Practical investigations
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Reactions and analysis involving iron(II) and iron(III) compounds
A large number of FAQs have now been received on this topic and published on this site. To help you find the information you want more easily, we have identified the main areas of interest and collected the relevant FAQs together below, under the following headings:
1. Iron(II) sulphate: preparation and reactions
2. Preparing iron(II) solutions
3. Analysis of iron tablets
4. Extracting iron from foodstuffs
5. Analysing mixtures of iron(II) and iron(III)
6. Determination of iron(II) using manganate(VII) (permanganate)
7. Titration calculations
8. Colorimetry involving iron complexes
9. Aspirin determination using iron(III)
10. Other titrations involving iron(II) or (III)
11. Miscellaneous reactions of iron and iron compounds
Use these headings to go to the area of interest to you. Some FAQs appear under more than one heading.
1. Iron(II) sulphate: preparation and reactions
Preparation of Mohr's salt, (NH4)2SO4.FeSO4.6H2O, and the thermal decomposition of hydrated iron(II) sulphate form part of at least one current A-level chemistry course.
What is hydrated iron(II) sulphate and what is it commonly used for?
What happens when iron(II) sulphate is heated on its own? Please give the reaction.
What is the chemical formula for hydrated iron ammonium sulphate?
Pages 13-14 in the Nuffield Student Book describe the preparation of a hydrated salt, ammonium iron(II) sulphate from iron filings, ammonia solution and sulphuric acid. Why is the actual yield of crystals likely to be much lower than the theoretical amount?
Please help me with the reactions concerning the preparation of ammonium iron(III) sulphate dodecahydrate
How can you prove that hydrated iron(II) sulphate contains some water of crystallisation? I thought it might be similar to copper sulphate which changes colour when it is heated. If it is, could you let me know the colour change for iron(II) sulphate?
2. Preparing iron(II) solutions
Many of you have asked about the need for the addition of acid when making up solutions of iron(II) and iron(III) salts.
How can I investigate whether iron(II) sulphate tablets get significantly oxidised depending on whether the tablets are crushed or whole?
I need to find the percentage iron, in cast iron, using a dichromate titration. What is the best way to dissolve the iron to form iron(II) without it oxidising to iron(III)?
Why is Mohr's salt preferred over iron(II) sulphate for titration purposes?
When making Mohr's salt in Topic 1, why is sulphuric acid added to the iron in excess? Wouldn't it be better to have excess metal and then filter?
I have 2 questions:
1) Why is it that potassium manganate(VII) solution is chosen to oxidise Fe(II) in the redox titration?
2) Why is iron(II) ammonium sulphate a relatively stable iron (II)compound? What is a double salt?
Im doing an experiment on the analysis of iron tablets. I havent done titration experiments before so Im lost. I need to know why the iron tablets should not be heated more than necessary when dissolving in sulphuric acid, and why they are being dissolved in sulphuric acid and not water.
We carried out a redox titration experiment to determine the percentage of iron in a sample of ammonium iron(II) sulphate. When we were dissolving the ammonium iron(II) sulphate crystals in sulphuric acid (approx 1 mol/dm3) we were told not to heat it to assist dissolving. Why was this?
Why are some iron(II) salts less suitable for use as a primary standard in volumetric work than others?
I am researching the iron content of different green vegetables. I am puzzled as to why you need to boil the vegetable in sulphuric acid before you titrate with KMnO4. I have searched text books and the internet and found nothing. Please help.
I am trying to find the percentage purity of cast iron using a redox titration with potassium manganate(VII). I plan to first dissolve the iron in 1 molar sulphuric acid but how do you determine what amount of cast iron to crush and dissolve if you do not know its formula mass?
I am planning an investigation to compare the iron content of various breakfast cereals and am having some difficulty finding out the best way to prepare the cereal in order to be able to titrate it using potassium manganate(VII) solution. What would you recommend?
I am investigating the iron content of broccoli for my advanced higher chemistry project. I know that the vegetables must be boiled in sulphuric acid, but I have no idea how much acid to use or for how long to boil the broccoli. I feel that if I do not know these basic facts my results will be inaccurate.
I am doing an A2 Chemistry planning exercise to determine the % purity of iron in a sample of cast iron by a redox titration. I know the cast iron should first be dissolved in sulphuric acid, but I do not know how to prevent the further oxidation of ferrous sulphate, or how to determine whether all the iron has all dissolved. Is it when there are no more gas bubbles? Do I need to write it as a complex ion, Fe(H2O)6SO4, when I write the equation of it with KMnO4? If not, then how should I write it? Thank you very much.
I am doing an A2 practical planning to find the % mass of iron in iron cast by redox titration. How do I go about it? Do I react the iron cast with conc. H2SO4? How can I make sure all the iron are in the form of Fe2+ before I titrate it with potassium manganate? To titrate with MnO4, one of the solution has to be acidified, so do I add H2SO4 again? Should it be concentrated or dilute? Thank you!
The question is to plan an investigation about the percentage purity of iron in iron cast by redox titration and I think I have to use potassium manganate in the presence of acid but I'm not sure what the formula is for cast iron can you help?
When you are carrying out an experiment to find the iron content of a solution using potassium manganate(VII), why do you have to boil the vegetables in acid and make this solution up to 100 cm3 with sulphuric acid before performing the titration? Why cant you just use water? And why, after working out the iron content of the vegetables, is the value obtained way above the daily iron requirement?
3. Investigating iron tablets
There are various methods for determining the iron content of 'iron tablets'. One student asks if pure iron(II) sulphate can be used instead of iron tablets for some experiments.
Im doing an experiment on the analysis of iron tablets. I havent done titration experiments before so Im lost. I need to know why the iron tablets should not be heated more than necessary when dissolving in sulphuric acid, and why they are being dissolved in sulphuric acid and not water.
What is hydrated iron(II) sulphate and what is it commonly used for?
I am doing a practical investigation into what affects the oxidation of Fe2+ ions to Fe3+ ions in iron tablets. I want to vary temperature and pH, but I'm really stuck on how to do it. And also is it possible to find the molarity of Fe2+ in the tablets and then use the appropriate amount of iron(II) sulphate powder instead of using the tablets each time? I am using a titration with potassium manganate(VII).
I am doing my A2 practical investigation for OCR Salters on the amount of iron in iron tablets. I have decided to use three different methods and compare different brands of iron tablets. I have already identified 2 methods - titration and colorimetry. However, I cannot think of the third and last one. Also what are the three best brands of iron tablets to use and compare as I have heard that not all iron tablets work in the titration experiment?
4. Extracting iron from foodstuffs
Questions about getting the iron out of spinach, bananas, meat and cereals.
Hi, I am planning an investigation into the iron content in organic and inorganic forms.
1) Will vitamin C release Fe2+ ions from the oxalate.
2) In what form/state is Fe found in vegetables- how can I find this out?
Will it affect whether I use sulphuric acid in which to dissolve it? Will it affect whether I can use KMnO4 to indicate its presence during a titration?
My investigation for Advanced Higher chemistry is the analysis of iron content in foods, from organic vegetables to meats rich in iron. I found and researched two procedures to carry out using potassium manganate and sodium thiocyanate. I have my fingers crossed that these will work. My problem is that no one has had experience of doing this sort of experiment so I am wondering if you have any suggestion of other experiments I could try to calculate the iron content. It would great help!
I'm investigating the concentration of iron in bananas. After heating a sample in concentrated sulphuric acid, clearing it with hydrogen peroxide I've to then adjust the pH with sodium hydroxide before adding hydroxyammonium chloride followed by phenanthroline reagent. What is the function of the hydroxyammonium chloride in this experiment?
I am planning an investigation to compare the iron content of various breakfast cereals and am having some difficulty finding out the best way to prepare the cereal in order to be able to titrate it using potassium manganate(VII) solution. What would you recommend?
I am investigating the iron content of broccoli for my advanced higher chemistry project. I know that the vegetables must be boiled in sulphuric acid, but I have no idea how much acid to use or for how long to boil the broccoli. I feel that if I do not know these basic facts my results will be inaccurate.
When you are carrying out an experiment to find the iron content of a solution using potassium manganate(VII), why do you have to boil the vegetables in acid and make this solution up to 100 cm3 with sulphuric acid before performing the titration? Why cant you just use water? And why, after working out the iron content of the vegetables, is the value obtained way above the daily iron requirement?
I am researching the iron content of different green vegetables. I am puzzled as to why you need to boil the vegetable in sulphuric acid before you titrate with KMnO4. I have searched text books and the internet and found nothing. Please help.
5. Analysing mixtures of iron(II) and iron(III)
When both iron(II) and iron(III) ions are present in solution, the iron(III) ions first have to be reduced to iron(II) using zinc. The total iron content and iron(II) content of the original mixture can be determined by titration.
'Estimation of iron(II) and iron(III) in a mixture containing both' is one of my pieces I have to investigate. The thing is, I haven't got a clue on how to do it! Please would you be able to give me a helping hand and tell me how it all works and how I would go about doing it? All I get told is that I am provided with 200 cm3 of a solution containing 1.1 to 1.3 g of iron ions as a mixture of the two. It has present at least 30% by mass of the two ions. I know have to give detailed method on how to work out the amount present, but I am unable to do so, so please could you help?
I'm doing a practical of trying to find out the percentage by mass of iron(II) and (III) in the same solution of 200cm3. I know the procedure but need to know some suitable concentrations and volumes of solutions and zinc to use through the experiment.
2Fe(3+) + Zn(2+) --> 2Fe(2+) + Zn
I am carrying out the above experiment.
The known volume of Iron is 100 cm3 at 0.08 mol dm-3.
Zinc must be in excess but how do I find the minimum mass of zinc required in order to justify my choice of 5g?
What is the overall equation when (acidified) iron (III) is reduced with granulated zinc?
I need to be able to carry out a volumetric procedure so that I can estimate the amount of iron(II) and iron(III) in a mixture containing both. I have looked at the recently asked questions and there is a similar one, however, this person seems to understand what they are doing - I dont even know where to start! Please help.
6. Determination of Fe2+ by titration with manganate(VII)
This titration makes use of the redox reaction in which manganate(VII) ions, in the presence of acid, oxidise iron(II) ions quantitatively to iron(III). The two half-reactions involved are:
Oxidation: Fe2+(aq) --> Fe3+(aq) + e-
Reduction: MnO4-(aq) + 8H+(aq) +5e- --> Mn2+(aq) + 4H2O(l)
Combining these two half-reactions so that the loss and gain of electrons cancels out, gives the balanced ionic equation for the reaction.
No indicator is needed as the titration is a self-indicating one, with the appearance of a pink colour when the last drop of manganate(VII) solution is added at the end-point.
I have 2 questions:
1) Why is it that potassium manganate(VII) solution is chosen to oxidise Fe(II) in the redox titration?
2) Why is iron(II) ammonium sulphate a relatively stable iron (II)compound? What is a double salt?
I am confused with the method of using colorimetry to determine the amount of iron(II) present in iron tablets. I will carry out a titration with potassium manganate(VII). however, I do not know what to do with the solution once it turns pink. Do I keep it for colorimetry or not? Also, how do I convert iron(II) into iron(III)? What would be a suitable filter to add for colorimetry? With what solution e.g. sodium thiocyanate? Please can you also help me what would be suitable quantities of whatever solution to add and the concentration required if any?
I am doing a Chemistry prac investigating the amount of Fe(II) present in iron tablets. I am investigating this using a titration with potassium permanganate, after adding distilled water and 2M suphuric acid to the ground iron tablet in the titration flask. According to the amount of iron that is given on the bottle, the titre should be about 3 ml. However, I cannot reach an end-point because the permanent pink colour that indicates the end point of the reaction is never reached. No matter how much potassium permanganate I add, the colour fades out (sometimes this takes overnight). The colour still fades out even when titres of 100 and 200 ml are added. I was wondering what may have occurred to cause this unforeseen consequence. Thanks heaps.
How do I standardize a solution of potassium manganate(VII) by an iron(II) salt (ammonium iron(II) sulphate)?
Hi. I am using a technique called a back titration in order to determine the amount of iron(II) in my iron tablets. I was advised to titrate the iron(II) (iron(II) is already mixed with sulphuric acid) with excess potassium manganate(VII), and then to react the left over potassium manganate(VII) with hydrogen peroxide. How would I determine the amount of hydrogen peroxide to react with the excess potassium manganate(VII) and the concentration of iron(II)?
I need to find the percentage by mass of iron in a sample of cast iron, with the method involving a redox titration. Im going to use KMnO4 and will dissolve the cast iron in the acid so that it converts into Fe2+ ions. I also understand how to do the calculations to work out the percentage by mass. However, Im not sure of how to work out the inital concs and initial mass - do I just randomly pick a number or is there actually a way of working out how much you should use? thanks a lot!
I am doing a planning exercise to determine the % mass of iron in cast iron by redox titration. Its all ok apart from one section where I have to use calculations to justify the amounts/concentrations used, or just to show what is a suitable amount/concentration.
In my calculations to determine the mass of iron and then the % mass of iron, I have used 0.04 molar KMnO4 and 0.50g of cast iron. I chose these values because 0.50g of iron seemed a sensible amount, and 0.04 molar KMnO4 is a value I found in a question about a similar reaction.
I'm assuming the classic equations such as No. of Moles = Mass/Mr and NOM in solution = volume x concentration must be used, but where do I start?
I have recently done a practical experiment and had quite a large error margin with my results. I dissolved iron(II) sulphate crystals in 50 ml of sulphuric acid, all mixed with distilled water to get a final volume of 250 ml. 25 ml of this solution was then added to 20 ml of sulphuric acid, and titrated using a burette, with potassium manganate(VII) (permanganate). Can you help me with finding where the errors for the potassium permanganate titration volume errors might have come from?
I'm working out some chemistry results from a iron sulphate titration how I can tell if my iron sulphate in my iron tablets is anhydrous or hydrated ... what should I be looking for?
Where can I find out about the reactions between potassium manganate(VII) and iron(II) oxalate. This is for an experiment investigation the iron(II) content of different vegetables using a titration with patassium manganate(VII).
7. Titration calculations
To tackle titration calculations succesfully you must be confident in using the relationship which links moles in solution to the volume and concentration used of the solution used:
moles in solution = volume (in cm3)/1000 x concentration (in mol dm-3)
It is important before you start hitting calculator buttons to see the link between the titre value obtained and the final answer required in terms of a series of simple steps. For example, in the manganate(VII) - Fe2+ titration , these steps would be:
From titration:
--> Volume of MnO4- titre
--> Moles of MnO4- used (see above formula)
--> Moles of Fe2+ in pipette sample (e.g. 20 cm3), using balanced equation
--> Total number of moles of Fe2+ in solution made up (e.g 200 cm3)
--> Mass of Fe2+ in solution
--> Mass of Fe or FeSO4 in original tablet or sample.
I would like to know how to calculate the number of moles of H2O of FeSO4.xH20 from titration with KMnO4. I titrated 25 cm3 of dissolved hydrated iron(ll) sulphate with added 20 cm3 1 mol dm-3 of H2SO4, with KMnO4. Mean titre= 22 cm3. The mass of the crystals which were dissolved = 3 g
Five iron tablets with combined mas of 0.9 g were dissolved in acid and made up to 100 cm3 of solution. In a titration 10 cm3 of solution reacted exactly with 10.4 cm3 of 0.01 mol dm-3 potassium permanganate, What is the percentage by mass of iron in the tablets?
I need to find out the % purity of iron in cast iron. I have found out how to do the calculations but it says that you have to work out the number of moles in potassium manganate. To do this I understand that you need the relative molecular mass of the potassium permanganate. Is that going to be KMnO4 or just MnO4-, and when adding up all the numbers because the oxygen is O4 does that mean that 4 is added onto the value of 16 because it has gained four electrons or does it stay at 16. Help would be appreciated. Thanks.
I am investigating the percentage purity of iron in cast iron. I know the method of titrating it against potassium manganate but I am a bit unsure about the calculations. Please can I have some guidance?
I have recently done a practical experiment and had quite a large error margin with my results. I dissolved iron(II) sulphate crystals in 50 ml of sulphuric acid, all mixed with distilled water to get a final volume of 250 ml. 25 ml of this solution was then added to 20 ml of sulphuric acid, and titrated using a burette, with potassium manganate(VII) (permanganate). Can you help me with finding where the errors for the potassium permanganate titration volume errors might have come from?
8. Colorimetry involving iron
The concentration of Fe2+ and Fe3+ in solution can be determined colorimetrically by converting them into more intensely coloured complex ions. First you will need to select an appropriate filter (depending on the colour of the complex) and produce a calibration curve for your colorimeter using solutions of the iron complex over a range of concentrations.
With Fe3+ a deep red complex is formed with thiocyanate (CNS-) ions, which should be present in large excess. With Fe2+ a complex with 1,10-phenanthroline can be used, but more usually at this level Fe2+ is first oxidised to Fe3+, using manganate(VII) - see previous section - or hydrogen peroxide, in the presence of acid. The Fe3+ formed is then determined as the thiocyanate complex.
Another coloured complex, formed between Fe3+ and salicylic acid (2-hydroxybenzoic acid) is used in the colorimetric determination of Aspirin - see next section.
Hello, I am carrying out my investigation into the purity of synthesised aspirin. I encountered a problem when setting up my calibration curve using iron (III) chloride. I came across another method that suggested making a buffer solution of HCl and KCl with the iron (III) chloride. When I did this it worked and a calibration curve could be drawn. Can you tell me why the buffer solution was used and the difference between it and just the normal iron (III) chloride? Does it interact differently with the 2-hydroxybenzoic acid? Thanks in advance :)
I am confused with the method of using colorimetry to determine the amount of iron(II) present in iron tablets. I will carry out a titration with potassium manganate(VII). however, I do not know what to do with the solution once it turns pink. Do I keep it for colorimetry or not? Also, how do I convert iron(II) into iron(III)? What would be a suitable filter to add for colorimetry? With what solution e.g. sodium thiocyanate? Please can you also help me what would be suitable quantities of whatever solution to add and the concentration required if any?
Hi, Im currently doing an investigation concerned with iron tablets. I want to carry out colorimetric analysis of ferrous ions and ferric ions separately in iron tablets. So far, to make the calibration graph for ferrous ions, Ive decided to use ammonium iron(II) sulphate and potassium ferrocyanide. For the ferric ion calibration graph, Ive decided to use ammonium iron(III) with potassium thiocyanate . However Im not too sure what concentrations of the chosen chemicals to use, to make up the needed calibration graphs.
I'm currently doing my advanced higher chemistry project, and have run into a few difficulties. I'm trying to determine the iron content of an iron tablet by a colorimetric method. I've tried converting the Fe2+ to Fe3+ using potassium permanganate the adding a drop of KSCN, but the mauve of the excess KMnO4 gets in the way. I'm now trying H2O2 to oxidise the Fe2+ as its colourless but I'm really confused about concentrations and methods. Please help!
I am carrying out a colorimetric analysis into the amount of iron in spinach after it has been cooked for varying amounts of time. I need to make an iron(III) solution in order to produce my calibration curve. It has been suggested that I use Fe(NO3)3.9H2O but will need to standardise it as it is not a primary standard. How do I standardise an iron(III) solution or is there a better Fe(III) compound to use for my standard?
I am doing a practical to find the concentration of iron(III) ions in a sample of river water using sodium thiocyanate. I do not know the standard procedure for this experiment. Thanks.
I want to determine % of iron in iron tablets using colorimetric method for my project, and compare it with normal titration. I am not sure how to make solutions and work out ppm.
9. Aspirin determination using iron(III)
Aspirin (2-ethanoyloxybenzoic acid) is synthesised form salicylic acid (2-hydroxybenzoic acid. Samples of aspirin made in a school laboratory often contain some salicylic acid as an impurity. Salicylic acid (or the salicylate anion) is produced when aspirin is hydrolysed by acid (or alkali) - see the revised React tutorial on Aspirin Investigations
The purple complex formed by salicylic acid or salicylate and Fe3+ ions is used to determine the concentration of this impurity in the aspirin sample, or after hydrolysis, the concentration of aspirin itself.
Once again a suitable calibration curve for the colorimeter used must be determined first.
Hello, I am carrying out my investigation into the purity of synthesised aspirin. I encountered a problem when setting up my calibration curve using iron (III) chloride. I came across another method that suggested making a buffer solution of HCl and KCl with the iron (III) chloride. When I did this it worked and a calibration curve could be drawn. Can you tell me why the buffer solution was used and the difference between it and just the normal iron (III) chloride? Does it interact differently with the 2-hydroxybenzoic acid? Thanks in advance :)
Hi, I'm really stuck. I'm doing my individual investigation on methods for testing the purity of aspirin. I did colorimetry with copper(II) nitrate. I got the method from school. I have been trying to find out about the Cu-aspirin complex and how the method works, but I can only find information on colorimetry with iron(III) chloride. Is my method correct? And what does the Cu-aspirin complex look like?
I am confused with my aspirin investigation. For the acid-base titration and colorimetry I have dissolved my aspirin tablet in ethanol. I thought this was a good idea as aspirin is not very soluble in water. However, I am now confused as to whether the ethanol affects the structure of the aspirin. Is it converted to salicylic acid?
I have also carried out iron (III) colorimetry. However, I think I have done this wrong because I used a variety of concentrations of salicylic acid to plot a calibration curve. I then dissolved the aspirin tablet in ethanol and added the iron(III) chloride. This formed a coloured complex, but I did not expect this as the aspirin would have needed to have a phenol group to form the purple complex.
Any help would be appreciated.
I've done an investigation on the purity of aspirin and completed a colorimetry procedure with ammonium iron(III)sulphate. If you could help by outlining the calculation steps for deducing the purity of the aspirin from salicylic acid that would be great!
I was wondering if you could explain why iron(III) NITRATE is able to comples salicylic acid. I understand that transition metals are able to form complexes but this specific reaction is hard to understand.
Im investigating the purity of aspirin using colorimetry. I know how to set up a calibration curve with iron(III) chloride. However once I have hydrolysed my aspirin do I have to add a certain amount of iron(III) chloride to this in order to compare it with my calibration curve?
I am planning my A2 chemistry coursework, and I have decided to investigate the different methods used to determine the purity of aspirin. The sheet suggests the use of colorimetry and iron(III) chloride. I am familiar with colorimetry, but don't know why I need to use iron(III) chloride in colorimetry nor do I know its relation to aspirin. If you could tell me how I relate it to colorimetry and/or aspirin, it would be a great starting point for me.
10. Other titrations involving iron(II) or (III)
How do I find the percentage mass of iron (III) when reacting it with iodine followed by titrating the iodine with sodium thiosulphate?
I'm investigating the concentration of iron in bananas. After heating a sample in concentrated sulphuric acid, clearing it with hydrogen peroxide I've to then adjust the pH with sodium hydroxide before adding hydroxyammonium chloride followed by phenanthroline reagent. What is the function of the hydroxyammonium chloride in this experiment?
I am trying to find the exact concentration of a saturated aqueous solution of chlorine by back titration, and it has been suggested to use an Iron(II) compound as the reducing agent. However I am unsure as which suitable iron(II) compound to use (can I use any?), and how to determine its excess in only knowing the appoximate concentration of the chlorine (7g dm-3).
I am doing a back titration to find out the concentration of an aqueous solution of chlorine using ammonium iron(II) sulphate. Should the ammonium iron(II) sulphate be acidifed or does that have no bearing on the results?
I am doing an investigation on the analysis of wine. I have looked at an article from Chemistry Review (Jan 98) by Derek Denby and in it he talks about how to find out what the alcohol content of wine is. One method is to oxidise ethanol using an excess of potassium dichromate and then doing a back titration using ammonium iron(II) sulphate. I really need an indicator for this experiment or to know when the end point is for the indicator sodium diphenylamine sulphonate. I haven't had much luck with the internet so I thought you could help!
11. Miscellaneous reactions of iron and iron compounds
I am doing my chem investigation on the percentage of copper in coins, and some of the coins contain Fe3+ ions. I know that excess ammonia is needed to remove the iron(III) ions, but please can you give me an equation that shows this reaction between the iron and the ammonia?
I'm doing an experiment on how to identify a drug as aspirin, paracetamol or ibuprofen by reacting them with neutral iron (III) chloride, noting colours and then heating it and noting any other colour changes. Ibuprofen doesn't have an OH group in it, so it doesn't change colour. However, paracetamol does, and so does aspirin after heating.
I'm having trouble explaining how the Fe3+ ion reacts with the enol group - what exactly does it do to make it change colour? (I've read the chemguide tutorial and didn't understand a word of it ... ) Help much appreciated!
Im revising benzene at the moment and I don't understand how iron catalyses the reaction between methylbenzene and bromine. Hope you can help. Thanks.
In a practical I added dilute sulphuric acid to aqueous potassium manganate(VII) followed by aqueous iron(II) sulphate. Could you tell me what are the inferences in this practical? Is this a test for anion or cation?
When aqueous sodium hydroxide is added to a solution of iron(II) sulphate a dark green gelatinous precipitate is formed. Is the colour of the green precipitate changing to brown after hydrogen peroxide is added, due to oxidation of iron(II) ions to iron(III) ions?
What happens to cobalt(II) chloride, copper(II) nitrate, iron(II) sulphate, zinc carbonate and sodium nitrate, when they are all heated (separately)? Also, can you tell me what gases if any, evolve?
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updated: 14 March 2007
