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Jadaav
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Does sodium reacting with water produces sodium hydroxide directly or produces sodium oxide first then hydroxide ?
I don't think you can get sodium oxide in that way.Borek said:Simplest method - roasting sodium carbonate.
lightarrow said:I don't think you can get sodium oxide in that way.
lightarrow said:I don't think you can get sodium oxide in that way.
Jadaav said:Thanks:)
Does Sodium reacting with Steam produces Sodium hydroxide ? I mean the metals above Magnesium can react with steam to form metal oxide or hydroxide ?
I'm confused, because here, for example, they say something different:Borek said:Most carbonates decompose in high temperatures yielding CO2 and metal oxides, sodium carbonate is not different. It is just a matter of temperature applied. From what I remember, standard Bunsen burner should be enough to heat the carbonate high enough.
lightarrow said:The rest of the Group 1 carbonates don't decompose at Bunsen temperatures, although at higher temperatures they will.>>
TGA curve of anhydrous sodium carbonate shows weight loss at the temperature range of 900–1175 K under static air atmosphere [9].
Thanks for that link, but sincerely from that diagram it's not so obvious for me to understand that it starts decomposing at 800°C (it even acquires weight some tens of degrees before that point) and that the decomposition is in Na2O and CO2.Borek said:That could be easy to check if someone has access to thermogravimetric curves for these carbonates. I am limited to what I can google, first one that I found
http://www.muhlenberg.edu/depts/chemistry/webmaps/sodacarb.htm
shows that sodium carbonate doesn't decompose up to 800 deg C, but Bunsen goes higher.
And are you sure that it refers to loss of weight due to decomposition in Na2O and CO2 and not, again, to the decomposition in carbonate and water?Google finds also this paper: http://www.tandfonline.com/doi/abs/10.1080/00986440108912851#preview. It can't be read without paying, but google shows quote from the paper:
TGA curve of anhydrous sodium carbonate shows weight loss at the temperature range of 900–1175 K under static air atmosphere [9].
lightarrow said:Thanks for that link, but sincerely from that diagram it's not so obvious for me to understand that it starts decomposing at 800°C
that the decomposition is in Na2O and CO2.
And are you sure that it refers to loss of weight due to decomposition in Na2O and CO2 and not, again, to the decomposition in carbonate and water?
Ok. So we still don't know for sure at which temperature Na2CO3 decomposes according to:Borek said:Never stated that, please reread my post - I wrote that it doesn't decompose before 800, it doesn't mean it starts at 800, it means it can be expected to decompose at some higher temp.
Maybe it's my english, I don't knowSorry, no idea what you mean by "decomposition in something". Could be my English fails me.
lightarrow said:Ok. So we still don't know for sure at which temperature Na2CO3 decomposes
and if it really does it
I mean that sodium carbonate at room temperature is never pure because it always has some water in it. So when we say it "decomposes" it's not clear which is the reaction
In that sense, ok. Yes, even water, e.g., decomposes into hydrogen and oxygen at temperatures high enough, but in the case of sodium carbonate how do we know that it doesn't decompose, e.g., in Na and CO2 and O2 or Na+, CO3-- and then in Na, C, O?Borek said:I have no doubts it does. There are no other stable substances it can decompose to, and it has to decompose at some point. Everything does, it is just a matter of temperature.
Right, even because, otherwise it wouldn't be possible to make standards heating it...No, it is clear. Water leaves crystals at much lower temperatures, before we get to 800 deg C (which we know is still not enough) carbonate is dry as Sahara desert.
When sodium reacts with water, it produces sodium hydroxide (NaOH) and hydrogen gas (H2) as byproducts. This process is highly exothermic, meaning it releases a large amount of heat.
Sodium is a highly reactive metal, and it reacts with water because it wants to lose its extra electron in order to become more stable. This reaction helps sodium reach a more stable state and releases energy in the form of heat.
Yes, the reaction between sodium and water can be dangerous if not performed carefully and under controlled conditions. The release of heat and flammable hydrogen gas can lead to explosions if not handled properly.
The chemical equation for the reaction of sodium with water is: 2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)
An exothermic reaction releases heat energy, while an endothermic reaction absorbs heat energy. In the reaction between sodium and water, the release of heat energy is a sign of an exothermic reaction.