKa of 2.95 for its 1st ionization, making it one of the strongest of the 'weak' acids.
Chopped up and baled steel cans ready for detinning.
Undesirable: Dissolution of metallic tin from the inside of a can into the food content has a major influence on the food quality and may cause toxicological effects.
This can be caused by two major food acids: citric and tartaric acid. In an interesting paper, Professors Rabab M. El-Sherif and Waheed A. Badawy from Cairo University, Egypt, found that the presence of certain amino acids in canned fruit inhibited the dissolution of tin by tartaric acid. Their paper "Mechanism of Corrosion and Corrosion Inhibition of Tin in Aqueous Solutions Containing Tartaric Acid" was published in the International Journal of Electrochemical Science , 6 (2011), p 6469-6482. They found that the amino acids alanine, glycine, glutamic acid and histidine were inhibitors for the tin dissolution process. The bonus is that these amino acids are environmentally safe and could be added to food.
You could see how the presence of different concentrations of (just one) of these amino acids inhibits detinning for a given concentration of tartaric acid (eg 1.00M tartaric) at a constant temperature over a given time. To get you started it may be helpful to know that: 0.001M glycine concentration in the tartaric acid gave a corrosion efficiency of 19% (that is 19% less dissolution than without any glycine) whereas 1.00M improved this to 67% efficiency (after 2 hours at 25°C) - with a dramatic increase in efficiency from 0.001M to 0.01M. Secondly, when the scientists compared the corrosion efficiency of different amino acids (at 1.00M for 2 hours at 25°C) they found that glycine was the best at 67%, followed by alanine at 38% and histidine at 35%.
Desirable: Recycling tin-coated steel cans to recover and separate the tin and steel is industrially important particularly when the price of tin is high.
Today most detinning is done by an electrolytic process whereby the steel cans are soaked in a hot caustic solution and the tin electrolytically deposited onto cathode or anode plates. An earlier method has some interesting chemistry. Research done in 1930 by Scott and Davis at UCLA (California) found that tartaric acid could be used to de-tin steel cans without dissolving any of the iron. The tin could then be recovered from the stannous tartarate solution by precipitation. Their paper was published in Industrial and Engineering Chemistry , 1930, V22 (8), p 910-911. Scott and Davis found that 3 hours of immersion in a 5% tartaric acid solution with constant aeration gave optimum detinning. You could see how the concentration of tartaric acid is related to the amount of tin removed. Question: how will you aerate the solution (does an aquarium pump give you an idea)?
HOW TO MEASURE TIN CONCENTRATIONS?
See the note in the EEI suggestion above about colorimetric analysis of tin.
Fermentation and the alcoholic content of wine - analysis by titration
The quality of Queensland wines is now recognised as amongst the best in Australia. Overseas exports are increasing, particularly to international markets seeking premium quality boutique wines. The Queensland wine industry has grown significantly over the years to cover a total of 1400 hectares. The majority of this growth has occurred during the past 20 years with significant plantings throughout the southeast corner of the State. However, winemaking is still somewhat of an art but is strongly informed by science. Thus an interesting student experiment can be undertaken in this context.
Sirromet Winery Mt Cotton: 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 Winery, Mt Cotton, Redland City.
THE MOST IMPORTANT INDEPENDENT VARIABLES IN FERMENTATION
There are two key independent variables worth considering:
(a) Sugar concentration. After crushing the grapes the next step in the making of wine is the fermentation of the grape juice and pulp with various yeasts and bacteria. Most books say that the amount of ethanol produced is dependent on the sugar concentration of the starting juice but then give four different equations depending on assumptions made (such as the ratio and purity of glucose and fructose; or whether the fermentation gets 'stuck' at the primary stage). The most common relationship is a linear one (y = 0.6x -1) - so an EEI could investigate that.
A good student experiment would be to use fresh grape juice or simulate grape juice with 150-250 g/L glucose (or an equal mix of glucose and fructose), adding a controlled amount of yeast and wine acids and fermenting to stillness at constant temperature.
(b) Temperature. Fermentation is an exothermic reaction so heat is generated during the process. To control the heat, the winemaker must choose a suitable vessel size or else use a cooling device. In the case of a student experiment , you might control the temperature by use of a water bath (or a refrigerator). Typically, white wine is fermented between 18-20°C though a wine maker may choose to use a higher temperature to bring out some of the complexity of the wine. Red wine is typically fermented at higher temperatures up to 29°C. Fermentation at higher temperatures may have adverse effect on the wine in stunning the yeast to inactivity and even "boiling off" some of the flavors of the wines. How temperature and the final concentration of alcohol are related would make an ideal student experiment.
OTHER INDEPENDENT VARIABLES WORTH CONSIDERING
(c) Acidity. This is not such an important one and the effects may be small - but nevertheless important. You could repeat it with acidity as the independent variable and controlling the amount of sugar, yeast, temperature and so on.
(d) Type of yeast. As the alcohol concentration rises the yeast cell membranes become susceptible to rupture by the ethanol. Some yeasts are more susceptible than others. Baker's yeast is very susceptible and will die at just a few % alcohol; brewer's yeasts (for beer) are okay up to 5% but some can survive in up to 9% alcohol; and wine yeast usually go from about 13% (Sav Blanc), Riesling (16%) and a sherry yeast can tolerate about 17%. Or you could look at the susceptibility of yeasts to [SO2 ] - winemakers use SO2 in the form of sodium metabisulfite to kill off wild yeasts as these are less tolerant than wine yeast to the SO2 . The interesting thing is that you could breed a yeast to survive higher alcohol concentrations (like they do at wineries for their malo-lactic fermentation of sparkling wines) by slowly increasing the amount of alcohol in the brew from say a few % to 10% over a period of a week so that the yeast adapts. This sounds more like a Biology EEI so I'd better stop here.
(e) Sulfur dioxide. Sulfur dioxide is widely used in winemaking because of its antioxidant and antibacterial properties. You could hypothesise and test how SO2 affects the performance of yeast. Two common methods for determining SO2 in wine are included in the "sulfite" section later.
SCHOOL WINEMAKING IN GENERAL
Any fruit (or juice) works just fine although some require more sugar to be added. James Palframan HOD Science at MacGregor High School Brisbane adds a salutary note: "My year 12 chemists last year made a range of wines including lychee, lemon, ginger, dragon fruit, apple and mango. The dragon fruit wine ended up being very expensive at approximately $10 to $12 a bottle as the dragon fruit was quite expensive to start with and then the students discovered when that chopped it up that it was mostly water and not a lot of fruit pulp". Two successful used Wine EEI task sheets are from Otto Craig Wine EEI #1" and Melissa Dixon Wine EEI #2 (but remember, these are not exemplars , they are just from submissions that have been to QSA District or State Panels for the purpose of review). Also, my winemaking unit for chemistry teachers is available online.
It is worth stressing here that you should not taste the wine being produced; this is not because of the alcohol but rather because of the unsanitary conditions under which your wine is being made (in a lab, not the Home Economics kitchen). If you intend tasting your wine then your risk assessment should state and evaluate that. If I was your teacher I'd say "no" as it would not be worth the headache.
THE DEPENDENT VARIABLE
The most common method is by redox titration. In this analysis, you add an excess of standardized acidified potassium dichromate solution to the wine which converts the ethanol to ethanoic (acetic) acid. The amount of unreacted dichromate is then determined by adding an excess of potassium iodide solution which is also oxidised by the potassium dichromate to form iodine. The iodine is then back-titrated with a standard solution of sodium thiosulfate and a starch indicator. The titration results are used to calculate the ethanol content of the original solution. It is complex but works well and is very impressive.
You have a problem if you are dealing with red wine as the red pigments mask the colour changes. In that case you have to extract the ethanol from the wine (in effect, by various forms of distillation) and carry on, as above, from there. Canterbury University NZ has quite a simple method for red wine: see Canterbury - alcohol titration . Chemistry teacher Emma Hodginkson from Mountain Creek State High has performed the Canterbury ethanol titration with her Year 12s for a few years and has found it very successful. When analysing commercial wine, her students get very close to the %alcohol on the label. As it is a redox titration she says it works better in Year 12 when the students have completed some redox theory whereas Year 11s find the calculations a bit heavy going. To overcome the practical difficulty in locating a small container to suspend above the dichromate - they use a plastic water bottle lid suspended with cotton thread. Cool!
At All Hallows' School, Brisbane, chemistry teacher Matthew Stuart uses a more compact method: a boiling tube is used to hold the 20mL wine sample, and 8mL dichromate solution is placed in a small "fusion tube" (a small diameter test tube) carefully slipped into the wine sample. Some parafilm is used to seal the boiling tube, or a suitably sized stopper if possible. With care, the small inner test tube will float on top of the wine sample, and these are left in a drying oven (50°C) overnight to drive (evaporate) the ethanol to the dichromate solution. A pair of forceps is used to remove the dichromate tube without spilling into the wine sample, the outside rinsed with water, and then titrated. 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.
Yeast type and alcohol concentration in wine
You may have seen different types of yeast for different purposes; for example, there are beer yeasts (top-cropping ale strains and bottom-cropping lager strains); baker's yeast for bread-making; red and white wine yeasts; and even genetically engineered yeasts used for industrial alcohol production. Many are just different strains of the same yeast but grow differently. A good research question might be: how do different yeasts affect the production of alcohol from grape juice?
Baker's Yeast - just $4.59 for 12 sachets from Woolies. The packet says "yeast 98%, emulsifier (sorbitan monostearate) and ascorbic acid". The Lowans yeast to the right gives a faster rise I've found. Both types are "Active Dry Yeast" and consist of coarse oblong granules of yeast, with live yeast cells encapsulated in a thick jacket of dry, dead cells with some growth medium. You used to be able to buy 'compressed yeast' which is essentially a yeast suspension with most of the liquid removed. It is a soft solid, beige in color in foil-wrapped cubes.
Winemaker's yeasts (Saccharomyces cerevisiae ) come in different strains and are selected depending on the type of wine grape varietals being used. White wine yeasts are optimised for growth at 20°C and allow for the development of fruity esters. Red wine yeasts work best at 29°C. You can buy these from a 'Home Brew' shop or on-line for about $3 for 25 g. Today there are several hundred different strains of S. cerevisiae dentified but not all are good for wine making. Some give a 'stinky fermentation' (as they say). Just your luck!
There is a problem however. The majority of senior chemistry investigations use 'continuous' variables; that is, any value is possible within the limits the variable ranges. For example, temperature, pH, sugar concentration, amount (mass) of yeast, surface area and so on. But when you are comparing two different categories of a variable, such as type of yeast (wine yeast vs. baker's yeast) then you have a categorical variable. The choice of the categorical variables is not as common as using a continuous variable in Senior Chemistry experiments, partly because categorical variables have their peculiar difficulties and this makes the design of this EEI far more complex than it looks at first glance. It is not just a matter of comparing equal amounts of the two yeast products on the amount of alcohol produced.
A decision has to be made about the amount of each product to use to get some sort of equivalent mass of yeast for comparison (and how this is arrived at; is there any indication of the % composition of the two products). Do the yeasts each have an optimum pH and if so what pH will be chosen for the grape juice (and why)? I know that the Lalvin BGY yeast from Burgundy, France is hopeless at pHs lower than 3.2 but other work at higher pHs. Is surface area a concern (maybe if one is a bottom fermenter, and another a top fermenting yeast). What temperature will be used (and why) if the yeasts have their own optimum temperature for growth; for example the BGY Lalvin yeast from Burgundy, France works best at 24°-28°C. Will a low sugar or high sugar juice be used - important as it may be the alcohol itself that inhibits the yeast. For example, the Lalvin CLOS yeast from Spain is high-alcohol-tolerant up to 15% alcohol but others give up before that. And what about the dependent variable (alcohol concentration): will the rate of alcohol production be measured, or just the amount of alcohol present when the yeasts die or the sugar runs out; or will the alcohol be measured after a set time, eg 7 days? Is time important? Some yeasts are slow (eg the CY3079 Slow White yeast from France takes its time but gets there in the end; it would be a brave decision to cut it off after 7 days). Lastly, some yeasts convert malic acid to alcohol (as well as converting the sugar). Imagine using a yeast such as the Lalvin C from France which partially degrades malic acid. Of course you'd get more alcohol out of this one.
It could also be taken further by using a continuous variable (amount of yeast product) in conjunction with the categorical variable. The possibilities are quite endless really. What a fabulous investigation - if you have the time.
Which ferments best: glucose, fructose or sucrose?
Another terrific idea for a wine student experiment. The two common fruit sugars used in winemaking are glucose and fructose. Grape juice is made up of these in roughly equal quantities. Another sugar used in the fermentation industry is sucrose (cane sugar). Sucrose is frequently used as a cheap carbohydrate by breweries, wineries and other fermentation-based industries employing yeast. It is a di-saccharide composed of D-glucose and D-fructose linked by an alpha-1,4 glycoside bond. In the initial stages of fermentation, sucrose is rapidly hydrolyzed into glucose and fructose by the action of the enzyme invertase on this bond. Then the sugars are transported across the cell membrane where they ferment and form alcohol. So investigating the fermentation of sucrose is really also about studying the fermentation of a mixture of glucose and fructose. This suggests a terrific student experiment: to look at the reaction rates of glucose, fructose and sucrose. I am indebted to Savannah from Coolum State High School for suggesting this to me.
Savannah's photos: Gases from the three fermenters was captured and measured using water displacement in measuring cylinders. As Savannah found - CO2 is water soluble and will be the source of errors that have to be allowed for.
Glucose fermenter. May 2014.
Measuring sugar concentration with a Brix hydrometer.
Saccharomyces cerevisiae consumes glucose faster than fructose because this yeast is know to prefer glucose (and so is called a glucophilic or 'glucose-loving' yeast). As fermentation proceeds the ratio of fructose to glucose increases. Therefore, fermented grape juice will contain more fructose than glucose as residual sugar. Fructose is the sweetest hexose sugar, approximately twice as sweet as glucose, and thus the wine gets an undesirable fructose sweetness, unbalanced by the sweetness of glucose.
Source: Rodicio, R. Heinisch, J.J. (2009). Sugar metabolism by Saccharomyces and non-Saccharomyces yeasts, chap. 6, pp. 113-134. Springer-Verlag Berlin Heidelberg.
Savannah's results from her Coolum State High Year 12 Chemistry EEI.
So here is the starting point if you are doing a student experiment. Pose the Research Question: "How do the fermentation rates of glucose, fructose and sucrose compare?" Potentially it is a fantastic student experiment. Studies have found that glucose and fructose ferment at equal rates when they ferment separately; and when their concentrations are above 1%. However, when the glucose is below 1% it reacts faster than fructose. One delicious complication though is that when glucose and fructose are mixed (as in the case of fermenting sucrose; or in an artificial mixture) the glucose ferments faster than fructose. Glucose seems to inhibit fructose. Hmm - now that is tricky.
Your investigation could look at the fermentation rates of glucose and fructose separately - keeping everything the same except the independent variable of initial concentration. Note that 1% seems to be some sort of cut-off so examining concentrations either side of this would seem laudable. An then look at a mixture - or just look at sucrose as a natural mixture. What are you going to measure? The progress of the fermentation can be assessed by measuring the concentration of residual sugar or of the ethanol, or by the amount of CO2 produced. Sugar concentration can be measured using a Brix refractometer; or in the case of the two reducing sugars glucose and fructose, you can do a Fehling's titration; and there is a titration for ethanol. The density can also be used as an index. You could also monitor the reaction with a gas pressure sensor. I've used the Vernier sensor with a Texas Instruments CBL and that works well. There are lots of other gas sensors too.
A good start into the experimental glucose/fructose/sucrose fermentation is Leanie Mocke's 2013 Master of Science (Biochemistry) Thesis from Stellenbosch University, South Africa entitled Kinetic Modelling of Wine Fermentations: Why Does Yeast Prefer Glucose to Fructose? [available on-line].
Sulfites in wine - titrations
Sulfite/sulphite is a general term for sulfur dioxide (SO2 ) where sulfur has the oxidation state +IV. Sulfites are widely used in winemaking as a preservative to protect wines from oxidation and microbial spoilage but some people are sensitive to them. A friend of mine has enjoyed white wine with sulfites for decades without a problem, but has become sensitive in his older age. This sensitivity can cause a reaction that range from mild to severe: a dry 'asthma' cough is the most common. Thus, the government requires labeling of any food or beverage with a sulfite concentration of more than 10 ppm.
The term 'sulfites' is an inclusive term for SO2 (sulfur dioxide), HSO3 - (also called hydrogen sulphite or bisulfite) and SO3 2- (sulfite). These three species are pH dependent but in wine with a pH of 3-4 the first two species predominate. Sulfur dioxide is widely used in winemaking because of its antioxidant and antibacterial properties. But you need the right amount: too little and the wine goes "off", too much and it has a pungent taste and smell. Two titration methods for the analysis of sulfur dioxide can be downloaded here . They are the Rankine Aspiration and the Ripper Method (also included below).
Sulfites in white wine - a gravimetric method
After reading an article about analysing for sulfite in wine I spoke to German wine chemist Dr Tom Mortier from the Faculty of Health and Welfare, University Colleges Leuven-Limburg, Herestraat, Belgium. He has written about it and I thought it may make a good student experiment. Tom said that the most common method for the analysis of sulfites in wine is the Ripper method (see above), which utilizes a direct titration of the wine with iodine using a starch indicator. Tom suggested we try a new analytical technique he has been working on (see May 2015 Journal of Chem Ed download here).
You can precipitate the sulfite in wine by adding strontium chloride solution. However, the normal pH of wine is about 3-4 and at that pH the sulfite is in the HSO3 - form which won't react with the Sr2+ . But if you raise the pH to 8 it will react and form a strontium sulfite precipitate which can be fltered and weighed.
Tom suggests: Take 100mL of wine, add 2.5 M NaOH dropwise until its changes to a dark brown colour. Check its pH - it should be about 8. Add 10% strontium chloride solution and it will go white if sulfites are present. Filter and weigh the precipitate. Calculate the %sulfite using stoichiometric calculations. This method is a bit vague but that's for you to experiment with.
This image, courtesy of Dr Tom Mortier, shows the colour change when white wine (left) has its alkalinity raised to pH 8 by the addition of NaOH. It changes colour and the hydrogen sulfite ion changes to sulfite. When strontium chloride is added a white precipitate of strontium sulfite forms (right) and the solution goes cloudy.
Lastly, what could your hypothesis be? You could look at sulfutes and temperature; sulfites and whether they drop if a wine sample is exposed to air. This would be a fabulous student experiment. I'm going to try it after the holidays.
Oxidation of wine to vinegar 1 - titratable acidity 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.
Natural oxidation: leaving a sample of wine
open to the air for a couple of days may make it taste "off" but this may not be a change in acidity but merely the flavour components oxidizing. It will take a bit longer to have sufficient change in acidity to be picked up by an
acid-base titration. I took about 100 mL of white wine and filled a small bottle to the brim and put the lid on. I left another 100 mL on the bench in an open beaker. A week later I titrated them both using 0.100 M NaOH and a 20 mL aliquot of the wine (by pipette) and phenolphthalein indicator. The sealed wine had a pH of 3.02 ND gave a titre of 17.80 mL. The wine left open to the air had a pH of 3.13 and gave a titre of 21.65 mL. So, a good student experiment can be made of this.
Acetobacter oxidation: You could also give the wine a kick along and instead of relying on the acetobacter floating around in the air, you could add some. This is called 'inoculating' your wine. A simple way is to get some organic vinegar from a supermarket or health food store that states on the label that it contains the "mother" - this means acetobacter . Then just and filter out the cloud (which is the suspended acetobacter) and use some of that in your wine. See photos below.
Two organic vinegars from Woolworths. The first is Italian. It looks clear but the bacteria have settled to the bottom. The second "Macro" brand is made in Australia. It is cloudy because I shook it up for the photo. The prices are for 250 mL bottles - current at October 2015.
Organic vinegar still has the acetobacter suspended in it. That's what the label means by "Mother of Vinegar". This would be good for inoculating your wine.
Bragg Organic Apple Cider Vinegar is about $6 for 473 mL. The label says "Raw - unfiltered" and "With the Mother".
I didn't want to filter out the acetobacter so I just added 2.0 mL of the vinegar to a beaker-full of wine. This would change the pH and titres I know but I though I should try. I put one sample in a sealed air-free bottle and the other in a beaker on the bench. After a week I titrated them against the 0.100 M NaOH as mentioned above. The sealed wine gave a titre of 18.15 mL and the air-exposed wine gave a titre of 24.95 mL. So, this wine oxidized too. I didn't do the pHs as there was no point. Great for a student experiment .
Some independent variables worth considering:
Aeration: As well, you probably need to give the wine plenty of aeration. Acetic acid bacteria are aerobic microorganisms and thus will not grow in an aerobic (without air) conditions. So, it might be a good idea to keep
it stirred, or having a large surface area, or bubbling air through it (from a fish tank aerator perhaps). You need to have several treatments (different surface areas, or different lengths of exposure to the air) so you can graph the results. I suggest 5 treatments with duplicates of each.
Independent variables. You may guess temperature , but what about acidity (from the natural tartaric
acid) or preservatives (perhaps SO2 ), 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 inoculated 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 Acetobacter is 25-30°C, with no growth observed at 40°C. Weak growth was
observed even as low as 10°C, but none at 8°C.2. Acidity or pH 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 (ask your teacher before you bring any to school). From the density of the
distillate (use an SG bottle) you can calculate how much to add to some
fresh wine.4. Additives. 2 in wine consists of the free form (molecular SO2 , bisufite ions HSO3 - and
sulfite ions SO3 2- ) 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 bisulfite 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.
Only the finest cask
wine is used. Use a cheap 'young' white wine so that the SO2 levels are reasonably high to start with.
Sample in the
refrigerator. Note the big surface area.
Jen calibrates the pH
meter. The pH may not tell you much as there are many different acids, many triprotic, so the relationship between acidity and pH is complex.
TITRATABLE ACIDITY - METHODS
There is one main way for measuring total acidity and that is with a titration against sodium hydroxide using a phenolphthalein indicator. There is a 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 a colour-change indicator). The method he uses (O24A-FD) comes from the Association of Analytical
Communities , AOAC. Click here to download
a copy. A pH of 8.2 is the international standard for TA endpoints in wine.
Oxidation of wine 2 - sulfites as inhibitors of oxidation
The previous section suggested that sulfites can prevent the growth of acetic acid bacteria and hence restrict oxidation. In this suggestion we look at the effect of different initial SO2 concentrations, (0, 125, 250 and 375 ppm) on the oxidation of ethanol, specifically with reference to acetaldehyde (ethanal) bonding. My thanks to Year 12 Chemistry student Codi Baker-Lahey from St Andrew's Anglican College, Sunshine Coast, Queensland, Australia for this suggestion and some of her results. It was done in 2016 under the supervision of her chemistry teacher Mrs Larsen. Codi wrote:
The oxidation of ethanol occurs in two phases. The first is the oxidisation of ethanol to ethanal. Here, if SO2 is present it binds to the ethanal and prevents the next stage from occurring. If there is no or limited SO2 then the next stage will occur - which involves further oxidation - causing the ethanal to turn into ethanoic (aceti)c acid.
The investigation supported the hypothesis that an increase in SO2 concentration would reduce oxidisation. Generally, apart from a very small number of anomalies, a higher SO2 concentration resulted in a lower titratable acidity and higher ethanol content. The most likely reason for this yielded relationship was that a higher SO2 concentration allowed more SO2 to bond to ethanal, preventing a larger amount of ethanol from oxidation to the next stage of ethanoic acid.
Four identical samples of mock wine were prepared. Mock wine was used to remove other possible bonding substances other than acetaldehyde. Sodium metabisulphite was added to change the amount of SO2 . Ethanol was monitored through an alcoholmeter and titration. It was predicted that the ethanol in all four samples would decrease as the SO2 cannot prevent this stage from occurring. This hypothesis was confirmed.
Codi's setup. She measured the free and total SO2 concentration using Ripper Method iodine titrations (see earlier); ethanol by dichromate/thiosulfate titrations (and by alcoholmeter readings of SG); and acidity by NaOH titration using phenolphthalein indicator.
This graph shows the ethanol concentration for each solution on day 8 - the final day of testing. It should be noted that each solution started with 13.5%. It can be seen that a higher SO2 initial concentration resulted in a higher final ethanol percentage; that is -higher SO2 results in less oxidation.
Oxidation of Wine 3 - rate of reaction 2 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.
WINE CHEMISTRY TEACHING RESOURCES
NON-EXPERIMENT RESEARCH QUESTIONS
If you are looking for a couple of non-experimental research questions then here are some witten as thesis statements:
1. "The pH and titratable acidity of wine are two measures of the same thing".
WINE CHEMISTRY PROFESSIONAL DEVELOPMENT
Queensland College of Wine & Tourism at Stanthorpe.
For teachers who would like to have a basic introduction to wine chemistry or would like to have a refresher, the Queensland College of Wine Tourism at Stanthorpe offers a professional development workshop that may be useful. No recommendations are made as it is up to teachers to check it out (but reports have been good). It is called the "Teachers' Wine Chemistry Professional Development ". Teachers are provided with an overview of the wine production process and the chemical processes/analysis that are conducted at each stage of wine production with an emphasis on the importance of chemistry concepts to quality wine production.
Ginger Beer - avoiding the headaches
Investigating the production of alcohol in wine can give you a few headaches - particularly if you think someone will drink your experiment. Dr Gary Turner, HoD Xavier Catholic College (Hervey Bay) suggests that for Chemistry experiments a good alternative is ginger beer using a 7-day fermentation recipe (see later). This may be also suitable if your school doesn't allow you to have alcoholic beverages on campus - even in an experiment.
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 or in the lab using one of the
many recipes available on the internet (but do NOT drink it ; not for an EEI).
Bundaberg Brewing uses the real ginger 'bug' plant.
Brew up some batches in the school lab - but don't drink it.
The suspended yeast makes it look cloudy.
Gary 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. 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
"The students can try the variables of yeast, sugar, and temperature (by dividing batches into samples) and titrate for alcohol at intervals (many hands makes light-work). You need a fridge and an incubator to give you three temps (including room temp) for the temp-as-variable". His method is one of many that can be easily obtained from the internet:
You need to create what is called a Ginger Beer Plant. Put 15g of general purpose dried yeast into a large jar or bowl, add 300mL water, 2 teaspoons ground ginger and 2 teaspoons sugar. Cover with a sheet of cling film and secure with a rubber band. Each day, for seven days, add 1 teaspoon of ginger and 1 teaspoon of sugar to the mixture in the jar. Now strain the mixture through a piece of fine muslin and add the juice of two lemons to the liquid. Add 50g or sugar to the liquid and make up to 4.5 litres with cold water, stirring to dissolve the sugar. Bottle into a plastic bottle Keep for 7-10 days when the ginger beer is sparkling and ready for drinking.
You can keep the sediment that you have left after straining the ginger beer plant. Divide into two jars and give 1 plant away to a friend with the instructions. To the sediment add 300 mL water, 2 teaspoons sugar and 2 teaspoons ginger and carry on as before.
Alcohol is produced but its concentration is likely to be under 1%. Also, most fermented soft drinks are acidified to inhibit bacterial growth. Does this also inhibit the yeast? You could investigate the effect of pH on the rate of fermentation using lemon juice or better - citric acid. The juice of 1 lemon contains about 12 g citic acid. Be warned - you should not be drinking the ginger beer unless you have approval from your teacher (and this is unlikely). Drinking stuff made in a laboratory with no hygiene controls is DEFINITELY NOT PERMITTED.
Lastly, what sort of yeast is best? Chemistry teacher Torsten Pluschke of Atherton State High School (North Queensland) said that you could buy genuine ginger beer plant (a SCOBY – symbiotic colony of bacteria and yeast, rather than just baker/brewer's yeast) online. He said that it must be looked after, and consumed in a shorter time frame than baker/brewer or wild yeasts. One online site (no recommendations given though) is The Ginger Beer Plant who can provide a 50 g sachet of the SCOBY for about $20 including postage to Australia. I asked at my local brew shop and the guy said none was available in Australia and it had to be bought online.
However, what yeast do you think Bundaberg Brewed Drinks use for their Brewed Ginger Beer? Here's what Richard Cowdroy-Ling, General Manager Business Technologies, told me about their product:
We use a standard bread yeast for the Ginger Beer and in the sparkling fruit range we use various yeast from the wine industry depending on the flavour we are chasing. In the interests of maintain a pure culture and having good control of the process for quality issues we have stayed away from the mixed cultures of yeast and bacteria [scoby] and continue to use fresh yeast for every batch. We have previously trialled a "ginger beer plant" culture in our laboratory from the UK - that was several years ago - but without much success and the flavour was not any better than on the baker's yeast that we use now.
Banana ripening kinetics - a great EEI
When you pick fruit of a tree or vine it remains alive even though it is separated from its parent. They can't get water or nutrients from the parent so they have to use their own stored chemicals to continue. Bananas are like this. Mature green bananas are full of starch, but when you pick them they begin to ripen: their average starch content just before ripening reaches 25% and drops over a few days of ripening to less than 1%. But little is known about the mechanism involved.
Bananas ripening over 7 days.
It seems that the hormone ethylene triggers the production of enzymes. For example, the alpha (1 to 4) bonds of starch may be hydrolyzed by amylases (and glucosidases) or broken by starch phosphorylases. It has been observed that sucrose starts to accumulate first, before glucose and fructose, and parallel to starch disappearance. See N. Terra et al, 1983. "Starch-Sugar Transformation during banana ripening". Journal of Food Science, V48(4) p1097-1100.
This suggests a great EEI (which could be equally as good for Biology). The question can be asked "What conditions affect the kinetics of banana ripening?" The obvious one is temperature, but an intriguing one is the presence of ethylene (ethene). This gas is produced as bananas begin to ripen so it would be instructive to compare bananas ripening in a plastic bag where the ethylene is trapped in with the fruit, versus ripening in a breeze where the ethylene is blown away. I've used a little fan out of a computer. That's all you need.
What would you measure as an index of ripening? Some books suggest iodine but that is too inaccurate for a Chemistry EEI. I suggest measuring the concentration of glucose. There are a couple of methods of measuring this but the two main ones would have to be: (a) using glucose test strips either by comparison against colour standards, or by use of a glucometer as used in diabetes monitoring; (b) by titration using either using Benedict's solution, or an iodine/thiosulfate titration (download).
Antioxidants in food and the effect of temperature
Antioxidants are a wide variety of compounds found in fruit and vegetables that are beneficial for health because they destroy free radicals and prevent tissue damage, especially in blood vessels. They are often the colorful pigments in these foods. For example, the anti-oxidant beta-carotene provides the orange hue in carrots, lycopene gives tomatoes their red appearance and anthocyanins make blueberries and certain grapes look dark purple.
The temperatures used for cooking in most household kitchens are enough to destroy particularly heat-sensitive antioxidants such as vitamin C, but the antioxidants in some foods actually become more potent with heat. For example, when tomatoes are cooked for 30 minutes at 88°C, they lose almost 30% of their Vitamin C, but 35% more of the anti-oxidant lycopene becomes available. Beta-carotene levels in carrots also increase with moderate heat. The reason seems to be that the heat breaks down the plants' thick cell walls and makes the nutrient available. [See below for reference].
This suggests a great EEI. The problem is - how can we measure antioxidant concentration in the laboratory. The Briggs-Rauscher reaction, an oscillating chemical reaction that exhibits a vivid colour change from colourless to amber to a sudden dark blue, can be used to determine relative antioxidant concentration in foods (and wine). The reaction is very complex and involves both iodide ions and iodine molecules. It is an "oscillating" reaction that goes from blue to colourless to yellow to blue in a period of time that depends on the relative concentrations of these species.
WARNING: Before you get too carried away with this experiment, you should be aware that one of the chemicals involved is malonic acid. This is not too hard to get but you need to plan ahead. I bought some through Labtek (Australia) but it took almost 4 weeks. The one I got was Ajax 'Unilab' brand, code 305-100G, $63.80 for a 100 g bottle. Griffith University would have helped as part of their outreach program to schools out but they would have to order it in too. I know of three girls at Palm Beach Currumbin State High School (hello to Jissa, Momoka and Niki) on the Gold Coast, Queensland, who started to do this EEI but ran into time constraints. They told me that they found a website "that may supply future students with a kit, inclusive of malonic acid, that will allow for them to do the BR reaction: http://www.teachersource.com/product/fascinating-oscillating-reaction-kit/chemistry ."
One complete oscillation of the Briggs-Rauscher reaction is shown here. The time it takes to go from dark blue to the next dark blue will be lengthened by the presence of anti-oxidant molecules. If one oscillation took 30 seconds in the absence of anti-oxidants then when a fruit sample is added it make take say 50 seconds.
To be more accurate you could let it go for say 10 oscillations. Image http://advan.physiology.org/content/ajpadvan/35/4/427/F2.large.jpg
Using this reaction, it is possible to determine the level of antioxidants present in everyday foods and drinks by starting a timer when the reaction mixture is blue and noting how long it takes for the reaction mixture to go through one full cycle (stopping the stopwatch when the mixture is blue again). The addition of antioxidants increases the time taken for this reaction. The longer the time taken for the reaction cycle, the more antioxidants the food contains. A terrific resource which includes a detailed method is available at Learning and Teaching Scotland .
Reference: "Thermal Processing Enhances the Nutritional Value of Tomatoes by Increasing Total Antioxidant Activity" by Veronica Dewanto, Xianzhong Wu, Kafui Adom and Rui Hai Liu, published in the Journal of Agricultural Food Chem ., 2002, 50 (10), pp 3010-3014 http://pubs.acs.org/doi/full/10.1021/jf0115589 .
Factors affecting the viscosity of cellulose gum
Carboxymethylcellulose (CMC) is a modified cellulose gum that can be used as a food thickener (Code E461). It produces a clear, slightly gummy, solution in water. It is used to thicken dry mix beverages like soups, as well as syrups and ice cream. In the wine industry it has recently been approved as an additive to keep wine clear. CMC acts upon the face of a growing crystal, restricting further growth while ensuring that nothing is visible to the naked eye. Metal ions with different charges significantly affect the viscosity of aqueous sodium carboxylmethylcellulose solution. And herein like the essence of a great EEI.
You could take an aqueous sodium CMC solution and add various metal ions to it. You could try metal salts such as LiCl, NaCl, KCl, CaCl2, AlCl3 and make aqueous 0.50% (w/v) solutions in the range 0.00-0.10 mol/kg. To measure viscosity you could use a tube (say 2 cm diameter) and about 100 cm long (or even a 100 mL measuring cylinder). Put marks as far apart as possible, maybe 50cm apart (or at the 80 and 20 mL marks on the measuring cylinder). Fill it with the solutions and drop small (plastic) balls and time between the marks. You want balls that are not too dense but will not float (about 1.020 g/cm3 seems okay). Better note the temperature.
The viscosity of the sample was then calculated from the following (Poiseuille) equation η/ηo = ρt/ρo to where ηo (viscosity) is in mPa s, ρo (density) is in g cm-3 , to (flow time) is in s, ρ is in g cm-3 . The standard values of viscosity (ηo) and density (ρo ) for water at 25 °C can be taken from the literature, while the densities of samples can be measured by whatever method you like.
Effects of metal ions on viscosity of aqueous 0.50% (w/v) CMC solution at 25 °C.
I bought some CC from a health food shop and it cost $44 for 200g. You can get it cheaper from Aliexpress but the postage is high.
The graph shows that metal ions have significant effects on the viscosity of aqueous 0.50%w/v CMC solutions and its extent is particularly dependent on metal ion charges. NaCl and LiCl and the same as KCl. The cation size effects of 1++ , Na+ , and K+ on viscosity are all very similar. The viscosity of the aqueous CMC solutions gradually decreases with the increase of alkali metal chloride concentration. On the other hand, the effect of calcium ion with 2+2 concentration. However, in the case of 3+3+ you will get a real surprise.
My thanks to Seng Set and Masakazu Kita, Department of Chemistry, Okayama University, Japan, and David Ford Department of Chemistry, Royal University of Phnom Penh, Vietnam. Refer J. Chem. Educ. 2015, 92, 946-949.
Nitrification in Soils - a great titration prac.
Nitrogen is one of the most essential nutrients for plants and is most frequently the limiting factor in crop productivity. The vast majority of the total nitrogen in soil (>98%) is in organic matter, which can't directly be used by plants. It must be converted to inorganic forms as either ammonium (NH4 + ) or nitrate (NO3 - ). This conversion occurs via a biological process called nitrification involving soil microorganisms.
Nitrification in nature is a two-step oxidation process of ammonium ion or ammonia to nitrate ion catalyzed by two types of bacteria. The first reaction is oxidation of ammonium to nitrite by ammonium oxidizing bacteria (AOB) represented by the Nitrosomonas species. The second reaction is oxidation of nitrite (NO2 - ) to nitrate by nitrate-oxidizing bacteria (NOB), represented by the Nitrobacter species.
The progress of nitrification can be investigated using a fairly common titration procedure. It involves making a filtered solution of the soil, boiling off free ammonia, adding an excess of sodium hydroxide solution (which reacts with the ammonium ion) and back-titrating the excess hydroxide with standardized hydrochloric acid. The method can be easily found on the internet and may even be in your textbook under a heading of "analyzing fertilizer". The photos below are courtesy of Yr 12 student Jamie from Our Lady's College, Annerley, Brisbane.
Sieving gets rid of the larger bits of organic matter.
Weighed and ready to go into the incubating flasks. If you use an unsealed container you may need to weigh the sample each day and add a few drops of water to keep the mass constant.
Getting ready for titration. Yr 12 students in the Chemistry lab at Our Lady's College, Annerley.
A good mixed source of nitrifying bacteria and ammonium ions is potting mix. It usually contains peat moss, sand, and other organic material such as wood chips. It would have reasonable amounts of ammonium ions present. Your EEI could be to measure the ammonium ion concentration after a period of time (say a week) under different conditions that may affect the bacteria (temperature, oxygen availability, pH, salinity, light).
Filtering the soil and water mixture at the end. The filtrate contains the soluble ammonium ions that are to be back-titrated.
Temperature: nitrifiying bacteria are supposed to be at their optimum from 25°C-30°C. Aquarium operators have a rule of thumb that says: 50% activity at 20°C, 25% at 10°C, 0% at 4°C and are death at 0°C. When heated, their activity decreases after 30°C and they die by 50°C. That suggests a great investigation.
Oxygen: the bacteria need oxygen to produce energy to live (respiration). If you cut off their oxygen supply their activity decreases. Farmers know that in waterlogged soils the bacteria are less productive. That suggests waterlogging your potting mix samples to prevent the bacteria getting oxygen and comparing them to the control. It would be good to know how much water is present in the potting mix and you can do this by taking a weighed sample and drying it in the oven at 100°C (gravimetric analysis).
Because the hydroxide is in the flask, the end point will be when the pink just disappears.
The flasks at the end point.
Moisture: nitrifying bacteria are most active in a soil that is moist but not saturated (see above). Is there an optimum amount of moisture? If completely dry they go into hibernation - but what is optimum?
pH: there is an optimum for the first stage bacteria (nitrosomanas ) of 7.8 to 8.0, and for stage two (nitrobacter) it is a little bit less. This investigation would require careful design. Potting mix is usually pH 5-6.5 so you'll have to add some alkali to increase the pH. This will affect the acid/base titration. You'll need to think about a controlled sample without potting mix but with the added base. As a matter of interest, as the ocean becomes enriched in anthropogenic (human activity) CO2 the resulting decrease in pH could lead to decreasing rates of nitrification. That's another context for this EEI.
Salinity: the optimum is said to be zero to 0.6%.
Light: the nitrifying bacteria are supposed to be sensitive to blue/UV light. But how far into the soil would light penetrate (a good design consideration).
My thanks to Chemistry colleague at Our Lady's College, Annerley, Brisbane, Kayleen Solomon and Yr 12 student Jamie Nguyen for providing ideas and photos for this EEI suggestion. A copy of Jamie's method can be downloaded here .
Nitrification in Soils - spectrographic determination
In the above suggestion, a back titration was used to determine the ammonium concentration. If you have access to a spectrophotometer and the reagents required you could determine nitrate ion spectrophophotometrically.
Physically, nitrite is a colorless and odorless ion. However, there are spectrophotometric methods for its determination. One example is to react it with an acidified sulfanilamide and N-(1-naphthyl)ethylenediamine dihydrochloride to form an intensely coloured dye having maximum absorption at 540 nm. There are lots of other methods. Often you can get them in the form of test kits where you compare the colour to a test strip. Before you get started check that you can get the reagents.
Stability of Vitamin C in solution - analysed by titration. + ], 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 with time 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.
Here are some data to get you started. In an examination of some Vitamin C products for the pharmaceutical industry Uprety and Revis (1964) reported that the factors below were important (but not necessarily linear in their response to the manipulated variable). There were a lot of anolomous data in their results.
Temperature: after 30 days 20% of the ascorbic acid was lost at 37°C, whereas 46% was lost at 55°C.
pH and acidity: after 30 days 33% of the ascorbic acid was lost at pH 2.5 and 35% lost at pH 6.5, thus low pH seems to stop the loss of ascorbic acid. However, there were lots of anomalies and in some samples the losses went down as the pH increased. Also note that when fruit is cut, the enzyme polyphenol oxidase is released from the cells and reacts with the oxygen in the air causing the fruit to deteriorate. When you decrease the pH by adding citric acid you tend to stop the polyphenol oxidase working as its optimum range is from pH 5 to 7. In fact, below a pH level of 3.0, the enzyme is inactivated. In the Uprety and Revis article it was found that citric acid was protected ascorbic acid well. An apple juice sample lost 98.6% of its ascorbic acid after 60 days, but when 0.2% citric acid was added it lost just 59% (they didn't say the pHs. The pH of lemon juice is in the 2.0 range, making it very effective against browning. Now there's a great EEI.
Antioxidants: Citric acid can be considered to be an antioxidant but mainly for its low pH rather than anything special about the citric acid molecule (other fruit acids would work if their pH was low enough). However, it is also known that EDTA (ethylene diamine tetraacetic acid) protects juices in certain cases. It is believed that the EDTA binds any copper ions present. A 0.01% solution seems to work with apple juice. Sodium chloride also seems to work - a 0.9% solution is believed to counter the oxidation of ascorbic acid.
Natural juice vs water: A sample of lime juice at pH 3.5 for 30 days at 37°C lost 76% of its ascorbic acid, whereas distilled water at the same pH and temperature lost only 50% in the same time. Now that discounts the pH argument. There must be something else that goes on too.
Different juices: some juices have substances which help destroy the ascorbic acid. For example, apple juice held at 37°C, pH 3.5 for 60 days lost 98.6% of its ascorbic acid, whereas pineapple juice under the identical conditions lost just 71.1%. This is attributed to the presence of polyphenol-oxidase in the apple juice which catalyses the oxidation of ascorbic acid.
Head Space: the bigger the headspace in a sealed container the greater the loss of ascorbic acid. The more access the juice has to oxygen the greater the degredation. Bubbling air through a sample would speed things up you would think.
Colour change: as ascorbic acid breaks down the juices got darker in colour. This is a good index of ascorbic acid concentration if you have a colorimeter. See my smartphone colorimeter suggestion later on this webpage. There's a free app from your iPhone or Android that will turn them into a colorimeter. Works well.
Titrations
A DCPIP method that has been trialled extensively
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 .
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 and filter it the next day.
THE COLOUR OF DCPIP IN ACID
The blue DCPIP solution goes pink when it mixes with acid - any acid.This is its acid colour. It does this even when no ascorbic acid is present. Students get confused because they think the pink is the colour DCPIP changes to when it reacts with ascorbic acid. It is not. When it reacts with ascorbic acid the pink DCPIP goes colourless. When all of the ascorbic acid is used up (at the end point) there is no ascorbic acid left to make it go colourless so it goes pink (permanently).
You can see the blue DCPIP falling into the flask from the burette. When it hits the acid it is blue until it swirls around and reacts to go pink. The acid is not ascorbic acid - it is just extracts solution.
STANDARDIZING DCPIP WITH STANDARD ASCORBIC ACID SOLUTION
1. A 20 mL aliquot of ascorbic acid solution in the flask, along with some acidic extraction solution. No DCPIP has been added.
2. The blue DCPIP is added from the burette and the solution goes pink where the DCPIP meets the acid solution. If you stop adding DCPIP it goes colourless.
3. At the end point all of the ascorbic acid has been used up by the DCPIP so it remains pink. You should go from colourless to a very faint pink with just one drop.
4. This is way past the end point. As you keep adding more DCPIP it just keeps getting a darker pink. This is because the other acids in the flask make it go to its acid colour which is pink.
TITRATING ORANGE JUICE WITH DCPIP
1. A 2 mL aliquot of orange juice is placed in the flask with some extract solution. No DCPIP has been added. My thanks to Tara Robinson for her orange juice.
2. The blue DCPIP is added from the burette and the solution goes pink where the DCPIP meets the acid solution. If you stop adding DCPIP it goes colourless but this is a bit hard to see because of the orange juice colour.
3. At the end point all of the ascorbic acid has been used up by the DCPIP so it remains pink. You should go from orange (from the juice) to a very faint pinkish orange with just one drop.
4. This is way past the end point. As you keep adding more DCPIP it just keeps getting a darker pink/orange. This is because the other acids in the flask make it go to its acid colour which is pink.
IODINE TITRATION OF ASCORBIC ACID IN FRUIT JUICE
The
second method is by iodine titration and uses cheaper and more easily obtained
chemicals. In essence, you make up a potassium iodide (KIO3 )and potassium iodate (KI) solution in acid.
The iodate ion IO3 - reacts with the iodide ion I- in acidic H+ solution to form iodine I2 and water H2 O.
2 IO3 - +10 I- + 12 H+ → 6 I2 + 6 H2 O
During the titration the iodine reacts with ascorbic acid C6 H8 O6 in your sample to form dehydroascorbic acid C6 H6 O6 , iodide ions I- and hydrogen ions H+
6 I2 + 6 C6 H8 O6 → 6 C6 H6 O6 +12 I- + 12 H+
When all of the ascorbic acid is used up by the iodine at the end point, the excess iodine turns the starch blue.
The overall reaction can be seen by the addition of the above reactions and the coefficients simplified:
IO3 - +5 I- + 3 C6 H8 O6 →
6 I- + 3 C6 H6 O6 + 3 H2 O
The stoichiometric ratio is:
1 IO3 - + 3 C6 H8 O6
or
n (IO3 - )/1 = n (C6 H6 O6 )/3
C iodate x V iodate = m vitc/ M vitc/3
25 mL orange juice and the first drops of the iodine solution
Way past the end point
For a copy of this method with sample calculations click:Vitamin C Iodine titration .
My thanks to Tara Robinson - Yr 12 Chem at MBC - for the photos.
Soft drink ('soda' or 'pop' in the US) goes flat when you leave it out for a day or so. When the lid is on there is an equilibrium between the dissolved carbon dioxide in the liquid and the CO2 gas in the headspace. When the lid is opened the pressure is relieved and the equilibrium is disturbed. According to Le Chatelier's Principle the system tends to oppose the removal of the gas from the headspace so more and more carbon dioxide comes out of solution to replace it. As the gas can never build up sufficient pressure above the surface of the liquid, the gas keeps coming out of the solution until it is 'flat'. But we can measure the [CO2 ] in the drink by titration as the CO2 forms the weak acid H2 CO3 and this can be titrated against a solution of the strong base NaOH. The problem is that soft drink has citric acid in it as well. We can measure the total acidity of the soft drink by titration, and then decarbonate it (make it 'flat') and measure just the citric acid by titration. The difference in titres is a measure of the [CO2 ].
This is a great experiment and will develop you analytical techniques. However, it doesn't make a good student investigation as there is no mention of relationships between variables as expected in a student experiment. You could measure the [CO2 ] of different drinks but that hardly forms a good reasearch question. For a good student experiment based on this technique, see the next suggestion further down. Here are some hints, some typical results, and worked calculations.
HINT 1. You can standardized the approximately 0.1 M NaOH solution against the solid acid potassium hydrogen phthalate (KHP) - a weak monobasic acid that is used as a primary standard. If you don't have KHP, you can prepare approximately 0.1 M HCL and standardize it against the sodium hydrogen carbonate NaHCO3 and then standardize the NaOH solution against the standard HCl solution.
HINT 2. When using a coloured indicator such as phenolphthalein for the titration, you need a clear uncoloured drink so you can see the end point. Lemonade works well.
HINT 3. If you are testing a coloured drink you will need a pH meter to determine the end point (pH = 8).
HINT 4. It is impossible to pipette an aliquot of fizzy soft drink as the bubbles keep expanding inside the pipette. You need to weigh out a sample on a balance. I measured out 20.00 grams and titrated that against 0.100 M NaOH using phenolphthalein indicator.
HINT 5. It is hard to get three titres exactly the same as the pouring of the sample (aliquot) into the flask disrupts the rest of the liquid. You need to think of a way to remove three samples from the container without disturbing the container too much. I measured out three samples into separate flasks at the one time. That seemed to give low spread of results (low uncertainty = high precision).
TYPICAL RESULTS. For 20.00 g aliquots of Schweppes lemonade I had titres averaging 10.60 mL. Then I flattened (decarbonated) the lemonade and measured out 20.00 g again and had average titres of 7.50 mL. This gives a citric acid concentration of 0.240% (which is about right), and a CO2 concentration of 0.00775 M which is 0.034 g/100g (0.034% w/w). This is the CO2 concentration about 40 minutes after opening. The CO2 concentration of the drink inside the can before opening would be much higher. It is likely to be about 0.37 % w/w or 0.084 mol L-1 . Calculations are shown below:
CALCULATIONS FOR CITRIC ACID 2 COOH-C(OH)-COOH-CH2 COOHcitric = CNaOH VNaOH /3 = 0.100 x 0.00750/3 = 0.000250 mol citric acid in 20.00 mLcitric = nM = 0.000250 x 192.12 = 0.0480 g citric acid in 20.00 mLCALCULATIONS FOR CARBON DIOXIDE 2 + 2 NaOH → Na2 CO3 + H2 OCO2 = n NaOH /2 = CNaOH VNaOH /2CO2 =0.100 x 0.00310/2 = 0.000155 mol in 20.0 g (20.0 mL)CO2 = n/V = 0.000155 mol/0.020 L = 0.00775 MCO2 = nM = 0.000155 x 44 = 0.00682 g in 20.00 mLCO2 = mCO2 /20.00 x 100 = 0.034% (or 0.034 g /100 g)
Rate of carbon dioxide loss in soft drink - by titration. A nice modification As I said above, measuring the CO2 concentration is only useful for a student experiment if it is a part of a bigger experiment with a valid research question. It would be interesting to see how fast soft drink goes flat. Thus, a question could be asked: what is the relationship between the CO2 concentration of soft drink in an open container, and time elapsed?
HINT 1. You need to decarbonate some soft drink to use as a control. Heating it to say 80o C is the simplest way to achieve this. No need to boil it though.
HINT 2. The rate of carbon dioxide loss from the drink is fast initially and then tapers off exponentially as time elapses. Try to do 4 or 5 trials on the one day and one more the next day. At 24 hours the sample should be basically flat but worth including.
HINT 3. The 5 x 3 approach is useful to ensure you collect 'sufficent' data. That is, five variations of the independent variable (the time intervals, eg 30, 60, 90, 120, 240 minutes, and a final one at 24 hours or so) and three replicate titrations for each of the 5 trials. For replicates, you could just take 20.00 g aliquots one sample, or have three separate 20.00 g samples for each time interval. One approach is to weigh out 20.00 grams of the soft drink into each of 15 containers (beakers or specimen jars) at the start. At each time interval titrate three of the samples. That gives you three equivalent replicates and the variation between the replicate titres is likely to be small. Ideally, you would like the replicates to be within 1 drop of each other (that is, +/- 0.05 mL). Aim for that but don't be disappointed if they are +/- 0.5 mL as it hard to control bubbly drinks.
HINT 3. For the theory behind decarbonation of soft drink, see the next section where I look at the less complicated situation of soda water going flat.
Rate of carbon dioxide loss in soda water as it goes flat. A great modification.
To avoid the complication of residual acidity caused by the presence of citric acid (or phosphoric acid in the case of cola drinks), why not try the same experiment using plain soda water - which has no added acid or acidic fruit syrups. This would be a great modification to either of the two experiments above. It is not a redirection as you are measuring the same dependent and independent variables. It is more of an extension. You are extending the scope of the experiment by including drink without residual acidity.
HINT 1. Check the label and make sure the drink is just water and carbon dioxide. It will say 'carbonated water'. Other carbonated drinks such as Club Soda and 'sparkling mineral water' contain added or dissolved minerals such as potassium bicarbonate, sodium bicarbonate, sodium citrate, or potassium sulfate. Use 'soda water' if you can to avoid the presence of other substances. If you can't find soda water, use sparkling mineral water.
Woolies Soda Water. Ingredients: carbonated water
HINT 2. The titre for flat soda water should be (almost) zero but if it is not then you have some further checking to do. Check the label again to see that it is 'carbonated water' and not 'carbonated mineral water'. If it is just carbonated water then it is likely you have contamination of the solutions. Perhaps the flasks had some residual HCl in them from an earlier standardisation. Rinse everything again and redo the titrations.
HINT 3. If your titres are less than 10 mL, you may like to use 0.05 M NaOH instead of 0.1 M NaOH. It's up to you.
HINT 4. Keep the lid on the 0.1 M NaOH solution unless you are filling the burette. NaOH solution will absorb CO2 from the atmosphere and react to form sodium carbonate. This reduces the concnetration of the NaOH solution and will add a systematic error to your titres.
THEORY BEHIND RATE OF DECARBONATION POINT 1. With the lid on, there is an equilibrium between carbon dioxide gas in the
headspace and the dissolved carbon dioxide in the water. Henry's Law gives a relationship between the pressure of the CO2 gas in the headspace and the amount of CO2 gas dissolved in the water. Henry's Law is P = kT C. Here's a sample question and calculation to show how it applies. 2 in 1.00 L of soda water under a headspace pressure of 3.50 atmospheres (atm). Note: Henry's Law constant kT = 29.76 atm mol L-1 .T C) can be rearranged to give C = P/kT = (2.50 atm)/(29.76 atm mol L-1 ) = 0.084 mol L-1 .-1 = 3.70 g (3 sf).
POINT 2. You can use Henry's Law to work out the theoretical concentration of CO2 in water when exposed to just normal atmospheric pressure (1 atm, or 101.3 kPa). The current atmospheric concentration of CO2 is 360 micro-atmospheres (360 μ atm). It is rising because of human action (or inaction) but let's just use the accepted value for the moment. Henry's Law: C = P/kT = (360 x 10-6 atm)/(29.76 atm mol L-1 ) = 1.21 x 10-5 mol L-1 -5 mol x 44 g mol-1 = 5.32x 10-4 g in 1 litre (3 sf).
POINT 3. the mole ratios in the titration reaction are :NaOH VNaOH /2 = CCO2 VCO2 NaOH = 2 x (CCO2 VCO2 )/CNaOH = (2 x 1.21 x 10-5 x 20.00)
POINT 4. When CO2 dissolves in water, the following sequence of reactions takes place:2 (g) → CO2 (aq) 2 (aq) + H2 O (l) → H2 CO3 (aq)2 CO3 (aq) + H2 O (l) → HCO3 - (aq) + H3 O+ (aq)3 - (aq) + H2 O (l) → CO3 2- (aq) + H3 O+ (aq)2 (g) + H2 O (l) → CO3 2- (aq) + 2H+ (aq)2 (g) comes out of solution in the form of bubbles.
POINT 5. The rate of decarbonation depends on three processes:3 2- ) and hydrogen carbonate ion (HCO3 - ) revert to carbonic acid (H2 CO3 ) and thence to dissolved CO2 (aq) and finally gaseous CO2 . This is controlled by the activation energies of the reverse reactions and any rate determining steps.2 from the liquid to the bubble. Diffusion is affected by the temperature and pressure of the gas above the liquid.
• the distance they have to travel (depth of the liquid),
POINT 4. The invisible diffusion of dissolved CO2
through the water surface.
This will be proportional to the surface area of the open container.
So, you can see, there is no single factor that determines the relationship between the rate of CO2 loss and time elapsed. It will be rather complex.
RESULTS. Here are some of my results. Yours will have different values but the shape should be similar. I didn't do triplicates for each data point as you would, so I can't show error bars.
This appears to be an exponential relationship (exponential decay) so we can linearise this by plotting the natural log (ln) of the concentration vs time elapsed. Note: you would not be expected to know how to linearise an exponential graph in senior chemistry, and you would not be penalised for not showing it. It is just an added piece of evidence to support your claim about the relationship - but is not necessary.Carbon dioxide in soft drinks - effect of temperature. More titrations. 2 in soda water or 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. This is a redirection as the independent variable has been changed from 'time elapsed' to '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 H2 CO3 , carbonic acid: CO2 (aq) + H2 O(l)
→ H2 CO3 (aq)
The reaction is reversible; that is, the
products can react to form the reactants. 2 exists as carbonic acid H2 CO3 . The other 99.5% exists as
CO2 (aq). Carbonic acid is a weak acid which dissociates in two steps. Step 1: H2 CO3 → H+ + HCO3 -
Ka1 = 4.2 x 10-7 ; Step 2: HCO3 - → H+ + CO3 2-
Ka2 = 4.8 x 10-11 . When titrated, all the CO2 (aq)
is reacted so a titration is a measure of total CO2 content: H2 CO3 (aq)
+ 2NaOH(aq) → Na2 CO3 + 2H2 O.
POINT 1. Students often say 'if only a tiny amount of CO2 exists as carbonic acid H2 CO3 , won't a titration just pick up the H2 CO3 and not the vast amount of the dissolved carbon dioxide (CO2 (aq))?' The answer is that as the H2 CO3 gets used up in a titration reaction, more is produced as the equilibrium for the reaction CO2 (aq) + H2 O(l)
→ H2 CO3 (aq)
shifts to the right.
POINT 2. The enthalpy changes for the various reactions will allow a prediction to be made about the effect of increased temperature (that is, when energy in the form of heat is added). The ΔH (heats of reaction) value for the reactions are as follows:2 (g) → CO2 (aq); ΔH = -400.0 kJ mol-1 exothermic2 (aq) + H2 O (l) → H2 CO3 (aq) ; ΔH = +21.8 kJ mol-1 endothermic2 CO3 (aq) + H2 O (l) → HCO3 - (aq) + H3 O+ (aq) ;ΔH = -35.8kJ mol-1 exothermic3 - (aq) + H2 O (l) → CO3 2- (aq) + H3 O+ (aq) ;ΔH = -14.8 kJ mol-1 exothermic2 (g) + H2 O (l) → CO3 2- (aq) + 2H+ (aq) ; ΔH = -428.8 kJ mol-1 exothermic2 (g) comes out of solution in the form of bubbles when it is heated.
HINT 1. Try to do the samples from different temperatures as quickly as possible. The carbon dioxide concentration decreases with time elapsed once the bottle is opened, so you need to control time as best you can. For example, it drops by about 0.5 g/L in 30 minutes after opening.
HINT 2. Normally, you'd aim for titres of approximately the same as the aliquot being used (20 mL). When using 0.1 M NaOH you may find the titres are only half that or less. You may like to use 0.05 M NaOH instead.
Cooling the drink on ice.
Titrate against 0.1M or 0.05 M standardised sodium hydroxide solution
Phenolphthalein is the best indicator for a weak acid/strong base titration (pH 8).
A faint pink color should persist for 30
seconds at the end point. Use a white tile or filter paper for a background.
My thanks to Year 12 Chemistry student Codi Baker-Lahey from St Andrew's Anglican College, Sunshine Coast, Queensland, Australia for sharing the results of her EEI (2016) done under the supervision of her teacher Mrs Larsen.
Determination of total carbon dioxide by titration. Another marvellous modification. 2 in soft drink involves heating the beverage in a flask and capturing the CO2 in a balloon placed over the top (that contains an excess of sodium hydroxide solution). The evolved carbon dioxide gas that passes into the balloon is absorbed and converted into an equivalent amount of sodium carbonate. The resulting mixture consisting of the excess sodium hydroxide and sodium carbonate can be titrated with standard HCl.
Titration to the first colorless phenolphthalein endpoint neutralizes the excess sodium hydroxide and converts all of the sodium carbonate into sodium bicarbonate. Methyl orange indicator is added and the titration is continued to the second endpoint (where methyl orange changes colour) as the acid converts the sodium bicarbonate to water and carbon dioxide. The difference in milliliters between the first and second endpoints is used to calculate the carbon dioxide present in the sample or the grams of substance being sought.
The NaOH in the balloon absorbs the carbon dioxide driven off from the heated soft drink.
Turn the hotplate off when the balloon stops growing and let it cool overnight.
As a former industrial chemist at Golden Circle Cannery (soft drink manufacturer) I tried this method: I poured 100mL of chilled soft drink into a chilled 500mL conical flask and added a magnetic stirring bar. Then I poured exactly 40 mL of 1.00M NaOH into a balloon and immediately stretched the balloon over the neck of the flask - being careful not to spill any of it into the flask. If the balloon is positioned as illustrated in the photo above there is no danger of the NaOH entering into the reaction flask.
I stirred the flask at slowly first, then after the vigorous fizzing has ended, I continued with greater agitation for 15 minutes and then turned it off and allowed the flask to stand overnight. In the morning the balloon had collapsed - demonstrating the absorption of the CO2 by the NaOH.
I carefully transferred the contents of the balloon into a 250 mL conical flask and rinsed the balloon out with some distilled water. Then I titrated the solution (using 4 drops of phenolphthalein indicator) with standard HCl in the burette (approx 1M) until it reached a colorless end point. This titre (V1) was recorded. Then I added four drops of methyl orange (that turns yellow when added) and completed the titration until an orange color was obtained. I recorded the volume (V2) used to the second endpoint.
It can be shown that the difference in titres (V2-V1) is related to the amount of NaHCO3 produced which is equal to the amount of CO2 produced, that is: (V2-V1) x CHCl = n CO2 = mCO2 /Mr CO2 .
For Schweppes Lemonade I had the following results using 1.09M HCl: V1 = 11.1 mL, V2 = 15.6 mL, V2-V1 = 4.5 mL.HCl = n CO2 = m(CO2 )/M(CO2 )2 ) = 0.196g (in 100mL beverage) = 1.96 g/L (1L gas per 1L of beverage).
This method is based on one reported by Crossno, Kalbus & Kalbus from California State University in the Journal of Chemical Education V73(2), Feb 1996, p175. It is now available on-line here . If that link gets broken, download it off my site here .
Inversion of sucrose.
Table sugar 'sucrose' is a disaccharide made up of a glucose and fructose molecule joined together by a glycosidic linkage. Hydrolysis breaks the glycosidic bond and converts it into glucose and fructose:
C12 H22 O11 + H2 O + H
+ → C6 H12 O6 + C6 H12 O6 + H
+
There are two main ways of hydrolysing sucrose:
1. ACID HYDROLYSIS - This method involves the addition of weak acids, such as tartaric acid or citric acid (lemon juice). Likewise, gastric (stomach) acidity hydrolyses sucrose. Dependent variable: glucose concentration Independent variables: sucrose concentration, pH (using weak acids), temperature (rate of inversion increases exponentially with temperature). Method: I'd start with a 10g/100mL sucrose solution, temperature 25°C, pH7. Measure the glucose every 10 minutes with the glucose meter.
2. ENZYMES - This method involves the adding the enzyme sucrase (invertase) as found in baker's yeast. Controlled variables: Baker's yeast suspension of 0.5g/100mL Method: I'd also start with a 10g/100mL sucrose solution, temperature of say 25°C, pH7. It would also be good to add some nutrient so a Sorensen buffer (NaH2PO4 0.066 M, K2HPO4 0.066 M) made to the various pHs would be handy. Measure the glucose every 10 minutes with the meter.
NOTE: if you'd like to measure the glucose more accurately (by titration) then see the suggestion below about the ripening of bananas.
Conductivity of mixtures
The electrical conductivity of aqueous solutions has been extensively studied and reviewed for the past several decades. However, most of the theoretical treatments and experimental investigations have been directed toward systems containing only a single cation/anion pair. There are lots of data tables around showing the conductivity of aqueous solutions as a function of their concentration. For example the conductivity of a 5% (w/w) potassium nitrate solution is 58 μS/cm but a 10% solution is just 96 μS/cm (not double as you might expect). The same is true of sodium nitrate 52 uS/cm for a 5% solution but 92 for a 10% one.
The point is that conductivity is not a linear relationship with concentration. Anyone can just look up the tables though. However, there seems nothing published on mixtures. For example, if you mixed together equal volumes of KNO3 and NaNO3 would the resultant conductivity be the average (96 + 92)/2 = 94 μS/cm? You could also argue that each solution has its volume doubled so is half the original concentration. Thus, you could just add the two 5% values and get 58 + 52 = 110 μS/cm. A good EEI would be to explore the additivity of conductivities of aqueous ionic solutions. I can't wait to try it when I get back to school. Download it here if you're game .
Conductivity of solutions that precipitate
Imagine if you were making solutions of electrolytes in water for use in dialysis and a precipitate formed. What a mess you could be in. The quality of the water used in the mixing of the electrolyte and the maintenance of dialysis systems is also critically important and this can be measured by conductivity.
Chemical constituents of the water can change the ionic composition of the dialysate thus altering the concentration gradient in the dialyzer; react with constituents of dialysate or blood changing the chemical composition of the dialysate prescription or generating unwanted precipitates. This is just to highlight how important knowing how conductivity changes when precipitation occurs.
A good EEI can be made out of a fairly typical 1st year university chemistry prac. I would take 50 mL of 0.1 M copper (II) nitrate and place it in a beaker with a conductivity probe. Then I would add 0.1M sodium carbonate solution from a burette and note the conductivity as the carbonate was added. A plot of volume of carbonate added vs conductivity would be fascinating. Of course, as the carbonate is added it causes copper carbonate to precipitate so the copper ions and carbonate ions are removed from solution. The conductivity should decrease until the equivalence point (50 mL of each). Here's a graph of the theoretical calculations.
Heat of combustion C values of methanol, ethanol, propanol; or even of the three C4 H9 OH
isomers. Why might the accuracy be different? What does this tell you about
intramolecular bonding? Are there any correlations with BPt?
A note of caution: teachers report that they have experienced safety issues with these burners (cracking and catching on fire, or carelessly spilt liquid igniting, and so on). Testing petrol or other highly volatile liquids seems fraught with danger but many teachers have reported that there were no problems. Strict supervision would seem necessary.
Heat of combustion of mixtures 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
ΔHc values (for example if Fuel A has a
ΔHc of 1000 kJ/mol and Fuel B has a ΔHc of 2000 kJ/mol does a 50:50 mixture of the two have a ΔHc of 1500 kJ/mol.
A basic setup for calorimetry. You need to consider a chimney and a lid on the calorimeter. My thanks to Lily for her help.
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 ΔHc 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 ethanol: water mixtures (and expert commentary)
Ethanol is currently being considered as a potential alternative to traditional fuels. However, ethanol offers a low return in terms of energy output per dollar invested when compared to fossil fuels. More than one-third of the cost associated with bio-ethanol production is devoted to distillation and water removal. A good EEI would be to look at the heat of combustion of ethanol-water mixtures. The research question of interest is "is hydrous ethanol a practical fuel to be used in lieu of anhydrous ethanol?" Hydrous ethanol is the chemical you have in your high school science lab: it is 95% ethanol and 5% water (E95/W5). The other type - anhydrous or absolute ethanol - is 100% ethanol and is more expensive.
In his 2012 Master of Engineering thesis from Louisiana State University, science graduate Baine Breaux investigated how water affected the combustion of ethanol. His study titled "The Effect of Elevated Water Content on Ethanol Combustion", validated up to 20% water in ethanol as a practical fuel for continuous flame applications. He said that such a fuel can be produced at a lower capital cost than pure ethanol and would provide an economic benefit despite increased volumetric consumption. The use of up to E80/W20 (that is 80% ethanol, 20% water), he said, offered a reduction in exhaust NOx concentration and a reduction in peak flame temperatures without reducing combustion efficiency or exhaust gas temperature. Baine is now Lead Development Engineer at Hiltner Combustion Systems, Washington, and gave me permission to use his graph and photos below.
Graph taken from Baine Breaux's thesis. It doesn't show enthalpy data but does show The flame temperature of the various mixtures. The images show the shape of the flame.
Images of the flames at various mixture combinations. There is no flame with the 30% water mixture at bottom right. The ER value is about the proportion of air allowed for combustion.
Baine found that mixtures with up to 60% water can be burned but require pre-heating to 140°C (but he had a special calorimeter for this whereas you are just using a spirit burner). While ethanol has a heat of combustion of -1367 kJ/mol, for it to burn it has to be vaporised and that requires a mere +38.19 kJ/mol (= Lv ). Water too has to be vaporised and that takes just a bit more (Lv = +40.66 kJ/mol). However, when you do it on a mass basis ethanol has a Lv of 846 kJ/kg whereas water has a massive Lv of +2257 kJ/kJ, almost three times as much. The poor old ethanol has to use up most of its heat of combustion just to vaporise the water. No wonder it won't ignite with too much water present. What a fantastic EEI - and think of the fantastic calculations and equations you can impress your teacher with.
I wrote to Baine asking for some advice for Senior Chemsitry students planning to do this as an EEI. His reply is as follows (7 June 2016):
1) You might not be able to burn very high water percentages; I was forcing fuel and air into a very strong mixing environment with less ambient heat loss than you will have.
2) I essentially showed lower temperatures at the same heat rate with high water content. So as you go up in water content the energy is basically transitioning from a higher temperature over a smaller mass to a lower temperature over a larger mass (water is now included). I would expect the heat transfer effectiveness into your calorimeter will change as this transition occurs, and that that heat transfer would be dominated by the temperature 'into' the calorimeter. My guess is the increased mass flow per unit fuel energy will not be captured well, if at all, and you will see a difference with different water content.
3) I don't know exactly what 'controls' the fuel feed rate in a wick. Is it the consumption rate by the flame or capillary action? Or most likely some combination of the two. You might end up supplying a similar mass or volume of fuel for your different cases, but that would give you much less fuel energy for the high water cases. If this expectation is correct then it should magnify the differences you see with different fuels. In my experiment I was able to overcome this since I was controlling fuel flow. You can at least track the fuel flow by measuring the weight of your burner+fuel. The mass-to-fuel-energy calculations ought to be good for a high schooler and add some time to your study.
Baine Breaux
He provided me with a revised and more concise version of his thesis. It can be downloaded here (with permission).
I took the following photos in the lab at Moreton Bay College. They show the flames for various mixtures of ethanol and water. E80/W20 means 80% ethanol, 20% water by volume. I have tried to crop the photos so they are all the same size so you can see the relative heights of the flames. The E55/W45 burnt for about 10 seconds and went out. The colours are different, perhaps indicating different temperatures.
E100/W0
E80/W20
E70/W30
E60/W40
E55/W45
E50/W50
If you are doing this investigation as an EEI you'd need to work out how to state the ΔH, as per g , or per mol . If per mol, then per mole of what - ethanol, water? And what should you control: time of burning, temperature rise? What a conundrum.
Soybean biofuel manufacture and testing at school
Over 90% of Australia's transportation energy is supplied by petroleum-based fuel. Enormous amounts of (non-renewable) diesel fuel is consumed annually. Alternatively, biodiesel has fuel properties similar to petro-diesel and can be used directly in a diesel engine. Biodiesel improves lubricity and reduces toxic emissions during combustion. This suggests a good EEI.
Your research question could be along the lines of which vegetable oil produces the best biodiesel in comparison to commercial biodiesel? You can make biodiesel from soybean by the following method: Weigh accurately 20.0 g of soybean oil into a round bottom flask and add a few boiling chips, 6 mL of methanol and 1.2 grams of potassium carbonate and reflux for 25 minutes (at a low intensity). Then allow to cool. Add 18 mL of 1 M acetic acid to the flask and pour it all into a separating funnel. Allow the layers of the reaction mixture to separate overnight. Drain the lower glycerol layer into a waste beaker and collect the upper layer containing biodiesel into a tared beaker. Record the mass of collected biodiesel.
Testing: place a weighed amount (about 20 g) in a clean spirit burner and set it alight (as in the two suggestions about heat of combustion above). Let the heat be absorbed into a metal can containing a small amount (say 100 g accurately weighed) of tap water. Do the usual measurements to calculate the heat of combustion per gram, and the total energy in the sample of biodiesel. There's a good article: "Soybean Oil: Powering a High School Investigation of Biodiesel" by Paul De La Rosa, Katherine A. Azurin, and Michael F. Z. Page. See attached "soybean-diesel-supporting-info" .
Heat of combustion of
isomers and
polyols 4 H9 OH
(butan-1-ol, butan-2-ol and 2-methyl-propan-2-ol) - all commonly available at
school - would their ΔHc values vary and why? Do they have different boiling points - and what does this
tell you about intermolecular forces. Is this related to
ΔHc 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 97°C
BP 82°C
Sugar syrup density
Sugar solutions or "syrups" are used extensively in canning, and sometimes referred to as "heavy" or "light" syrups. These terms refer to their density in grams per millilitre. Water is 1 g/mL whereas a concentrated (heavy) sugar syrup might be 1.4 g/mL. But syrups are also made up as percent by weight (%w/w).
Just how are density and percent by weight related? This might seem like a trivial question but it is important in industry because sometimes the syrup is weighed out and sometimes measure by volume. And you have to be able to move from one to the other. For instance, if you mix 100mL of sugar with 100 mL water you don’t get 200 mL of solution. The 100 mL of sugar has a mass of 160 g so you end up with 160g sugar in 260g of solution (= 62.5% w/w). The density works out to be 1.3069 g/mL.
Glucose powder
Glucose syrup
An interesting investigation that could make a good EEI is finding the relationship between density and weight%. This has been done for sucrose and the relationship is a cubic polynomial (see my graph below). The question is: does glucose have the same relationship? The method seems quite straight-forward: weigh out accurately 50 g glucose and add 50.0g water. Mix. This is a 50%w/w solution. Tare a 10 mL measuring cylinder and add the syrup up to the 10 mL mark. Note the mass and calculate the density. I'm sure you could work out a more accurate method. Try other mixtures: eg 20% up to 70%w/w. Plot (make sure you also measure the mass of 10 mL of water (ie 0% glucose) so you can plot that point too. For extra information see the article by Karen Peterson, Department of Chemistry, San Diego State University, California "Measuring the Density of a Sugar Solution" in Journal of Chemical Education, Vol. 85 No. 8 August 2008, pp1089-1090.
The results of my trials with sucrose - a cubic polynomial as expected.
An after dinner drink called a Pousse Cafe is made from bottom to top, red grenadine, yellow chartreuse and green chartreuse. This can be simulated by coloring sugar water of different concentrations. It demonstrates how less dense liquids float on more dense liquids, if they're kept from mixing by carefully and slowing pouring them.
Nickel remediation
The most common use of nickel is for electroplating. Vast amounts of the metal are used in nickel plating and waste streams of aqueous nickel sulfate have to be dealt with. Besides being an expensive metal, nickel ions are toxic in the environment and have serious health hazards. Most commonly, the nickel (II) ions are precipitated as the hydroxide. Recovery of nickel can be right up to 100%. This suggests a great EEI, partly because the method of recovery can influence the yield.
I contacted Professor Bridget Trogden from Mercer University, Georgia, USA, who said that the mere timing of filtration was a factor affecting yield. Here's my suggestion: weigh out accurately about 1g of nickel sulfate hexahydrate (blue crystals, Mr 282.85 g/mol), dissolve in 100 mL water and add 8 mL 1.0M NaOH. You need to check the stoichiometry of this reaction to see that you have an excess of the hydroxide. Stir it and then filter immediately through a weighed filter paper. Air dry and then oven dry and weigh. I'd be doing duplicates or triplicates.
However, to assess the effect of delaying the filtering, I'd repeat the above but leave it for a week before filtering. If you have the resources, you could do another one but filter after say 4 days. You may be able to graph something here. I've attached the method from Professor Trogden's article in J Chem Ed, V88(2) Feb 2011, pp192-194 here .
She says that there are risks with using nickel compounds but should be safe for high school. The experiment contains no chemical hazards. Nickel compounds are considered suspected cancer agents via inhalation, but exposure is not expected in this experiment. Additionally, nickel hydroxide substances do not cause chemical burns or other problems that harsher hydroxides can. A 2005 J. Chem. Educ. Letter addresses these concerns and points out appropriate literature [Bentley, Anne K. et al. J. Chem. Educ. 2005, 82, 1775]. You'd do a risk assessment anyway. If it seems too dangerous you could try copper sulfate. Nickel sulfate hexahydrate (Sigma Aldrich) costs about A$37 for 100g.
Investigating the Impact of Ocean Acidification on Calcareous Organisms
The increase in ocean acidity since preindustrial times may have deleterious consequences for marine organisms, particularly those with calcium carbonate structures. The intergovernmental Panel on Climate Change (IPCC 2013) states that "carbon dioxide concentrations have increased by 40% since pre-industrial times, primarily from fossil fuel emissions and secondarily from net land use change emissions. The ocean has absorbed about 30% of the emitted anthropogenic carbon dioxide, causing ocean acidification".1
A good EEI would be to take some calcite to represent the calcareous organisms. Calcite is the most common form of the mineral whereas marble is the metamorphic form. Really, I guess you could use either. To a weighed sample you could add some acidic buffer solution of various pHs and let it react for a set time (maybe 30 minutes). Dry and reweigh. If the pH drops during the reaction add more buffer. What does the graph look like? Is pH that critical? Perhaps you should do triplicates. What about temperature (ocean warming): you could try a change in temperature as a separate variable. I think you may be shocked. There is an interesting article titled "Laboratory Experiment Investigating the Impact of Ocean Acidification on Calcareous Organisms" by Alokya P. Perera and A. Bopegedera in the J. Chem. Educ. 2014, 91, 1951−1953.
Here's my rough attempt at ocean acidification.
1. http://www.climate2013.org/images/uploads/WGI_AR5_SPM_final.pdf (accessed October 2015).
Corrosion
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 .
Oxidation of steel wool
If you leave steel out in the air it will rust as the iron reacts with oxygen and water in the air. If you exclude water the iron will still react with molecular oxygen thus: 4Fe(s) + 3O2 (g) → 2Fe2 O3 (s)
To get the steel wool perfectly clean you wash it in acetone, dry it, wash it in vinegar, and then rinse it off with water and dry it. Then quickly put it in a dry test-tube.See here for a method.
As a reminder, let's imagine it could be 1st, 2nd, or 3rd order:
So, you'd plot all three forms of the graph to try and linearise the relationship, and if one gives you a straight line then bingo - that's it.
My graph of oxygen pressure vs time.
I tried the natural log of oxygen pressure vs time and it was linear.
Corrosion of iron using spectrophotometry
Corrosion of iron in salt water is a popular context for an EEI. It is topical, it is of great economic importance, and it makes a great experiment. A common way of determining the corrosion rate is to measure the weight loss of an iron nail after a wee in salt water. However, to get reliable results, a balance with a precision of 0.01 mg is required since the weight loss is quite small over the week. And who has a spare $10000 to buy such a balance? But here’s a good spectrophotometric method that can be done with some simple chemicals and a phone with a colour analyser app.
In essence, sandpaper a nail with a wet 1200-grit sandpaper and degreasing it with a dip in acetone. Air dry it and immerse in 100 mL of a solution containing 3.5% NaCl, 1% ascorbic acid, and 0.25% 1,10-phenanthroline and begin timing. The color of the solution turns orange as time goes by. This is because the 1,10-phenanthroline,4,5 molecule reacts with iron to form an orange tris(1,10-phenanthroline)iron(II) complex ion. The concentration of this ion is directly dependent upon iron concentration.
Fe2+ (aq) 3 phen → [Fe(phen)]3 2+ .
The phenanthroline presents no chemical hazards and is available in a 5g bottle from Chem-Supply (Australia).
You can carry out the reaction on a magnetic stirrer if the iron nail sample is magnetic, and itself will function as a stir bar. Take aliquots (samples) of the solution every 5 minutes, transfer into a test tube or plastic cuvette, and capture the image with a mobile phone. I tried both Color Picker (for iOS on an iPad or iPhone), and Color Grab for Android. They both worked well. Of course you’ll have to make up a set of standard solutions of iron (II) ions with 1,10-phenanthroline added and measure their spectra.
A good source of Fe(II) is ammonium iron (II) sulfate hexahydrate and a standard solution of about 1 g/L is good. Then make up a range from 0.1 to 8 mg/L in 25 mL volumetric flasks for the standard curve. If you use the RGB values from the phone, you can calculate the absorbance. In essence, you subtract the blue value from the green value for the sample. You divide this by the B minus G values for the water blank. Then take the negative log (10) of the result. The value used by scientists is somewhat different (J. Chem. Educ. 2015, 92, 1696−1699, Supplementary Material). For example: A = -log((0.76Gs+0.5Bs)/(0.76Gb+0.5Bb) where Gs is the Green reading for the sample and Gb is the green reading for the blank. Likewise for blue (B). The value for R is not used as it is assumed to be zero
My thanks to Edgar Moraes, Mario Confessor and Luiz Gasparotto, Federal University of Rio Grande do Norte, Brazil.
pH and photosynthesis 2 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. You could also try light intensity. If you don't have a "luxmeter" to measure intensity you could take advantage of the fact that as you double the distance of the light source to the plant, the intensity is quartered (but you'd have to cut out daylight). There are a lot of variables to control and complex
biochemical reactions to examine.
Chlorine loss in a swimming pool - due to sunlight intensity 2 or 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 hypochlorite 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 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 transmitted (passed) by each type of cellophane; this is called their "λTmax"
(lambda T max), that
is, the wavelength most transmitted. 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 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 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 (λ
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.
Test kit
label. Solution type.
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 3 ) and urea (NH2 )2 CO
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). 2+ ) 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 21°C. They found that a Cu2+ concentration of 7.9 x 10-7 M (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 •5H2 O). 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 •5H2 O 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 (= 6227°C);
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).
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). 4 solution (5 g
copper sulfate pentahydrate, CuSO4 •5H2 O 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 4 •5H2 O - 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 (KSP ) of
copper carbonate at 25°C
is 1.4 x 10-10 . The equation KSP = [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 10-10 precipitation 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 KSP 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 KSP 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 NaCO3 solutions are mixed a precipitate will form if the ionic product is less than KSP .
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.
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 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 •nH2 O,
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 -
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 4 )3 •18H2 O).
The name 'alum' is a bit confusing as there is also
a double sulfate of potassium and aluminium with the
formula KAl(SO4 )2 ·12H2 O)
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 •18H2 O,
sometime called octadecahydrate (Mr = 666.4). The other
common one is •14H2 O.
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 •18H2 O
(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 by the scattering of light #1
HCl + sodium thiosulfate → sodium chloride + sulfur dioxide + sulfur + water.
2HCl(aq) + Na2 S2 O3 (aq) → 2NaCl(aq) + SO2 (g) + S(s) + H2 O(l) Equation 1
This reaction rate depends on the concentration of the two reactants and the temperature. Because the sulfur particles are so fine they form a cloudy whitish colloid. The concentration of the sulfur can be measured by the 'cloudiness' of the colloid mixture. The rate of this reaction can be measured by looking at the rate at which the product solid sulfur S(s) is formed. You may have carried out this reaction in Year 10 during the chemical reactions unit. It is usually carried out in a flask placed on a piece of white paper with a black cross on it - and you time how long it takes the cross to disappear. It is an old favourite. In effect, you are making use of the Tyndall effect (or Tyndall scattering), whereby light is scattered by particles in a colloid [see note below].
My setup. A beaker would work just as well.
I tried it again using a light meter. I made up a cell from four microscope slides epoxy glued together along their long edges to make a hollow box, and then glued another slide on the end (see photo above). I put a solution of 0.8 g of sodium thiosulfate in 50 mL water into my cell and shone a beam of light from a LED torch through the mixture and measured the intensity (transmittance) of the beam using a light sensor. I also tried a laser pointer (green, λ = 566 nm) and that worked well too.
To measure the transmitted light through the base of the cell, you could use a digital light meter that comes with the various laboratory data kits (eg LabQuest, DataMate, Spark). I tried an Arduino with a light dependent resistor (LDR) and it worked well. I then added one drop (0.05 mL) of 1M HCl to the cell, quickly stirred it and started a stopwatch. As time went by the solution became more and more cloudy and the transmittance decreased. I 'normalised' this value by expressing the transmittance as a fraction of the initial transmittance (Tn = T/Ti).
As the reaction proceeds less light can pass through the solution. Notice that it starts off slowly, speeds up and stops after 75 seconds. After that the sulfur starts to precipitate and the solution becomes clearer (so transmission increases).
The rate of the reaction is the slope of the Intensity vs time graph on the left. The rate is zero for the first 25 seconds and reaches a peak at about 50 seconds before returning to zero at 75 seconds. The negative rate is an anomaly (it just means the solution is getting clearer).
The reaction is said to be first order with respect to sodium thiosulfate concentration (rate ∝ [A]), where [A] is the concentration of the sodium thiosulfate. However, both the overall chemical equation and the mechanism for the decomposition of sodium thiosulfate are more complex than suggested by Equation 1 above. The reaction is acid-catalyzed, which means that the acid concentration must have some bearing on the rate in terms of producing an equilibrium concentration of HS2 O3 – ions. The HS2 O3 – ion is a reactive intermediate, reacting further with additional S2 O3 2– ions to produce polymeric ions containing multiple S atoms. When the chain of S atoms in a polymeric ion becomes long enough, it "closes" in on itself to form a ring of elemental sulfur (S8 ). That is complicated and the order is anyone's guess. Here are the steps:
S2 O3 2– + H+ → HS2 O3 – 3 – + nS2 O3 2– → H—S—(S)n —SO3 – + nSO3 2– n —SO3 – → H+ + – S—Sn —SO3 – – S—S7 —SO3 – → S8 + SO3 2–
NOTE: For a dispersion of particles to qualify for the Rayleigh formula, the particle sizes need to be below roughly 40 nanometres, and the particles may be individual molecules. Colloidal particles are bigger, and are in the rough vicinity of the size of a wavelength of light. Tyndall scattering, i.e. colloidal particle scattering, is much more intense than Rayleigh scattering due to the bigger particle sizes involved.
Reaction rate by the scattering of light #2 Further to the above discussion, you could also look at the light scattered at right angles to the path of the beam. As shown above when light is passed through a suspension of sulfur it is attenuated . That is, the transmitted light is not as intense as the incident light. This happens for small colloidal particles in suspension such as sulfur. The process shown in the EEI above is called 'turbidimetry': it is the measurement of the degree of attenuation of a radiant beam incident on particles suspended in a medium, the measurement being made in the directly transmitted beam.
But light is also scattered in all directions as it passes through the suspension. You can also measure the 'scattered light' by a technique called 'nephelometry': the measurement of the light scattered by suspended particles, the measurement usually being made perpendicularly to the incident beam.
Turbidimetry or nephelometry may be useful for the measurement of precipitates formed by the interaction of very dilute solutions of reagents (great for this Chemistry EEI), or other particulate matter, such as suspensions of bacterial cells (Biology EEI).
A fabulous chemistry EEI would be to examine the light scattered (by Tyndall scattering) through the sides of the glass cell. You could even do it in a beaker. If you look at the photo in the EEI above (#1) you will see that I also sticky taped a light dependent resistor (LDR) on the side of a cell, as well as the one underneath it. I connected the wires to the analog inputs of an Arduino board (each in series with a 1000 ohm resistor) and used a 'sketch' that read the outputs across the LDRs. It is very simple and worked well. The circuit diagram and the sketch is here if you want it.
In the graph below I have plotted the scattered intensity against time. The scattered light will start at a low number and then increase as more sulfur is precipitated. It is like the inverse of the transmitted light graph (although there is no mathematical relationship between transmission and scattering).
A graph of 90° scattering for the thiosulfate reaction. The scattering increaases as sulfur is formed.
A copy of the graph from the EEI above showing the same experiment but with transmitted intensity plotted against time. It decreases as sulfur is formed and blocks the light.
Reaction rate by the scattering of light #3 - effect of concentration
If you are using a light meter then just record the the scattered intensity as a function of time (as above) as you try out different concentrations. I made up four solutions of sodium thiosulfate in 50 mL water (0.2g, 0.4g, 0.6g and 1.0g), added acid and measured scattered intensity. Here are my four graphs (left):
The blue graph is of the most concentrated solution (1.0 g/50mL) and it is the quickest. I noted the time taken to get to an arbitrary intensity reading (300) and plotted concentration vs time.
The concentration vs time graph showed an inverse relationship so I linearised it by plotting 1/concentration vs time, and got a straight line (above).
Reaction Rate of Glow Sticks
Different colours are available.
Safe for use in emergency rescue
situations as there is no risk of sparking any flammable gases that may be
present.
The glow stick contains a vial of peroxide and a fluorescent dye mixed with
diphenyl oxalate. By mixing the contents, a chemical reaction takes place. The
reaction releases energy that excites the dye, which releases photons of
light. The reaction rate (and hence the brightness or intensity) depends on the
concentrations of the two chemicals. Manufacturers can produce glow sticks
that either glow brightly for a short time or glow more dimly for a much
longer time. Different situations require different types. The rate also
depends on temperature: cold ones are not as bright but last longer.
An ORP-12 light dependent
resistor (diameter is about 8 mm). This one was $3.50 from Jaycar
Electronics.
The calibration curve for
an ORP-12 LDR. It is a log-log graph and not easy to read (click graph image
to see an enlarged version)
This suggests a great EEI. The hypotheses are obvious from the above, but
it is the measure of light intensity that needs care.
Judging brightness with your eye is difficult and not that accurate.
You may have a light
intensity probe available with some of the datalogger kits. If not, you could
use a light dependent resistor (LDR) such as the ORP-12 for which there are
calibration curves of intensity vs resistance (see below).
The key think is to get absolute darkness as you don't want ambient light
(room or sun) getting in.
I made a colorimeter by using a length of grey plastic tubing (eg electrical conduit) with an inside
diameter big enough to fit a test tube (eg 20mm). I cut a window in the side
the diameter of the LDR and glued the LDR in place and wrapped black tape
around it to stop light getting in. You'd be surprised how light makes its way
in through the glue. I connected the two leads to a multimeter and glued an
endcap on the bottom (you can buy them to fit or just use a metal screw-cap
from a wine bottle.
Putting the glowstick in the home-made
colorimeter tube.
Looking down the tube - quite intense.
After 15 minutes it was reading 2.14 kW
I let it run for almost 3 days and was
still getting a reading. However, after 55 hours it was black to my eye.
The biggest changes are in the first few hours.
It is quite possible that the total light energy produced for hot, warm,
cold light sticks is the same (bright for short time = less bright for longer
time). Without giving too much away, you could compare the areas under the
curves (what quantity is this?). A hint: the light intensity is high just
after activation, then exponentially decays (so I believe). What a great
experiment.
Reaction Rate and Surface Area - I
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 and react.
You could measure the sides of the chunks with a Vernier calliper and calculate
the surface areas. I get a diameter of 25.63 mm, thickness of 4.30 mm and a full surface area of 1305 mm2 . When split in halves I get a SA of 1461 mm2 .
However, the reaction time for a whole tablet (at 23.5°C) is about
63 s, halves 54 s, quarters 51 s, eighths 43 s. A completely crushed tablet
- a powder whose surface area is enormous - takes about 23 s. So it is not linear; that is, the reaction rate doesn't vary directly with surface area.
What's the surface area of a powdered tablet? It is impossible to measure this directly in a high school laboratory. However, based on an estimation of how fine the powder is ground up (in a mortar, with a pestle) you can state an approximate "Specific Surface Area" (SSA). This is expressed in square metres per gram. For example, coarse sand (0.5 - 1.0 mm diameter) has a SSA of 0.01 m2 /g; fine sand 0.06 m2 /g, and very fine sand 0.1 m2 /g. Once you get into the clays (diameter < 0.002 mm) the SSAs are huge; montmorillinite is about 800 m2 /g. So an Alka Seltzer of mass 3.278 g and ground "fine" has a total surface area of 0.200 m2 (which is a whopping 200000 mm2 ). Whew, that's big!
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. I get 1305 mm2.
Cut them up and calculate the initial surface area. For the quartered tablet above I get 1617 mm2.
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.
Reaction Rate and Surface Area - II
Schools often use soluble aspirin tablets as they are smaller and cheaper. Here are some results to get you thinking. These are the averages of 16 groups from my Year 8 class; mass of tablet 1.240 g, diameter 17.16 mm, thickness 3.70 mm, volume of water 50.0 mL, beaker 100 mL, gentle stirring. The graph to the right is drawn from research data on the reaction rate of powdered limestone (calcium carbonate) as a function of particle size (and hence surface area). The numbers on the lines show each of the four particle sizes used (in µm).
Reaction time for aspirin tablets as a function of surface area.
Tablet division
Surface Area (mm2 )
Time to react (s)
Full
100
63
Halves
163
54
Quarters
226
51
Eighths
352
43
Fine powder
74000
23
Reaction Rate and Surface Area - III
If you want to get away from fizzy tablets here are a few other surface area experiments you could try:
Cut a marble tile with a masonry blade
into several long strips
Make measurements with a Vernier
calliper
Break the strips into 2, 4, 6 or 8
pieces.
Let them react with acid.
Zinc and
acid Displacement
Reaction 4 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 •5H2 O
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.
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.
This is not an EEI suggestion but a way of measuring solution concentrations colorimetrically. It could be used for any solution whose colour intensity is a measure of its concentration, eg blue copper ions. It may be useful if you don't have access to a colorimeter (as above), or even a cheaper one that can be used with a data logger (such as Vernier connected to a LabQuest as shown below) then a smartphone app may be the answer.
I found it is quite simple to use Vernier Spectrovis spectrometer with a LabQuest data logger to get accurate absorbance readings. The cuvette and solution is towards the top right corner of the spectrometer. You can scan the visible spectrum quite easily to get a lambda max reading. My thanks to Maddie for helping me out.
To test out the smartphone app, I made up some standard copper ion solutions of 4.0 g of copper per 100 mL (0.6 M) using copper (II) nitrate (Cu(NO3 )2 .3H2 O). I then diluted this to make a series of dilutions from about 0.24 M up). The Color Grabber app on an Android, or Color Picker app for the iPhone work well and reports HSV values. The H is for Hue and reports it on a scale of 0-360° which can be used as an index of absorbance. The "S" is for Saturation in %
I used the Color Grabber app for Android. This photo shows the the solution in a cuvette. The reading is "Morning Glory" HSV (H 180°, S 23%, V 82%). It is the H value I used as an index of concentration. A good one for the iPhone is Color Picker.
Here's my standard curve using a standard copper nitrate solution with known dilutions. The line is a good fit (R2 = 0.9984) indicating agreement with the Beer-Lambert Law. Using this curve I can read off the concentration of my unknown solutions.
Solubility of salts in water and alcohol
A study of salt solubility in different solvents is very important for many industrial applications. More particularly, a knowledge of accurate solubilities is needed for the design of separation processes such as extractive crystallization or for the safe operation of different processing units such as distillation columns, absorption units, and extraction plants.
Aqueous electrolyte solubility is generally available for many salts, but aqueous-organic mixed solvents data is very scarce, obsolete, or not available at all. An EEI based around this sounds like a good idea. You could try salts NaCl, KCl, and NaBr and solvents methanol and ethanol. To avoid water salt contamination, salts could be dried at 100°C in a drying stove for at least 2 days before use.
Evaporating the water is best done in an evaporating basin with a watchglass on top to prevent splatter.
I like to do the final drying in an oven at 120°C because it splatters too much on a hotplate.
All you need do is to prepare a saturated solution at a desired temperature (with undissolved solid on the bottom) and keep it stirred for an hour or two. Let settle and take a sample of the supernatant liquid, weigh it, and evaporate the water (on a hotplate to dryness and then in an oven at 120°C for a day until mass is constant). Might take a day or two.
Then you could try binary mixtures of solvents. Have a look at these results*:
Note the linear relationship up to 50°C then it is almost constant after that. Below this transition temperature, the solid phase is NaBr.2H2 O, and above it, the solid phase is NaBr. Well, so they say.
Note the inverse salt solubility temperature dependence. This unusual effect of the temperature on the solubility may be explained both by Debye-Hückel and ion association Bjerrum theories because at such low ionic concentrations the electrostatic interactions are more pronounced and association is favoured.
*Click here for a good article: "Solubility of NaCl, NaBr, and KCl in Water, Methanol, Ethanol, and Their Mixed Solvents" in the Journal of Chemical and Engineering Data , Vol. 50, No. 1, 2005, p29-32. The authors are Simão P. Pinho and Eugénia A. Macedo, Laboratory of Separation and Reaction Engineering, Departamento de Engenharia Química, Faculdade de Engenharia, Rua do Dr. Roberto Frias, 4200-465 Porto, Portugal. Thanks guys.
Cetyl alcohol and water evaporation losses
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 )15 OH - 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).
Sometimes it may
take a while to get so cetyl-stearyl alcohol would be fine (you can often get
this from places that sell home-made soap suplies). It is merely a mixture of
cetyl alcohol (C16) and stearyl alcohol (C18).
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. It obviously is not water soluble so you would need to
make a solution by using a different solvent (try ethanol). 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
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 20°C
to 80°C 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 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
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.
Lactic
acid and the fermentation of milk 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. Start with heat.
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 Ka 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. As a rough guide, one of my students measured the titratable acidity of milk as 0.0288M as lactic acid (7.10 mL titre) at the start, and on day 7 the value was 0.0815M (15.65 mL titre). The in-between values did not give a linear graph; it was much more exciting than that.
Note about identifying variables :controlled variable or independent variable (or both) in this experiment (and others that involve collecting data
over a period of time). 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
Anode Reaction: Pb(s) + HSO4 - (aq) → PbSO4 (s)
+ H+ (aq) + 2e− 2 (s) + HSO4 - (aq) + 3H+ (aq)
+ 2e− → PbSO4 (s) + 2H2 O(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) 3 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 (NaHCO3 and
Na2 CO3 ) 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: 2 CO3 present in solution will react with the acid
as the titration proceeds until it is all converted to
NaHCO3 (aq):2 CO3 (aq) + HCl (aq)
→
NaHCO3 (aq) + NaCl (aq)
Bicarbonate titration: 3 (aq) + HCl (aq) → NaCl (aq) + CO2 (g) + H2 O
(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.
solution of sodium bicarbonate.
Breakdown of Antacid Tablets
Instead of investigating the thermal breakdown of sodium bicarbonate in baking soda (as above), you could consider the effect of heat on antacid tablets. These tablets are designed to neutralise stomach acidity so they are basic salts, usually calcium carbonate (eg Quick-eze or TUMS ). The process and analysis would be the same as above. You could try Mylanta tablets (magnesium hydroxide and aluminium hydroxide) but Gaviscon may be tricky (aluminium/magnesium trisilicate). Remember, in an EEI you are not just trying to simulate real life; in this case you are trying to extend the understanding of carbonate breakdown to a bigger range of temperatures than you would find at home or in a car glove-box. That's why you would use the lab oven to get high temperatures for the trials.
TUMS - 500mg CaCO3 per tablet
Soft Drinks and Tooth Decay
A soft drink such as lemonade or Coca-Cola is a drink that does not contain alcohol, as opposed to a hard drink, which does. Australians consume about 300 mL of soft drink per day on average but amongst 14-16 year olds the figures are 1000 mL for males and 500mL for females. Soft drinks are about 10% sugar so a young male typically consumes 27 teaspoons of sugar per day in soft drink; a girl, about half that. Just one can of soft drink has about 10 teaspoons of sugar in it. The resultant obesity (fat) epidemic is attributed in part to soft drinks. Health risks from over consumption include diabetes, kidney stones, obesity, osteoporosis, and tooth destruction.
Tooth decay is partly from the bacteria feeding on the sugar but also from the acids reacting with the tooth enamel. The citric or phosphoric acid in soft drink dissolves the calcium out of the enamel leaving a softened matrix for bacteria to enter the teeth and cause wholesale carious (tooth) destruction. So drinking sugar-free (diet) soft drink is not the answer.
A good EEI may be to look at the effect of drink acids on teeth. Teeth are a form of hydroxyapatite Ca5 (PO4 )3 (OH) but you can simulate this in the lab with calcium carbonate (marble chips). The problem is: you need to control the type of acid, whether it is phosphoric acid as found in cola drinks, or citric acid as found in lemonade. A study by Fraunhofer and Rodgers (2004) found that the rate of enamel dissolution of teeth was not dependent on pH but may be affected by titratable acidity. Remember that weak acids (phosphoric, citric) are not fully dissociated in water (so their pH is not that low) but they gradually release more hydrogen ions as they react. The "titratable acidity" will be a measure of this. Citric acid, for instance, is a tribasic acid which releases its H+ ions in four steps. It has a reversible reaction with CaCO3 and the reaction is controlled by diffusion of reaction products away from the 'tooth' surface; thus, consider keeping it stirred.
The study by Fraunhofer & Rodgers (2004) found that Coke was not as aggressive as lemonade in enamel dissolution. In fact, Mountain Dew had the greatest impact of all.
As a trial, I took a 1 cm3 cube of marble and placed it in 200 mL of 7.5%w/v citric acid solution (pH 1.8) at 50°C (with stirring) and after 60 minutes it had lost 0.30 g. Why not make up a synthetic soft drink from phosphoric or citric acid. What concentration will you choose? How does the reaction rate or the extent of the reaction vary with concentration? Does temperature have much effect on the rate? Does the product - calcium citrate or calcium phosphate - impede the progress of the reaction; that is, how soluble are the products (one is 4 times as soluble as the other). How do you measure the progress of the reaction (amount of carbonate consumed or change in titratable acidity of the solution)? Oh, the possibilities are endless. And you can drink the left-over Coke and rot your teeth a bit more at the same time. A perfect EEI.
CCA Treated Timber Leaching
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
Iron filings in fortified cereal
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
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
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)?
Distinct ethanol water layer before mixing (thanks to Bek and Eliza)
Adjusted to 0.0 mL
After mixing it reads 1.20 mL
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
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
Browning of apples
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 . Laboratories use a reflectance spectrophotometer at a wavength of 400 nm to measure the degree of browning but it is unlikely you'd have access to one of those. If you have a regular (transmission) spectrophotometer you could take a sample of your browned apple, blend it up, filter it and measure absorption at 400nm. You could compare various treatments (providing the surface area of the same was controlled). Failing this, perhaps let an apple go brown and take a photo and make that your official "brown" (or make a range of photos on a scale of 1-10 from 'not brown' through to 'very brown".
Effect of catalyst concentration on reaction rate: enzymes
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 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
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 a or Kb ) gives a measure of this dissociation.
The quantitative behaviour of acids and bases in solution can only be understood
if their Ka (or pKa ) 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 pKa 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 pKa 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 CH3 COOH, and then measure it's pH, you can
calculate the Ka 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
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° 2+ (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?
Haybale fires - one for farmers
Over the warmer months Rural Fire Services throughout Australia are called to extinguish fires in stored hay as a result spontaneous combustion. These fires are usually the result of a combination of storing hay that is too moist and warmer temperatures. These fires can lead to the loss of valuable feed and stored machinery as well. The phenomena of exploding haystacks has been with us for as long as we have been making hay. Pliny, the Roman Philosopher wrote in 60BC "When the grass is cut it should be turned towards the sun and must never be stacked until it is quite dry. If this last precaution is not carefully taken a kind of vapour will be seen arising from the rick in the morning, and as soon as the sun is up it will ignite to a certainty, and so be consumed". Old microbiology books often contain anecdotal evidence of haystack explosions, one from 1939 states "The stack that sets itself on fire does so in a curious way dependent at first upon both moisture and microorganisms. A really dry stack of hay won't heat spontaneously; a really damp stack can't be set fire to".
Even though hay bales may seem dry, they are actually quite moist. Hay that is considered 'dry' has up to 20% moisture content. Although hay should be stored at about 15% humidity for the best conservation, it is often baled at much higher moisture levels. There are many reasons for this: moist hay has less leaf loss (major nutrients are in the leaves), sun drying reduces its quality, and, most importantly, climatic conditions often do not allow adequate drying.
When hay is bailed and the plant material is either too green or has excess moisture (over 30%, as a result of rain, dew, flood water, etc) it continues to respire, generating heat and water which emerges as a vapour through the leaf pores. Within the confines of the bale, the water condenses and spreads by capillary action. This initiates and promotes fungal and bacterial growth in the bale. These micro organisms produce heat and along with higher external temperatures and humidity, and can reach a peak of 70°C at about 5-7 days. If the relative humidity in the middle of the stack is below 95% then the microorganisms become inactive and the temperature of the stack drops. If the relative humidity in the middle of the stack is above 97% then the resultant heat of vaporisation of the water dissipates the heat rapidly and the temperature of the stack drops. This explains why very wet silage does not explode. However if the narrow window of 95%-97% relative humidity is obtained then the microorganisms continue to produce heat, which cannot escape, which raises the temperature. This temperature rise accelerates the chemical oxidation of the hay releasing more heat quickly raising the temperature within the bale to ignition point (200-280°C) where it will burst into flames.
It all depends on the size and shape of the bale, the moisture content and how tightly it is packed. Even if it doesn't catch on fire the presence of micro-organisms can be a health problem: a disease known as Farmer's lung disease - a pulmonary affection caused by an allergic reaction to the inhalation of these thermophilic actinomycetes and molds. Farmer's lung typically produces shortness of breath, cough, and fever. If not adequately treated, this disease can lead to severe and irreversible lung damage. This suggests a great EEI and one with great social and economic importance.
A traditional method of monitoring hay stack combustion potential used by farmers has been to poke a section of steel rod into bales in the hay stack and leave for several hours before removing it and feeling for any heat build-up. But for an EEI you have to do better than that.
A truck carrying a load of hay bales caught fire on the southeast outskirts of Winnipeg (Canada). CBC News Manitoba October 3, 2010.
However, if you are using a standard rectangular bale you can just push a thermometer in to it. For a big bale you would need to get a length of metal pipe with holes drilled in it and bash it into the bale - and then lower a thermometer (attached to a string) gently down the pipe until it is centered.
The most common bales used in Australia are the two string bales, which are approximately 900mm long x 450 mm deep x 350 mm high (16 kg). A prime shredded lucerne hay bale costs about $13.
Strength of plastic
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
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
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
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
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 2 CO3 ,
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 Na2 HPO4 and 0.1M NaH2 PO4 .
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?
Electroplating
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 2 B4 O7 .10H2 O
(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 2 .
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
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
Ion
exchange resin
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 4 , to Portland cement prolongs the hardening. The most
important compounds present in cement are: 3CaO•Al2 O3 ,
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 H2 SiO3 .
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
The health of our river
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
Polarisation of light in acidified sugar solution
C12 H22 O11 (sucrose)
+ H2 O + H+ => C6 H12 O6 (fructose)
+ C6 H12 O6 (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.
Acidification of seawater by carbon dioxide [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
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
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
Corrosion of
roofing steel
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 BlueScope zincalume better than ordinary galvanising?
Paper chromatography of leaves
Diffusion of aqueous ions 3 )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 4 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 (K2 MnO4 )
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
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 •2H2 O or 1 g MgSO4 •7H2 O 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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