/sci/ - Science & Math

[Anonymous No.7083775]

i never lurk this board, so i am not sure the protocol for posting here but..

given my tenuous understanding of chemistry, i was thinking that one could saturate water with sodium carbonate (or bicarbonate), then add an acid (i am thinking hcl dilute) to remove the carbonate...

then if you evporated the water, would you be left with pure sodium?(perhaps contaminated with hcl that didn't fully evap?)

is this an actual method to extract pure sodium?

[Anonymous No.7083778]

pure sodium is a group 1A metal. Group 1A metals go boom with H20

[Anonymous No.7083780]

like, in my head it seems wrong as fuck because i thought you need electrodes to do some fancy electrolysis and other such fancy things i can't comprehend.

but i can't think of why it wouldn't work.

[Anonymous No.7083783]

>>7083778
right, which is why this seems so wrong to me.

but if you have sodium carbonate.

and acid removes the carbonate.....

so you remove the acid.....

what is left?

[None !ELBOWS6FNQ No.7083788]

>>7083775
No, no it is not. Read a really basic textbook on chemistry.

http://hyperphysics.phy-astr.gsu.edu/hbase/chemical/acidbase.html

basically your acid neutralises with the base by making water, and everything left over sticks together.

this is a retard-tier explanation and isn't totally accurate but I hope it makes sense to you. (The reality involves dissolved hydronium and hydroxide ions.)

[Anonymous No.7083789]

>>7083775
sodium carbonate isn't na, its na+

[Anonymous No.7083800]

short answer: no
long answer: to get pure sodium, you'd need to reduce it from the +1 oxidation state to the 0 oxidation state, which can't be done in aqueous solution because metallic (0 oxidation state) sodium reacts vigorously with water to form sodium hydroxide and hydrogen gas. You'd also need a pretty strong reducing agent, since metallic sodium is quite electropositive.
Your proposal would leave you with sodium chloride, ordinary table salt.

TL;DR: learn redox

[Anonymous No.7083832]

>>7083783
>I have a basic solution of sodium carbonate
>Add a sufficiently strong Bronsted-Lowry acid, H(anion) in molar quantities
>Acid reacts with the carbonate, producing carbon dioxide and water
>Carbon dioxide mostly goes away, some stays in solution.
>So now in solution what's left?
>There's water, obviously, both from the initial solution and from the acid-base reaction.
>The anion from the acid
>The sodium from the sodium carbonate
So what's left is just the sodium salt of whatever anion you added originally. That is, if you added sulfuric acid, you now have sodium sulfate, if you added hydrochloric acid, you now have sodium chloride. And so on.

If you wanted get sodium in pure form, you would have to find a sufficiently strong reducing agent to give its electron to the sodium cation.

[Anonymous No.7083842]

>>7083832
Also, if you want to see some seriously badass reduction, you can even reduce sodium past its neutral state to get sodium hydride. If you think sodium in water videos are pretty cool, you'll probably shit your pants when you see a spatula tip of NaH dropped into water.

[Anonymous No.7083851]

>>7083842
>implying Na-
last I checked hydrogen was significantly more electronegative than sodium, and sodium hydride is generally conceived of Na+ and H-

[Anonymous No.7084240]

>>7083851
There is actually evidence to suggest that in certain instances (where H+ is able to be sequestered by cage ligands, for instance) Na- ions can exist. There's a short stub on 'inverse sodium hydride' on the wiki entry.

[Anonymous No.7084244]

>>7083851
>>7084240
to clarify however, yeah, you're absolutely right that NaH is typically thought of as Na+/H-