how can 2g cast iron be dissolved in sulphuric acid?
457094
It will be very much easier to dissolve the cast iron in concentrated nitric acid (HNO3).  I would suggest somewhere in the amount of 15 to 20 mL of concentrated HNO3 to completly dissolve all of the iron.  Just remember to NOT inhale the fumes as they may be toxic.  Note:  Sulferic acid (H2SO4) will actually dissolve organics such as plastic polymers.
293072
You could probably just dehydrate it with concentrated sulfuric acid and heat.\n\nAnother method would be to generate a alkyl halide with thionyl chloride (SOCl2) and then eliminate it in a E2 reaction with sodium ethoxide in ethanol (especially strong base to be sure you get elimination, and not SN1)
863022
Here you are:\n\nFe+2 + 1/2O2 (g) + H2O (l)  ---> Fe (OH)2 (s)\n\nFe (OH)2 (s) + 1/2O2 (g) ---> Fe2O3 (s) + 2H2O (l)
595775
Previously cast iron was used. Now its Aluminium. Work is going on a Ti-Mg alloy.
853690
Heating the sand with a strong solution of Hydrochloric acid will remove most of the impurities:\n\nFe2O3 + 6HCl --> 2FeCl3 + 3H2O\n\nAl2O3 + 6HCl --> 2AlCl3 + 3H2O\n\nThe chlorides formed can then be removed by rinsing with warm water.\n\nSilica does not react.
952998
The reductive dissolution of goethite ( -FeOOH) and hematite ( -Fe 2 O 3 ) by phenolic reductants has been examined in order to improve the understanding of iron transformations in soils, sediments and aquifers. Rates of goethite reductive dissolution by hydroquinone increased as the pH was increased from pH 1.8 to 4.65, arid the following reaction stoichiometry was obeyed: 2 - FeOOH + QH 2 = 2 Fe 2+ + Q + 4 OH - . As the pH was increased from pH 4.5 to 6.0, the reductive dissolution rate decreased to below the detection limit. At pH 3.4, the reductive dissolution of hematite was two orders of magnitude slower than goethite. The relationship between structure and reactivity was examined for a series of mono-, di-, and tri-hydroxybenzene reductants. Rates of reductive dissolution decreased in the following order: catechol ~- hydroquinone 〉 3,4-dihydroxybenzoic acid 〉 resorcinol-phenol-4-hydroxybenzoic acid.
822936
Look here:\n\nhttp://jewelrymaking.about.com/od/solderinginformation/a/032806.htm\n\nand here:\n\nhttp://alchemy.chem.uwm.edu/amalgamator/NCW/ncw2001/results.html\n\nto get the following:\n\nThe pickle, used in the soldering metal process is a sodium bisulfate based solution. This pickle is particularly good for gold. Other metals use different pickles. For example, "Copper and copper alloys can be pickled in sulfuric acid solutions with the addition of an oxidizing agent (for heavy scales) such as chromic acid, nitric acid, ferric sulfate or sodium dichromate. NOTE: Chromic acid and dichromate require treatment for hexavalent chromium. Chromic acid or dichromate solutions tend to passivate the substrate and require activation prior to plating. Activation is commonly carried out in 5-10% sulfuric acid solution or 10-20% hydrochloric acid solution. If the latter is used to interchangeably pickle steel parts and copper parts, the dissolved copper must be closely controlled to avoid a typical problem. The dissolved copper will eventually plate out on the steel parts and cause adhesion problems."\n\nI hope this helps.  Be careful.
798471
Assuming a pure substance, (this method will not work on mixtures).  \n1. Dissolve in Carbon disulphide.  (The compound is now banned in schools)\n2. Allow to evaporate in a fume cupboard.\n3. The presence of yellow rhombic like crystals is a test for S.\n4. Melt a sample and do a MP test.\n5. Melt a few grammes and allow to cool in air.  Needle like crystals will be formed.
566520
Fe(s)+H2SO4(aq)excess--> Fe2+(aq) + SO42-(aq)+H2(g)\n\nratio is 1:1\n\n5gFe/55.85g/molFe= .090mol\n\nand\n\n1 molar solution =1 mole per liter\n\nso,\n\n.09mol(1 liter/1mol)= .09L or 90mls
320284
It is not possible.  Observe the chemical equation:\nAg2SO4(s) + AgNO3(s) --> Ag2SO4(s) + AgNO3(s)\n\nIn order for a substance to dissolve, it must be in an aqueous solution.
508593
Like Kak says - remove the iron filings with a magnet. Sieve the iron filings to seperate the filings from any chalk and copper sulfate that is clinging to them. Wash the copper sulfate and chalk through the sieve with distilled water.\n\nThe copper sulfate is soluble in water, but the chalk isn't, so add water, stir until the copper sulfate dissolves, then let the copper sulfate solution filter through filter paper into a beaker below. Wash any remaining copper sulfate solution through the filter paper with distilled water.\n\nBoil the copper sulfate solution until all the water evaporates, leaving solid copper sulfate residue.\n\nThe damp chalk powder sediment is left in the filter paper. Scrape it off the filter paper into another beaker. Boil off the water until you are left with a dry chalk powder residue.
655901
You really need to give a little more information about what you are attempting to convert.  However, in very general terms you remove the sulfate, usually by making a solution containing your material basic.  You then form the hydrochloride salt by adjusting the pH back to the acid side via the addition of hydrochloric acid.\n\nPlease note, this process applies to the salt conversion of materials that are basic in nature and water soluable as their salt form.
843269
Use the Nickel Metal..Put your steel into a electrolysis solution of Nickle  electrolytes and charge it with electric anode and cathodes.
378767
You would complete dissolve it in strong acids and then separate the metals by qualitative analysis techniques.  Not simple and very expensive.
103867
NaClO + 2FeSO4 + H2O -> NaCl + 2 Fe(3+) + 2 (SO4)(2-) + 2 OH(1-)
1029447
sulfuric acid is h2so4. convert 1600 kg to g so times by 1000. then take that number in grams and times by the molar mass of of sulfur (gSx1mol S/gS) then convert to moles of H2SO4 by using your balanced chemical equation(coefficients), with moles of S on bottom, moles of H2SO4 on top. everything cancels and your left with moles of sulfuric acid that you could make. if you need to take it another step to grams then you would then just use the molar mass of sulfuric acid.
654500
Many of them oxidize H2S\n\nso its like water but using sulphur, thats why a lot of these bacteria are purple. \n\nheres an abstract on how they work, or at least will give you a general idea ans some examples!\n\nSeven strains of heterotrophic iron-oxidizing acidophilic bacteria were examined to determine their abilities to promote oxidative dissolution of pyrite (FeS2) when they were grown in pure cultures and in mixed cultures with sulfur-oxidizing Thiobacillus spp. Only one of the isolates (strain T-24) oxidized pyrite when it was grown in pyrite-basal salts medium. However, when pyrite-containing cultures were supplemented with 0.02% (wt/vol) yeast extract, most of the isolates oxidized pyrite, and one (strain T-24) promoted rates of mineral dissolution similar to the rates observed with the iron-oxidizing autotroph Thiobacillus ferrooxidans. Pyrite oxidation by another isolate (strain T-21) occurred in cultures containing between 0.005 and 0.05% (wt/vol) yeast extract but was completely inhibited in cultures containing 0.5% yeast extract. Ferrous iron was also needed for mineral dissolution by the iron-oxidizing heterotrophs, indicating that these organisms oxidize pyrite via the "indirect" mechanism. Mixed cultures of three isolates (strains T-21, T-23, and T-24) and the sulfur-oxidizing autotroph Thiobacillus thiooxidans promoted pyrite dissolution; since neither strains T-21 and T-23 nor T. thiooxidans could oxidize this mineral in yeast extract-free media, this was a novel example of bacterial synergism. Mixed cultures of strains T-21 and T-23 and the sulfur-oxidizing mixotroph Thiobacillus acidophilus also oxidized pyrite but to a lesser extent than did mixed cultures containing T. thiooxidans. Pyrite leaching by strain T-23 grown in an organic compound-rich medium and incubated either shaken or unshaken was also assessed. The potential environmental significance of iron-oxidizing heterotrophs in accelerating pyrite oxidation is discussed.