what is formed when solid copper sulfate is mixed with water?
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Solid copper sulfate is a bluish crystal that is highly water soluble. It would therefore dissociate to form copper (II) and sulfate ions when placed in water. It DOES NOT react with water.
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Cu2+ (aq)  +  Zn (s) -->  Cu (s) + Zn2+ (aq)
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The precipitate is a solid, and it usually is in the form of a salt (i.e. NaCl)  If you are given an equation that begins with two liquids, for example, and one of the end products is solid (s), then you know that that is a precipitate.  \n\nPrecipitation can occur for a few different reasons.  One reason can be due to the insolubily of one of the reactants and wil result in precipitation  This can also be due to supersaturation of the mixture that causes the insoluble material to precipitate out of solution.\n\nWikipedia has a nice example using Silver Chloride (AgCl).
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2Cu + O2 ---> 2CuO,          \nThis is copper oxide or cupric oxide or copper( II )oxide\n\n4 Cu + )2---> 2Cu2O\nCopper also combines with oxygen as Cu2O, cuprous oxide , or copper(I) oxide
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1)\nSodium Sulfate (Na2SO4) will break into ions in solution,\nNa2SO4 (s) --H2O--> 2 Na+1 (aq) + SO4-1 (aq)\nFor every 1 mole of Na2SO4 dissolved, 3 moles of ions are formed.  If the concentration of Na2SO4 is 1.15 mol/Liter, then the ion concentration must be 3 times as much, 3.45 mol/Liter.\n\n2)\nTlCl3 is an ionic compound which when dissolved in water will break into ions since Chloride compounds are generally soluble,\nTlCl3 (s) --H2O--> Tl+3 (aq) + 3Cl- (aq)\n\nC2H5OH, otherwise known as ethanol, will easily dissolve (easily mincible) in water, but will not break into ions (as it is not an ionic compound).\n\nFe2(CO3)3 (although it is an ionic compound) in considered insoluble in water.  Carbonates in generally are considered insoluble unless they are bonded with a group 1 element or the Ammonium ion.\n\n3)\n.356 mol/Liter HCl means that for every 1 Liter of solution, there are .356 moles of HCl dissolved in it.\nMolarity  = moles / Liter --> moles  = Molarity * Liters\nSo if you originally had a 6.5 M solution with a volume of .0572 Liters, you originally had .3718 moles of HCl.\n\n.356 M = .3718 moles HCl / x Liters\nx Liters = .3718 moles / .1.044 Liters.\n\nSo if then old volume was 57.2 mL and the new volume should be 1044 mL, then 986.8 mL shoudl be added to solution.
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d) basic\n\n[Non-metal oxides do the opposite and form acidic solutions if they are soluble e.g. sulphur dioxide dissolves to form sulphurous acid)
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Copper reacts with carbon dioxide in the presence of moisture to form copper carbonate, which is green.  More precisely this is a cupric carbonate/cupric hydroxide complex:  CuCO3/Cu(OH)2.  Likely some cuprous carbonate is also present Cu2CO3.
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Zinc is a much more reactive metal than Copper.  Zinc is much higher on the activity series than Copper.\n\nWhen Zinc metal is added to a Copper +2 solution, the Zinc will displace the Copper in solution, precipiating out Copper metal while Zinc ions go into solution.\n\nZn (s) + Cu+2 (aq) --> Zn+2 (aq) + Cu (s)\n\nSince Zinc (II) Sulfate is soluble in water, the Sulfate (SO4-2) ion is a spectator ion in the reaction, it is in solution on both sides of the reaction.
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Once again, you can follow my directions for making anhydroyus copper sulfate in your previous post. Anhydrous copper sulfate is white or a very pale blue-green if you didn't keep it dry enough, which is as close to colourless as you are likely to get. Water of course is colourless.\n\nThe chemicals themselves are fairly easy to find.\n\nIf you have access to the school's chemical supply you should be able to get some phenolphthalein. Add a few drops of baking soda and Bob's your Uncle.
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The salt that is formed is Barium sulphate, which simply has low solubility in water (or acid or bases - you need a chelating agent for the Barium to dissolve it again). There is  no other significant reaction (like redox or altering pH).
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CuSO4•5H20\nThe dot notation denotes the hydrate.
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Wait...not mixed with H2O? As in, molten liquid ammonia and molten liquid sodium sulfate? In that case no reaction. But with water, then:\n\n2NH3 + Na2SO4 + 2H2O --> (NH4)2SO4 + 2NaOH
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Anhydrous copper sulfate and water would be the easiest two. \n\nYou can buy copper sulfate at any good garden supply shop and most larger hardware stores. The copper sulfate usually sold is hydrated, and is already blue in colour.\n\nYou can make it anhydrous by putting it the oven, uncovered on a sheet of tin foil at 120oC or hotter. You will know when it has become anhydrous beause it goes white. This is quite safe to do in your kitchen oven, all you are doing is driving off the water. No nasty fumes or residues are produced.\n\nStore the crystals in an airtight jar to stop them absorbing water from the air. When you are ready add some water to a few crystals they will retrun to their original blue colour.\n\nThe other possibility is to manufacture your own cabbage indicator. That will make a bluish purple colour when mixed with baking soda. I've included a link with instruction on how to do this.\n\n______\n\nErrata.\n\nThat's what you get for working from memory. Copper sulfate will only partly dehydrate at 120oc. You will need to set your oven to 160oC or higher to create anhydrous copper sulfate.
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This lab involves putting iron (steel wool) into a solution of copper sulphate. The reaction produces metallic copper and iron sulphate solution. I've never been really sure if it's iron(II) sulphate or iron(III) sulphate as the product. I've used both equations with good results. \n\n   Fe   +  CuSO4   --->  Cu  +  FeSO4\n   2 Fe   +  3 CuSO4   --->  3 Cu  +  Fe2(SO4)3\n\n\nFor prelab, make up a saturated copper(II) sulphate solution\nTrust me. Otherwise, the kids will be filtering for hours.  \n\nHave containers available for appropriate disposal of reagents. And make sure they clean up all of the bits of steel wool from the lab benches. Those little bits will haunt you for weeks if they don't.\n\n\n\n\nMethod\n\n1. Obtain 100 mL of copper (II) sulphate solution in\n   a clean, dry beaker. \n\n2. Obtain about 10 g of steel wool. Record the exact mass.\n \n3. Place the steel wool into the copper(II) sulphate solution.  Stir with a glass rod. When the solution loses its blue colour, add more copper(II) sulphate solution.\n\n4. Continue stirring (and adding solution, if necessary)until the strands of steel wool are no longer apparent, and only particles of a reddish brown solid are visible.\n\n5. Record the mass of a piece of filter paper.\n   Place it into a funnel.\n \n6. Pour the mixture through filter paper in a funnel.\n   Rinse any residue in the beaker into the filter        paper      with distilled water. Rinse the solid in the filter paper with distilled water.\n\n7. Allow the solid and filter paper to dry, then find the combined mass.\n\n\nAnalysis\n\n1. If iron was the limiting reactant, predict the mass of copper which should have been formed. Show work.\n  \n2. Calculate the percentage difference between  the mass of copper predicted in question 1 and the actual amount of copper produced. Show all work.\n\n3. Discuss possible explanations for your results in question 2.\n\n\n\nGood luck !
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Magnesium Sulfate\n\nMgSO4\n\nMg(OH)2 + H2SO4 -> MgSO4 + 2HOH\n\nHOH = H20\nJust in case...
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Cupric sulfate's heat capacity is described here.
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None. "A more active metal will displace a less active metal from a solution of its salt." Copper and lead are both less active than zinc. Obviously, zinc won't displace itself from zinc sulfate.