Preparing Vanadium in Various Oxidisng States
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Originally, practical text books used zinc amalgam to reduce the vanadium(V) state to the beautiful lilac V(II) state. Mercury is not suitable in school (or any other) chemistry now and reducing with zinc granules is slow. Here, I have used zinc wire which has been inserted into 0.1M copper(II) sulfate solution to get a copper coating on the zinc which acts as a catalyst. By moving the wire from one little test to another you can put 3 wires Io the tube reducing to V(II). Once the lilac colour is there, you need to keep a wire in so that oxygen in the air does not reoxidise the V(II) to the green V(III).
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Electrolysis of solutions
The large scale procures used large carbon/graphite electrodes and produced so much chlorine that students were inhaling far too much chlorine which affected and hospitalised students with breathing issues such as asthma.
In this procedure only about 0.5ml of 0.5M copper chloride solution is used which will produce about 6ml of chlorine gas, much of which reacts with the potassium iodide and potassium bromide reagents and the moist blue litmus paper.
The electrodes are made of carbon fibre and are very robust.. They are available from online kite shops.
You can a see a video of this on http://tinyurl.com/pjcsge4.
In this procedure only about 0.5ml of 0.5M copper chloride solution is used which will produce about 6ml of chlorine gas, much of which reacts with the potassium iodide and potassium bromide reagents and the moist blue litmus paper.
The electrodes are made of carbon fibre and are very robust.. They are available from online kite shops.
You can a see a video of this on http://tinyurl.com/pjcsge4.
If the electrolysis is stopped and the clips carefully removed and replaced by clips attached to a multimeter set on 20V, a reading of about 1V is obtained. Now there is a fuel cell which is the reverse reaction to the electrolysis.
Cu2+ + 2e− → Cu . + 0.340
Cl2 + 2e− → 2Cl− .+ 1.358
No doubt chlrone gas is absorbed in the carbon fibre electrode
I wonder if that can be carried out on other systems?
Cu2+ + 2e− → Cu . + 0.340
Cl2 + 2e− → 2Cl− .+ 1.358
No doubt chlrone gas is absorbed in the carbon fibre electrode
I wonder if that can be carried out on other systems?
Galvanic cell?: that is not the apparatus in the book
One of the comments passed to me about microscale experiments is “It’s not the apparatus or method they see in exam questions and in the text book. So students will be at a disadvantage if they carry out these versions of experiments” “Even if they get better results?” I even had one teacher who said yes to that but usually the conversation alters course at that point.
Here is a case in point. Traditional pictures (scroll down for the gallery) of galvanic cells usually have 2 beakers overfull with 1M solutions with salt bridge, usually a gel in a glass U-tube between them. The salt bridge takes a long time to make and then all that is required friom the instructions is to measure a voltage. Applications such a titrations and illustrating the Nernst Equation flitter in and out of the UK A-level chemisry syllabus
I have done versions in small containers and well plates with electrolytic paper and cotton used as a salt bridge. Electrolytic means that a salt such as potassium nitrate of sodium sulphate is absorbed in the material..
However, microscale methods also involve using self contianed electrodes with porous plugs dipped into an electrolyte acting as the salt bridge (https://www.researchgate.net/publication/253768319_Small-Scale_and_Low-Cost_Electrodes_for/figures) from Per-Odd Eggen and or http://www.mijst.mju.ac.th/vol2/535-541.pdf). Plugs are made or cotton wool, electrolytic agar or electrolytic plaster. See https://www.youtube.com/watch?v=TxSoXxlcswc from Peter Schwarz in Germany. Kazuko Ogino (see microscale-exp.csj.jp/MCEexperiments_eng.html) works at a small level than me with some wonderful experiment..
Then there are the paper methods as seen in a still from a Flinn video. It works really well and over be done in 10 minutes. But it is not what students see in text books or on the exam paper.
I have now made a flexible bridge using electrolytic agar inserted into silicone tubing. The 0.7% agar solution with potassium nitrate added (amount does not appear to be over critical) is dissolved in a microwave (10 second bursts). When a clear solution is obtained (HOT), a syringe is attached to 4mm diameter silicone tubing and the liquid is draw up the tubing and then left for 20 minutes to cool and set. I made a 1 metre tube and cut pieces off as I wanted.
So here we are with equipment which looks like the text book or in questions set in the exam. No it will not convince some. Note how international this little piece is with items from Norway, Germany, Thailand, Japan and the USA a well as the UK.
Here is a case in point. Traditional pictures (scroll down for the gallery) of galvanic cells usually have 2 beakers overfull with 1M solutions with salt bridge, usually a gel in a glass U-tube between them. The salt bridge takes a long time to make and then all that is required friom the instructions is to measure a voltage. Applications such a titrations and illustrating the Nernst Equation flitter in and out of the UK A-level chemisry syllabus
I have done versions in small containers and well plates with electrolytic paper and cotton used as a salt bridge. Electrolytic means that a salt such as potassium nitrate of sodium sulphate is absorbed in the material..
However, microscale methods also involve using self contianed electrodes with porous plugs dipped into an electrolyte acting as the salt bridge (https://www.researchgate.net/publication/253768319_Small-Scale_and_Low-Cost_Electrodes_for/figures) from Per-Odd Eggen and or http://www.mijst.mju.ac.th/vol2/535-541.pdf). Plugs are made or cotton wool, electrolytic agar or electrolytic plaster. See https://www.youtube.com/watch?v=TxSoXxlcswc from Peter Schwarz in Germany. Kazuko Ogino (see microscale-exp.csj.jp/MCEexperiments_eng.html) works at a small level than me with some wonderful experiment..
Then there are the paper methods as seen in a still from a Flinn video. It works really well and over be done in 10 minutes. But it is not what students see in text books or on the exam paper.
I have now made a flexible bridge using electrolytic agar inserted into silicone tubing. The 0.7% agar solution with potassium nitrate added (amount does not appear to be over critical) is dissolved in a microwave (10 second bursts). When a clear solution is obtained (HOT), a syringe is attached to 4mm diameter silicone tubing and the liquid is draw up the tubing and then left for 20 minutes to cool and set. I made a 1 metre tube and cut pieces off as I wanted.
So here we are with equipment which looks like the text book or in questions set in the exam. No it will not convince some. Note how international this little piece is with items from Norway, Germany, Thailand, Japan and the USA a well as the UK.
The Bromate Bromide reaction: why do we use phenol?
This experiment which I did with students in the 70s and 80s is coming back and being used as an exemplar experiment by UK exam Boards. I used the method from Atherton and Lawrence’s book (Experimental Introduction to Reaction Kinetics (Concepts in Chemistry)) But I always wondered "why use phenol in the mix?" The idea is that In acid solution, bromide ions react with bromate ions to produce bromine which then reacts with phenol to give tribromophenol with hydrogen and bromide ions. The still-produced bromine then decolourises a dye such as methyl red or methyl orange and the colouless solution then turns slightly yellow as the bromine is still being produced. So I looked at J Chem Ed’s only entry on this kinetics experiment (J R Clarke, Volume 47, Number 11, November 1970) and his opening lines are “This laboratory exercise was developed by modification of a kinetic method for the determination of phenols, with the emphasis on the kinetic rather than on the analytical aspects.” In fact the rate was first measured in the 1890s without the use of phenol! |
The mix is made from 1 ml of methyl orange solution, 4 ml of water, 2 ml of 0.1M potassium bromide and 0.1M potassium bromate (needs to be freshly prepared). The clock is started with the addition of 2 ml of 0.5M sulfuric acid. These just happened to be available in the CLEAPSS lab.
Look no phenol and in 35s the dyes is decolourised. Video is on https://www.facebook.com/profile.php?id=100012329655433&ref=bookmarks. So I think the whole phenol idea comes from an application to analyse (and still is) the concentration of phenol in waste waters and in the kinetics it only serves as a distraction to the student. Could be used as an excellent environmental study though! How would the CLEAPSS colorimeter cope with this? First indications are good with using Arduino technology to follow the reaction. |
Preparation of copper sulfate crystals
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So why bother. Surely this reaction carried out by nearly every student in the UK for nearly 100 years should be OK large scale. Well there are really inefficient and unsafe parts to the procedure described in many text books. These include the following.
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Here is an interesting observation, when the copper(II)sulfate crystals appear immediately, they are rather long but when the crystallization is slower, one obtained the familiar lozenge shape.
Although cited in text books as a method of making salts, it is pretty unique, only ethanoic (acetic) acid works in the same way. There are problems with complex ions forming with other well-known acids. What about changing the metal. It does work with nickel oxide but the hazards now associated with nickel oxide do make it rather worrying. |
Rusting of iron nails in a Petri dish
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The reactions take place in a plastic Petri dish on a plastic file containing the instruction sheet. The experiments could also be carried out directly on the sheet.
The yellow solution is ferroxyl indicator which contains potassium hexacyanoferrate(III), phenolphthalein and sodium chloride. Potassium hexacyanoferrate(III) shows the presence of iron(III) ions by forming Prussian blue; these are the anode sites and indicate that iron is oxidising (ie, loosing electrons) to iron(II) ions on the surface and then to iron(III) ions (see circle 7 where the reagent is added direct to iron(III) nitrate solution) Phenolphthalein solution indicates the cathode sites where oxygen molecules dissolved in water are reduced (gain electrons) to hydroxide ions. The process is accelerated by the presence of an electrolyte, (sodium chloride) and becoming more alkaline. In circle 6, iron(II) hydroxide is naturally oxidising to iron(III) hydroxide (green to brown colour change). Phenolphthalein is mauve in alkaline solutions but there is a second colour change to colourless at high pH solution as shown in circle 8 in the third slide. In boxes 2,3 and 4, the effect of attaching a metal to the iron nail can be followed. Zinc and magnesium are used as sacrificial anodes to reduce the onset of corrosion. The effects can be seen in closeup by using a Veho usb microscope
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Gas chemistry in a Petri dish
The gas is generated in a reaction vessel which may be an empty blister pack from or a very small watch glass. The gas then diffuses within the dish to react with indicator papers, aqueous reagents or even flowers. The volume gas generated should be less than 10mL for a 9cm-diamter Petrie dish and 5mL for a 5cm-diameter Petri dish)
This means that toxic gases can be investigated because the levels used are well below the accepted workplace exposure limits. Much of the gas made is taken up undergoing reactions in the Petri dish.
An example using hydrogen sulphide can be found on this video
https://www.youtube.com/watch?v=klgUCRbxyMk
Below is an example of ammonia chemistry, showing reactions with damp narrow range indicator paper, acid solution with Universal Indicator and an array of metal salts. It is rather an over the top version! You can decide which solutions you can use.
This means that toxic gases can be investigated because the levels used are well below the accepted workplace exposure limits. Much of the gas made is taken up undergoing reactions in the Petri dish.
An example using hydrogen sulphide can be found on this video
https://www.youtube.com/watch?v=klgUCRbxyMk
Below is an example of ammonia chemistry, showing reactions with damp narrow range indicator paper, acid solution with Universal Indicator and an array of metal salts. It is rather an over the top version! You can decide which solutions you can use.
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1) 0.01M hydrochloric aid and Universal Indicator (situated at "one o'clock")
2) 0.1M copper(II) sulfate(VI) solution 3) 0.1M nickel(II) sulfate(VI) solution 4) 0.1M cobalt chloride(II) solution 5) 0.1M iron(III) nitrate solution 6) Freshly made iron(II) sulphate solution, about 0.1M 7) 0.1M zinc sulfate(VI) solution 8) 0.1M lead nitrate solution 9) 0.1M barium chloride solution 10) 0.1M magnesium sulfate(VI) solution 11) 0.1M manganese sulfate(VI) solution |
Other gases to study in this way are chlorine (sodium hypochlorite and hydrochloric acid), hydrogen sulfide (iron sulfide and dil hydrochloric acid), sulphur dioxide (sodium metabisulfite and hydrochloric acid) and the nitrogen oxides (sodium nitrite and hydrochloric acid).
Chemistry on a plastic sheet
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The instructions are inserted into a plastic sheet and the chemistry carried out on the sheets byt adding drops with either plastic pipettes or dropping bottles (see equipment). The time is spent in the preparation but disposal is usually just wiping off the solutions with a paper towel.
An example can be found on this video https://www.youtube.com/watch?v=sk3ZolhPyWM |
The formation of a precipitate in the centre of a drop is interesting because the solids are pushed in from each end of the drop. The solids dissolve, the particles diffuse through the water and meet to form a precipitate if the chemistry is suitable. An example can be found on this video https://www.youtube.com/watch?v=oizwWsm43lY .
How to explain this beautiful effect? Usually we show precipitation reactions by mixing 2 liquids together and a solid appears. Here we start with the solids. Dissolving splits the solid apart and there must be the formation of a less soluble solid when the line forms down the centre of the droplet. There has to be a fundamental understanding of chemical processes at a nano-level to explain this. The next title might help.
How to explain this beautiful effect? Usually we show precipitation reactions by mixing 2 liquids together and a solid appears. Here we start with the solids. Dissolving splits the solid apart and there must be the formation of a less soluble solid when the line forms down the centre of the droplet. There has to be a fundamental understanding of chemical processes at a nano-level to explain this. The next title might help.
Conductivity of solutions
The conductivity equipment (See equipment page) is quite sensitive. There are various electrodes you can use, carbon fibre, steel from a paper clip and copper wire from stripped electrical wire. In pure water (distilled, deionised or RO water) the LED will not light up but in tap water (middle picture) it lights up. Add one grain of salt to a droplet of pure water and, after stirring, with the electrodes, the LED light up. How can this possibly happen? You might notice bubbling at the electrodes which indicates chemical reactions at the electrodes are required to allow the current to flow. This brings in the whole notion of IONS which with atoms and molecules, completes the model that students need to picture chemical activities at the nano-level. The third picture show an arrangement that can b sue to show what happens when tap water with a little Universal Indicator is conducting electricity.
pH and Indicators
This beautiful array of colours is formed in a Combo plate. The buffers, in the large wells, are made by combining two halves of the components of a buffer, various ratios of drops which always add up to 20 (1mL). 1 mL of pure water is added and the pH can be measured with a Hanna Chequer pH meter. A few drops are then moved to small wells and indicator can added across the wells.