Sciencemadness Discussion Board - Black gunk that should be NaNO3

Black gunk that should be NaNO3

Jylliana - 3-12-2014 at 05:49

I did something... well.. not really stupid imo, but a little stupid.
I tried to neutralize my 65% nitric acid solution, so I added a few spoons full of sodium carbonate... but the pH crept up too slowly, so I proceeded by adding sodium hydroxide... and then this happened:

It started to boil very quickly and violently and some sticky black gunk formed. It looks a bit like molten rubber.
The rise in temperature actually made my plastic, teflon coated stir bar all flimsy, like it melted or something, and on cooling down it went back to it's original state.

Internet stated that this reaction should yield NaNO3, but that stuff is white, and this was black.

Does anyone know what happened?

PS The solution also had copper ions present(I dissolved metallic copper in the acid), but I don't know if that influenced the reaction.

[Edited on 3-12-2014 by Jylliana]

forgottenpassword - 3-12-2014 at 05:57

I don't know any copper chemistry, but surely that explains it: you've made some black copper compound. If there was only a little in there perhaps you can recrystallize the sodium nitrate using a suitable solvent. What that solvent is, I couldn't tell you.

Magpie - 3-12-2014 at 06:14

If there was any extra un-neutralized Na2CO3 this would be a source of carbon. At high temperature this can lead to black gunk.

Jylliana - 3-12-2014 at 06:18

Of course... simple redox chemistry... why didn't I think of that >.<

Quote:

You are doing the classic "copper cycle" lab. I started using this with my students in 1972. Copper was dissolved in nitric acid to make a copper (II) nitrate solution, then neutralized to Cu(OH)2, then heated to CuO.

Source: https://answers.yahoo.com/question/index?qid=20090310142739A...

Copper oxide is black

[Edited on 3-12-2014 by Jylliana]

Jylliana - 3-12-2014 at 06:19

Quote: Originally posted by Magpie

If there was any extra un-neutralized Na2CO3 this would be a source of carbon. At high temperature this can lead to black gunk.

Maybe, but is nitric acid able to convert the carbonate to elemental carbon?

[Edited on 3-12-2014 by Jylliana]

forgottenpassword - 3-12-2014 at 06:37

If you have finished neutralizing the solution, perhaps you can dissolve it in hot water and filter out the black precipitate. If that clears up your solution you can boil it down a little then dry it over a low heat. You will probably need to find a recrystallization solvent to clean it up to an acceptable level; but I'm sure that that will be easily achieved.

[Edited on 3-12-2014 by forgottenpassword]

Bert - 3-12-2014 at 06:41

Welcome to the wonderful world of chemistry. There are no failed experiments, just unexpected learning opportunities!

Just looked at the biographical information in your profile-

You are involved in high school level science teaching as a technician? You have graduated both high school and some sort of college or technical school?

Could you please leave your original post unedited, but now write your exact lab procedure, explaining your intent and actions, providing stoichiometric information.

You proceeded on the basis of:

Quote:

Internet stated that this reaction should yield NaNO3

Please provide a link to your source.

After you write this up to include the level of information required, I will be happy to move the thread out of "beginnings"!

Jylliana - 3-12-2014 at 07:16

Procedure:
- Add 50ml 65% nitric acid to a beaker
- With a spatula, add a few grams of metallic copper(turnings).

- When the reaction is finished, add three or four tablespoons of sodium carbonate.
- Stir until dissolved
- Add one tablespoon of sodium hydroxide to the solution

I did all of this because I always neutralize solutions before I throw them in the (heavy metals) waste bin.
I didn’t measure out everything precisely, I described my actions precisely the way I actually did my experiment.

The conclusion that this reaction should yield NaNO3 is from this source:
https://answers.yahoo.com/question/index?qid=20081031171657A...

When I wrote that part of my post, I hadn’t thought of the copper ions in solution yet, that’s why I looked up ‘HNO3 + NaOH’.

-

Yes, I am a School Science Technician. I finished High School and also went to college in Analytical Chemistry, and took extra classes to prepare me for working at a school.

I hope this is enough information.

Bert - 3-12-2014 at 08:48

Quote:

I didn’t measure out everything precisely, I described my actions precisely the way I actually did my experiment.

An EXPERIMENT is an orderly procedure carried out with the goal of verifying, refuting, or establishing the validity of a hypothesis.

What was your hypothesis? By what standards would you have considered it to have been validated???

An orderly process would include measuring all reactants and keeping notes. Measuring of dry chemicals would be by weight, not "3 or 4 tablespoons"

Children might mix some unmeasured chemicals in a beaker with the vague intent of having "something happen". If you are an adult, a college educated chemist and work as a technician- Something seems very odd in all of this, 12 year old self taught members here will conduct operations in a more rational way.

Magpie - 3-12-2014 at 09:25

Quote: Originally posted by Jylliana

Maybe, but is nitric acid able to convert the carbonate to elemental carbon?

More likely it would go to CO2.

blogfast25 - 3-12-2014 at 10:11

Quote: Originally posted by Jylliana

Of course... simple redox chemistry... why didn't I think of that >.<

Redox chemistry? Where is the oxidation and the reduction? (There's none, it was a rhetorical question)

Redox, no, reductionist, probably.

During addition of alkali, Cu(OH)2 formed + neutralisation heat. Heat + Cu(OH)2 = CuO, as already noted.

Jylliana - 3-12-2014 at 10:14

Quote: Originally posted by Bert

Quote:

I didn’t measure out everything precisely, I described my actions precisely the way I actually did my experiment.

Fine, not an experiment. I didn't research anything. I expected the solution to be neutralized with an excess NaOH or carbonate. That is all.
And as far as I know, an excess is an excess, whether it's 50mg or 10 grams... still an excess. Thats why I used spoonfulls and not grams.

And I would appreciate it if you'd use a friendlier tone in your posts. I did not mean any harm and I always try my best to think before doing. We can't all be super smart people who anticipate every tiny problem.

Jylliana - 3-12-2014 at 10:16

Quote: Originally posted by blogfast25

Quote: Originally posted by Jylliana

Of course... simple redox chemistry... why didn't I think of that >.<

Redox chemistry? Where is the oxidation and the reduction? (There's none, it was a rhetorical question)

Redox, no, reductionist, probably.

During addition of alkali, Cu(OH)2 formed + neutralisation heat. Heat + Cu(OH)2 = CuO, as already noted.

Precipitation reaction, then? I don't know... I'm getting more confused by the minute.

Bert - 3-12-2014 at 10:30

I must apologize for my overly critical and narrow interpretations of your statements-

Mere words. No reason to be so fussy on exact definition in a scientific context! We all can fill in any small lacunae.

Speaking (writing?!) of words, I wish to complement you on your perfect use of the English language, and fitting use of USA standard idiom and phrasing. It is so rare for a person using English as a second language to use it as you do, with absolutely no trace of the sentence structure or idiom of their "milk tongue". It is a profound skill, especially in one so young.

MrHomeScientist - 3-12-2014 at 10:36

I'll chime in with another confirmation that it is copper(II) oxide. Your copper nitrate solution precipitated copper carbonate and hydroxide upon addition of the various bases, and the high heat decomposed the hydroxide to black CuO. I'm actually planning on using this reaction (in a more controlled manner) this weekend to introduce some kids to chemistry.

Magpie - 3-12-2014 at 10:54

Quote: Originally posted by Bert

It is so rare for a person using English as a second language to use it as you do, with absolutely no trace of the sentence structure or idiom of their "milk tongue". It is a profound skill, especially in one so young.

Yes, it is awesome. Garage chemist, a German forum member, also has this skill.

Jylliana, can you speak fluently in US idiom as well? Do you have an American accent, or is it British, or is it Dutch?

Bert - 3-12-2014 at 10:56

It's a very common high school or college introductory chemistry lab-

I recall doing it myself, back when high school kids in USA were allowed to handle more "interesting" reagents .

DraconicAcid - 3-12-2014 at 11:04

Quote: Originally posted by Magpie

Do you have an American accent, or is it British, or is it Dutch?

I'm sure she doesn't have an accent at all.

(Actually, people are usually very bad at judging their own accents, or lack thereof, so you might not a useful answer to that question. I once worked with a French woman who told me, "You 'ave zee vairy stonk akSENT. Zee peepl 'oo teach me Eeng-lessh, zhey 'ave no akSENT.")

Amos - 3-12-2014 at 11:05

I remember that lab! It happened just a month or two before I started home chemistry, and it's part of the reason I wanted to start in the first place.

And it might be a good idea to be a little more accepting and constructive towards newer members; the last thing we need is to drive people away from what is probably an already shrinking field. I was terrified of you veteran users when I first joined the forum.

Jylliana - 3-12-2014 at 11:19

Quote: Originally posted by Magpie

Quote: Originally posted by Bert

Here is a video I narrated for Rador Labs. You can hear my voice there, if you're interested.
I gotta add, I was nervous for this video. My day-to-day english is a bit more fluent and I normally don't stutter.

Bert, thanks for the compliment

. I got my english skills from watching too many american sitcoms and movies

blogfast25 - 3-12-2014 at 11:29

Quote: Originally posted by Jylliana

Precipitation reaction, then? I don't know... I'm getting more confused by the minute.

Yes. A simple precipitation followed by thermal dehydration of the precipitated hydroxide. A bit unexpected in your experimental context.

Magpie - 3-12-2014 at 11:42

I would say a very slight Dutch accent.

Accents vary all over the US. When I was in the military I used to have fun trying to guess the home states of other soldiers based on their accent. I usually came pretty close. One person thought I was from Pennsylvania - he was way off - I'm from Idaho.

woelen - 4-12-2014 at 04:37

I think that the following happened in Jylliana's "experiment".

First you made copper nitrate, dissolved in excess nitric acid.
Addition of the Na2CO3 creates CO2 and NaNO3. NO C IS FORMED! There is no redox reaction. That would require a reductor and not an oxidizer like HNO3 and it would require a STRONG reductor and elevated temperatures.
Addition of too much NaOH leads to formation of NaNO3, Cu(OH)2 and a lot of heat, due to the neutralization reaction between HNO3 and NaOH. You added excess NaOH. The heat makes all of the Cu(OH)2 decompose to CuO and water. Cu(OH)2 even can split off water molecules when it is under water!

Try this little experiment, it is very instructive, and may also be nice as a demo experiment for your classes:

- Prepare a solution of NaOH. Just add a spatula of solid to a few ml of water and dissolve. Allow to cool down. Fast cooling down can be done by keeping the test tube under slowly running tap water for half a minute or so, while swirling it around a little bit.
- Prepare a solution of CuSO4.5H2O. Be sure that you use less copper sulfate than NaOH.
- Mix the two solutions.

You get a beautiful bright blue precipitate of Cu(OH)2.
Now heat the test tube with the precipitate of Cu(OH)2, while gently swirling it around. It quickly turns black completely, at a temperature well below the boiling point of water. The Cu(OH)2 decomposes to CuO and H2O.

j_sum1 - 4-12-2014 at 04:59

Thanks woelen. I am sure I have done all of these steps before but not to deliberately investigate anything. It will be a good quick exercise next time I am in the lab. And yes, one that is graphic enough to be a good demo for students. Quite a few principles can be illustrated by this one.

Jylliana - 4-12-2014 at 04:59

For the record;
I managed to clean up the 'disaster' from yesterday and tried again today. This time, I added the NaOH more slowly, and finished with adding an excess sodium carbonate very slowly. I did this while squeezing acetone vapor out of a bottle to keep the foaming under control.
What I ended up with, was a beautiful neutral suspension/slurry of copper hydroxide

This is what I intended to do yesterday, but back then I added everything too quickly and the heat produced by the NaOH converted the formed copper hydroxide to copper oxide, along with some excess NaOH pellets and very little water(it all splattered out or evaporated). This is what probably made it look like molten rubber.

Thanks for the idea, woelen, i'll give it a try. I have both solutions in stock so it'll be nice, quick and simple to try ^_^

[Edited on 4-12-2014 by Jylliana]