Chemical Forums
Chemistry Forums for Students => High School Chemistry Forum => Topic started by: Big-Daddy on May 03, 2013, 03:18:39 PM
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How do we write a balanced equation (not ionic) for the reduction of NaClO3 into ClO2 (SO2 is oxidized into SO42-)?
I've gotten the idea that even when writing ionic equations, if we've got an acid such as HClO3, it's best to keep this acid together. e.g. don't write the half-equation starting with the ClO3- ion but rather with HClO3 as a whole, and this will change the number of H+ we add later to balance the hydrogens, etc.
For ionic equations, I would normally leave out the Na+ and treat the anions only for this case, as Na+ must be a spectator, but we want an overall balanced equation, not ionic, to be the final result. How do I work it out?
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Unless you are dealing with rate laws and the species are not fully ionizable, or you need to incorporate enthalpies/entropies of dissolution, then I see no reason to include spectator counterions. It's just something else to screw up.
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Unless you are dealing with rate laws and the species are not fully ionizable, or you need to incorporate enthalpies/entropies of dissolution, then I see no reason to include spectator counterions. It's just something else to screw up.
But the final equation shouldn't have any ions in it ... so it seems best just to work with the Na still in there.
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But the final equation shouldn't have any ions in it
Really?
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But the final equation shouldn't have any ions in it
Really?
What I mean is, they want an answer with the Na written in. I can't just say, oh, it's Na+ so let's leave it out.
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In most cases I've seen, spectator ions are specifically left out. UNLESS dissolution is part of the overall process. And it's in my mind wrong because the whole point of a reaction is to show change. Why show something that doesn't change? You might as well show everything else in the solution that doesn't change, either. Dissolved gas, trace metals, etc.
I mean, if you have to do it, do it, but there's no sense to it from my viewpoint.
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I mean, if you have to do it, do it, but there's no sense to it from my viewpoint.
How do I do it?
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If ClO2 is one product, should I just assume there is only one other product, therefore it must be Na2SO4 and we balance to 2NaClO3 + SO2 :rarrow: 2ClO2 + Na2SO4? This was the right answer, but it is not given anywhere in the question that there is only one other product or that our other product is Na2SO4, how was I supposed to know?
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SO42- is a typical final product of redox reactions in which sulfur is present as a reducing agent. Sometimes it is also elemental sulfur (mostly when we gently oxidize sulfides).
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SO42- is a typical final product of redox reactions in which sulfur is present as a reducing agent. Sometimes it is also elemental sulfur (mostly when we gently oxidize sulfides).
OK - but in this case it's not S but rather SO42- because if we oxidize SO2 then S has to go up in state from +4 and only +6 (sulphate) is left, S8 would be a reduction. Thanks for the help. :)
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Is it safe to assume that products with ions like S2O32-, S4O62-, S2O82-, etc. are unlikely to result when sulphur is oxidized or said to be involved as a reducing agent?
The question would then arise, how do we get these ions then?
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Is it safe to assume that products with ions like S2O32-, S4O62-, S2O82-, etc. are unlikely to result when sulphur is oxidized or said to be involved as a reducing agent?
Yes. Thiosulfate is not stable and on its own decomposes to sulfite and sulfur. In the presence of oxidizing agents it can get oxidized to tetrathionate or just sulfate. Tetrathionate is not stable either. Peroxydisulfuric acid decomposes like most compounds containing -O-O- group.
The question would then arise, how do we get these ions then?
They require well selected, specific conditions. Chances of producing them in a randomly prepared solution are quite low.