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Making Lead Nitrate

aga - 23-6-2014 at 10:57

Lead Nitrate was one of today's 'Make This' compounds, as the stuff to make it was at hand.

So,
Pb_(s) + 4HNO_3(aq) -> Pb(NO₃)_2(s) + 2NO_2(g) + 2H₂O_(l)

The Lead was a 10g solid chunk of lead, used as a ballast weight in a model airplane.
The Nitric acid was 6.8[M]and an excess of 30% was added over the stoichiometric quantity required.

Pretty orange gas evolved.

What i ended up with was a grey-white precipitate, with some (about 10%) unreacted metal, which was a surprise.

Reading that Lead Nitrate is soluble in water, i decanted the liquid from the reaction beaker, leaving mainly the grey/white precipitate, then added 100ml of water.
The precipitate refuses to dissolve.

What i get is a milky liquid, with most of the grey-white powder remaining as a solid.

Is it obvious to anyone where i went wrong ?

Zyklon-A - 23-6-2014 at 11:18

Lead impurities are the only things that come to mind.
Are you sure of the purity? It likely contains antimony, tin and other impurities if it was just a weight.
Either way, it's probably possible to use the solution you have, as it should be reasonably pure lead nitrate (aq).
Two recrystallizations should do the trick.

aga - 23-6-2014 at 11:20

ahhh. so the precipitate is SOS^* and the lead nitrate is still in solution !

Obvious ! Doh ! Well, obvious now.

So i'll filter out the SOS and boil down the liquid.

Many thanks.

*Some Other Sh1t

[Edited on 23-6-2014 by aga]

blogfast25 - 23-6-2014 at 12:37

Careful boiling in lead bearing liquors: avoid things getting airborne with an hourglass (or similar) on the beaker. Lead (II) salts are very poisonous.

The state symbols (s), (g) etc are not usually written as suffixes. Just plain old Pb (s) will do...

aga - 23-6-2014 at 12:50

I will be careful.
Fume hood, fan over the beaker, as low a heat as possible to remove the water.

Sub tags are a pain to put in, so i just left the (s) etc as _subscript.

Edit : Which is right ?

1. NH_3(aq)
2. NH₃(aq)

You mean #2 right ? Do it non-subscripted ?

[Edited on 23-6-2014 by aga]

jock88 - 23-6-2014 at 14:22

Had a similar problem. The SOS I think is Tin Oxide.
The 'Lead' sheet may be old and contain Tin. The Tin is worth a few bob.

How would you be sure if it is Tin Oxide?

HgDinis25 - 23-6-2014 at 16:53

After filtering and getting a clear solution you could add nitric acid to it. The nitrate precipitates out of solution and you can filter out. Give it a few washes with cold Ethanol (to clean excess acid) and let it dry. It's the safest way to make dry Lead (II) Nitrate. No eating required.

thesmug - 23-6-2014 at 23:01

Well I'd suggest not eating lead compounds. Checking for tin oxide is not so easy. Here's an old journal entry that might help: http://pubs.acs.org/doi/abs/10.1021/ja02042a009

sasan - 24-6-2014 at 00:38

Guys,it is too complicated and maybe all of you are right.The lead it self can form lead oxides in nitric acid(I mean PbO that has 2 colors and PbO2) like dissolving Fe in nitric acid that forms iron oxides and just a tiny amount of iron go to the solution and make ferric nitrate(this problem contains nitrate of Pb/Fe/Hg/Bi/Sn/Ag/and some other elements)

For this problem I read somewhere to use a special procedure that I have
forgotten but I think he was using a certain concentration and density of nitric acid

I don't think so if it has Sb or Sn impurities.because they are valuable enough to not
use in toys and lead has a easy procedure for purifying.

I suggest you guys to use mixture of acetic acid and hydrogen peroxide to make

soluble lead acetate instead of using nitric acid.

Nitric acid has more difficult chemistery.in solutions with different concentration it
behaves differently.It's not just a simple acid like HCl,it is more complicated

Industry production of lead nitrate involves treating of lead oxide and nitric acid,or
treating its ores with nitric and subsequent crystalization.

I have same problem with mecury that I have dissolved its orange oxide in nitric acid and got a white precipitate.and I don't know what is this white material(maybe mercury(I)nitrate).I should use acetic acid and hydrogen peroxide but it is too late and mercury have a high price in Iran

aga - 24-6-2014 at 01:54

Seems that some of the SOS is PbO judging by the colour and wiki image.

HgDinis25 - 24-6-2014 at 02:33

If you used excess Nitric Acid there is no room for PbO formation. Many Lead alloys have Tin that reacts wit Nitric Acid to form Tin Dioxide (insoluble).

MrHomeScientist - 24-6-2014 at 10:34

Yes, aga, non-subscripted for the state.
Lead is almost always alloyed with something to improve its mechanical properties - it's too soft to be useful for much of anything as the pure metal. Lead fishing sinkers are alloyed with antimony, so I'd expect your lead weight to also have some additives.

Brain&Force - 24-6-2014 at 10:39

In bags of those small crimp weights (I think that's what they're called, they're the tiniest size of lead weight) I have sometimes found individual weights made out of tin. It's a rarity, but it just goes to show that they aren't inspected for purity. (Tin itself needs to be alloyed to prevent the alpha form from destroying the weights.)

Metacelsus - 24-6-2014 at 10:45

When I made lead nitrate last year not everything dissolved. It's pretty common for lead to be alloyed with other stuff. Just filter it out.

blogfast25 - 24-6-2014 at 10:51

@Sasan:

What MrHomeScientist said. For real applications Pb needs to be alloyed for strength. Sb is a very common lead hardener. There are others. For battery lead I've even read about Ca as a hardener.

And I can confirm from experience that pure lead (> 99 %) dissolves in nitric acid quite easily and completely.

Lead acetate from lead, GAA and peroxide is no sinecure either: once the peroxide has been used up/depleted too much the reactions stop. You need to keep topping up with peroxide, much of which get wasted in side reactions. I've done this and it ain't easy.

[Edited on 24-6-2014 by blogfast25]

aga - 24-6-2014 at 11:20

Well, i filtered, got a clear liquid, then started boiling.

At 50% original volume there were many shiny white glistening crystals on the bottom of the beaker.

So i got greedy and boiled down to 10%.
Now there is just white powder that will not dissolve in water.

Bugger.

blogfast25 - 24-6-2014 at 11:29

Quote: Originally posted by aga

Well, i filtered, got a clear liquid, then started boiling.

At 50% original volume there were many shiny white glistening crystals on the bottom of the beaker.

So i got greedy and boiled down to 10%.
Now there is just white powder that will not dissolve in water.

Bugger.

Chances are that you've hydrolysed it, greedypants!

Not to worry: add a 'good dollop' of nitric acid and some water and gently heat with stirring to about boiling: the hydrolysate should redissolve but it could take some time. Add a bit more NA to dissolve the last stubborn bits.

Then allow it too cool and see if you get crystals. Lead nitrate has a strong solubility - temperature dependence, so boiling dry wasn't necessary: thermal recrystallization would force most of the lead nitrate out of solution anyway, after having evaporated say 50 % of the water.

http://en.wikipedia.org/wiki/Solubility_table#L

aga - 24-6-2014 at 11:36

Quote: Originally posted by blogfast25

Chances are that you've hydrolysed it, greedypants!

Not to worry: add a 'good dollop' of nitric acid and some water and gently heat with stirring to about boiling

Yay ! Cheers for that. I will give it a go.
Dollops away !

Thanks also for the Solubility table link.
I've been looking up each substance individually and then trying to find the solubility data in the blurb.

jock88 - 24-6-2014 at 12:19

Don't forget the Lead Salts Preparation thread.

http://www.sciencemadness.org/talk/viewthread.php?tid=5490

aga - 24-6-2014 at 12:29

Thanks for the Lead Salts Preparation link.

Next time i will know what search terms to use !

Edit:
@blogfast25 : that worked. Thanks !

[Edited on 25-6-2014 by aga]

phlogiston - 25-6-2014 at 11:22

One more thing:

Quote:

then added 100ml of water.
...
What i get is a milky liquid, with most of the grey-white powder remaining as a solid.

You used plain tap water, didn't you? The milkyness is due to a finely divided precipitate formed from the lead and carbonate/sulphate/phosphates/etc in tap water.

You should definately use deionized or distilled water when working with lead salts.
Solutions of silver salts also give a similar precipitate when mixed with tap water.

[Edited on 25-6-2014 by phlogiston]

aga - 25-6-2014 at 11:45

Nope.
I use Distilled De-ionised water in all things.

As soon as i threw the empty used beaker in the wash bucket, it turned the whole bucket of tap water milky white.

Amazingly i also have Silver Nitrate experiments on the go at the same time, and yes, they do the same to the tap water.

The milkiness in the lead nitrate beaker was lead oxide/hydroxide or similar, formed because i was greedy and tried to boil the liquid too far.

Edit: i should have mentioned that the solution had a stirbar running in it at the time.

[Edited on 25-6-2014 by aga]

Fantasma4500 - 25-6-2014 at 12:01

as i just some time ago decided to get my lead and copper back from one of my lead acetate 'cells' (Pb + CuAcetate) i had to decant it off a few times
to remove lead i reacted the liquid with NaHCO3 to get PbCO3
i decided to make Pb(NO3)2 from this
the solution seemed very cloudy, although there was excess HNO3 in it, and it was all totally dissolved
the salt of this then turned somewhat greenish when heated, heating it more removed all colour, strange.

i actually also recall some strange precipitate when i was making Pb(NO3)2 from HNO3 and Pb, but i think i managed to dissolve it.. if a solution goes cloudy you can add some acid to it, add too much and you will have a problem when you boil it away, though..

AJKOER - 29-6-2014 at 12:58

No HNO3, try starting with NH3, H2O2, and PbCO3. Now, per Atomistry.com, from an intermediate, Lead nitrite (see http://lead.atomistry.com/lead_nitrites.html ) to the nitrate to quote:

"Lead nitrite solution slowly decomposes thus:

3Pb(NO2)2 + 2H2O = Pb(NO3)2 + 2Pb(OH)2 + 4NO

a decomposition similar to that which nitrous acid itself undergoes"

I would assume one could treat Lead carbonate with Nitrous acid as a path to Lead nitrite:

PbCO3 + 2 HNO2 ---) H2O + CO2 (g) + Pb(NO2)2

Per Atomistry.com on Nitrous acid ( http://nitrogen.atomistry.com/nitrous_acid.html ), one of many paths to Nitrous acid to quote:

"Oxidation of ammonia with hydrogen peroxide produces nitrous acid, but there is always some ammonium nitrite present in the solution:

NH3 + 3H2O2 = HNO2 + 4H2O"

[Edited on 29-6-2014 by AJKOER]

aga - 29-6-2014 at 14:24

Er, Jokey A ?

The Lead Nitrate is already crystallised and in a jar, happily awaiting its next adventure.

Assume nothing, do the maths and do experiments.
Tends to work out as more fun, and more reactants !

Do you get your kicks out of randomly Googling chem stuff in order to respond to someone actually Doing some chemistry, or is it a Kick just to get some random chemist (as in an Actual Chemist) responding to to the stuff you've just googled ?

What is with the ammonia gig anyway ?
It's a chemical, yes, loosely bridging OC and IOC (in my mind at least), but nothing more.

Chemistry, to me, is Exploring the Theoretical, then Testing by Experiment.
Not Einstein's cheapskate Thought Experiments, but by Actual Experiments, to see what Actually happens.

You seem to be fixated on doing No experiments, and simply Googling to elcit any kind of response.

Brother or Sister, you are missing out on so so much, and appear to be learning less than i have in the past 3 months.

AJKOER - 29-6-2014 at 15:49

Aga:

Actually, there is more value in my comment than simply proceeding without HNO3. The nature of some of the unexpected results may disappear in a nitrite approach.
----------------------------------------------------

Since you mentioned experimenting, I just finish exploring the wonders of weakened gamma alumina/Al on a solution of royal blue copper ammonium hydroxide (prepared by the action of dilute H2O2 on Copper in aqueous household ammonia with some sea salt, an amazing rapid formation of the royal blue complex with a moment of sigificant bubbling from associated NH4NO2 decomposition).

The original royal blue transforms in 10 minutes to aqua blue and then clear upon the addition of the gamma weakened alumina/Al!

Talk about experimenting!

I suspect the formation of copper aluminate which breakdowns in solution leaving copper ions and Aluminum ions!!

Interestingly, if one dilutes the solution with distilled water, some color returns!
-------------------------------------

Sorry for the theory session on this thread, but I once worked with Lead salts and have since lost the desire to work with them further. This was more psychological than real as my brothers kids spent some time in Chile and returned to the USA with learning disabilities that slowly cleared up over years. I suspect heavy metal exposure to their developing minds when in Chile. Just scary stuff.

[Edited on 30-6-2014 by AJKOER]

blogfast25 - 30-6-2014 at 04:58

Quote: Originally posted by AJKOER

Since you mentioned experimenting, I just finish exploring the wonders of weakened gamma alumina/Al on a solution of royal blue copper ammonium hydroxide (prepared by the action of dilute H2O2 on Copper in aqueous household ammonia with some sea salt, an amazing rapid formation of the royal blue complex with a moment of sigificant bubbling from associated NH4NO2 decomposition).

The original royal blue transforms in 10 minutes to aqua blue and then clear upon the addition of the gamma weakened alumina/Al!

Talk about experimenting!

I suspect the formation of copper aluminate which breakdowns in solution leaving copper ions and Aluminum ions!!

Interestingly, if one dilutes the solution with distilled water, some color returns!
[Edited on 30-6-2014 by AJKOER]

This is for the most part just gobbledygook, as we've come to expect from you. No evidence presented, unsubstantiated claims made, contentious terms used.

Explain it step by step and people might get interested.

[Edited on 30-6-2014 by blogfast25]

aga - 30-6-2014 at 09:11

I think i have to agree with blogfast25.

Flapping about and waving an ammony-wand doesn't really engage anything with me at least.

It really would be far more interesting to see your detailed examination of tinfoil heated on a candle, then doused with OTC 3% ammonia, and some 3% OTC peroxide added, and a pinch of salt.

The hypothesis, equations, reactants (source, purity, possible contaminats etc), method, observations, conclusions ... you know, like a real scientific experiment, which would enable others to duplicate the experient and check the conculsions.

No matter if it is tinfoil and table salt : there is no shame in having impure reactants, so long as you Know that, state that, and maybe have some clue as to what the impurities are.

What is 'The' Royal Blue complex ?
I made some tetraamminecopper(II)sulfate early on, and that's a deep, rich, regal purply-blue ...

blogfast25 - 30-6-2014 at 12:36

Quote: Originally posted by AJKOER

I suspect the formation of copper aluminate which breakdowns in solution leaving copper ions and Aluminum ions!!

... for instance make no sense at all. As AJ should know well by now aluminate anions (Al(OH)4-

can only exist in conditions of high pH (at least at reasonable concentrations). At such pH values solvated Cu2+ cannot exist and will precipitate as Cu(OH)2 (hydrated). At very high pH values aluminate and copper (II) could coexist with the copper as the cuprate ion (Cu(OH)4-

.

I see no regime in which 'copper aluminate' can exist, either in solution or as a pure substance (precipitate). No to mention conditions in which the mysterious 'copper aluminate' would "breakdowns in solution leaving copper ions and Aluminum". The exclamation marks would be well deserved...

[Edited on 30-6-2014 by blogfast25]

AJKOER - 30-6-2014 at 21:40

Two points if I may. First, it should be easy to replicate my observation. Create a royal blue tetraamminecopper complex under alkali conditions (alkali per one of the half reaction in the electrochemical part of the reaction of Cu, H2O2, NH3 with sea salt as the preferential electrolyte, as a reference please see "Kinetics and Mechanism of Copper Dissolution In Aqueous Ammonia" fully available after signing on to ones Facebook account at http://www.academia.edu/292096/Kinetics_and_Mechanism_of_Cop... ). Next, burn to glowing red Aluminum foil to form, my contention based on chemical reactivity, weakened gamma Al2O3 and Al (reference, see white paper at https://www.google.com/url?sa=t&source=web&rct=j&... ).

Second, I am speculating on the chemistry, no references on this reaction (except perhaps http://www.sciencemadness.org/talk/viewthread.php?tid=20554 ). Combine with the tetraamminecopper complex and watch, it doesn't take long. What is evident is the gradual and complete break up of the royal blue tetraamminecopper complex and the formation of a clear, I suspect Aluminum, sallt.

[Edited on 1-7-2014 by AJKOER]

aga - 1-7-2014 at 03:14

Quote: Originally posted by AJKOER

it should be easy to replicate my observation

It would be easier for a total noob like me to have a detailed procedure to follow, showing exactly what the reactants are and what you do with them.

blogfast25 - 1-7-2014 at 05:02

Aga:

As a first step, mix some copper conductor wire (fine strands cut into pieces), some NaCl, some peroxide solution and some ammonia. Mix to dissolve the NaCl. Allow to stand and observe.

With both an oxidiser AND a complexing agent present, some copper is likely to dissolve, to dark blue copper (II) tetrammine complex. It's a 'going round the houses' way of preparing a bit of the latter.

Then we'll see.

Aluminium is notoriously difficult to burn due to that passivation layer that protects it against further attack.

I suspect AJ's solution loses colour simply because the copper (II) tetrammine complex is reduced to copper by the aluminium. Nothing mysterious... or new. It's known (there's a good thread about it somewhere) that that reduction can be very fast in the presence of chloride ions.

[Edited on 1-7-2014 by blogfast25]

AJKOER - 1-7-2014 at 06:32

Quote: Originally posted by blogfast25

Aga:

As a first step, mix some copper conductor wire (fine strands cut into pieces), some NaCl, some peroxide solution and some ammonia. Mix to dissolve the NaCl. Allow to stand and observe.

With both an oxidiser AND a complexing agent present, some copper is likely to dissolve, to dark blue copper (II) tetrammine complex. It's a 'going round the houses' way of preparing a bit of the latter.

Then we'll see.

Aluminium is notoriously difficult to burn due to that passivation layer that protects it against further attack.

I suspect AJ's solution loses colour simply because the copper (II) tetrammine complex is reduced to copper by the aluminium. Nothing mysterious... or new. It's known (there's a good thread about it somewhere) that that reduction can be very fast in the presence of chloride ions.

[Edited on 1-7-2014 by blogfast25]

Blogfast:

I cannot fault your logic as it was mine first!

My intention of performing the experiment was a quick and inexpensive path to fine copper.

However, as copper or its oxide is not visibly formed, it may have dissolved into some colorless(?) compound under alkali conditions. Aluminum also has apparently entered the solution.

To speculate further is clearly way off the topic of this thread.
---------------------------------

Here is a good reference on electrochemistry that at times includes Pb. See http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch20/el...

[Edited on 1-7-2014 by AJKOER]

Zyklon-A - 1-7-2014 at 09:04

Quote: Originally posted by aga

Er, Jokey A ?

Do you get your kicks out of randomly Googling chem stuff in order to respond to someone actually Doing some chemistry, or is it a Kick just to get some random chemist (as in an Actual Chemist) responding to to the stuff you've just googled ?

Who cares if he does that? I google things all the time and use the information I just learned to reply to posts. Lots of people here do that.
There's no reason to do an experiment for every question you have, and you'll learn a lot more from looking up actual scientific sources then any experiment. The reason is simple, when you do an experiment, there is so much room for error just from inaccurate measurements, impure reagents, wrong stoichiometry, different reactions happening than expected etc.

Quote: Originally posted by aga

What is with the ammonia gig anyway ?
It's a chemical, yes, loosely bridging OC and IOC (in my mind at least), but nothing more.

Ammonia is a pretty cool chemical. You saying that it's unimportant simply because you haven't used it much, only makes you look ignorant and arrogant.

Quote: Originally posted by aga

Chemistry, to me, is Exploring the Theoretical, then Testing by Experiment.
Not Einstein's cheapskate Thought Experiments, but by Actual Experiments, to see what Actually happens.

Or you could just read a paper that has all the information you need and save you a lot of time. Then you can spend your time doing reactions that benefit you in some way (like teaching you, or synthesizing a new, useful reagent (like ammonia

) )

Quote: Originally posted by aga

You seem to be fixated on doing No experiments, and simply Googling to elicit any kind of response.

Again, who cares? Maybe he dosn't have time to do experiments. Maybe he doesn't need to, cause he can find all the information he needs online.
He is contributing. Have you ever even followed the references he sites? I do, they are full of very relevant information generally.
So, instead of complaining about what he does post, why don't you try to learn something from it?

aga - 1-7-2014 at 11:17

If you do an experiment yourself, you learn much, especially when it goes totally wrong.

I can't see where i say 'ammonia is unimportant', most likely cos i didn't say that at all.
NH₃ is pretty cool, i agree, and making it i has been great fun the 5 and only times i have done so, especially when it went totally wrong the first time ...

At my 3-months-in level of learning, much of the published info is way over my head, fascinating though it is.

I agree that AJ's references are well worth reading, and i cannot remember one that was boring.

My boggle is basically the vagueness.
I can't mix 'vaguely some' gamma Fluffitol with a half-quoundle of Amoonyflip and even imagine to repeat the experiment accurately.

The last time AJ actually wrote down a do-able reaction, i did 5 of them, and reported the findings.

Quote: Originally posted by aga

blogfast25 - 1-7-2014 at 12:17

Quote: Originally posted by AJKOER

[However, as copper or its oxide is not visibly formed, it may have dissolved into some colorless(?) compound under alkali conditions. Aluminum also has apparently entered the solution.

There are no colourless, aqueous Cu(II) compounds, even less so when chloride and/or ammonia is present, alkaline or neutral.

The fact that the solution became colourless points to reduction of the Cu(II) to copper, it's the only reasonable way to explain the loss of blue.

The initial solution may have contained very little Cu(II) to begin with, as a the tetrammine complex is very intensely coloured and very small amounts will cause colouration (and the dissolution of copper metal in saline, peroxydic ammonia cannot be a fast process). If so, the amounts of reduced copper may be hard to see.

Al would have entered the solution as unburned Al reduced Cu(II) to copper, the Al being oxidised to Al(III).

There is in essence nothing to see here...

[Edited on 1-7-2014 by blogfast25]

blogfast25 - 1-7-2014 at 12:24

Quote: Originally posted by Zyklon-A

[ He is contributing. Have you ever even followed the references he sites? I do, they are full of very relevant information generally.

Sorry, but he also posts a lot of irrelevant, obscure and sometimes demonstrably faulty stuff too. AJ's a champ for cherry picking material that seems to confirm his ideas, even if the sources are far from credible. On other occasions he pulls things horribly out of context.

He's posted references to 'nascent hydrogen' for instance. I could go on.

AJKOER - 1-7-2014 at 15:03

Yes, true, and here are the details on hydrogen formed on an Aluminum surface as published in a 2012 prestigious journal, link http://www.researchgate.net/publication/221934434_Chemical_r... , titled "Chemical reduction of an aqueous suspension of graphene oxide by nascent hydrogen" by Viet Hung Pham, Hai Dinh Pham, ... in Journal of Materials Chemistry (Impact Factor: 5.97). 05/2012; DOI:10.1039/C2JM30562C .

To quote from the abstract:

"ABSTRACT One of the major challenges in the chemical reduction of graphene oxide is increasing the C/O atomic ratio of the chemically-converted graphene. In this paper, we report a simple and effective method to reduce aqueous suspensions of graphene oxide using nascent hydrogen generated in situ by the reaction between Al foil and HCl, Al foil and NaOH and Zn powder and NaOH. The nascent hydrogen-reduced graphene oxides (nHRGOs) were characterized by elemental analysis, UV-vis spectra, Raman spectra, X-ray photoelectron spectroscopy, thermogravimetric analysis and electrical conductive measurements. The reduction efficiency of graphene oxide strongly depended on the reaction medium and the rate of nascent hydrogen generation. The best nHRGO achieved a C/O atomic ratio greater than 21 and a bulk electrical conductivity as high as 12,500 S/m, corresponding to the nascent hydrogen generated from the reaction between Al foil and HCl. Since nascent hydrogen could be produced on a metal surface upon oxidation in solution, other metals with low standard reduction potentials, such as Mg, Mn, and Fe, can be applied to reduce graphene oxide."

The reason this article was published, in my opinion, relates to its value as a potential source of graphene.

For some background on why graphene is potentially important, see for example, http://www.telegraph.co.uk/finance/businessclub/10936423/Gra...

By the way, I for one, do not include the Journal of Materials Chemistry in the group of "sources are far from credible". However, when I do quote Atomistry.com as a source, beware, as it contents, per a recent discovery of mine, are extracts taken from noted journals of the time with, however, references deleted and no supplied date of publication.

[Edited on 2-7-2014 by AJKOER]

arkoma - 1-7-2014 at 22:47

How many times does poor aga have to say,"I am a n00b, please 'dumb it down' for me so I can get a handle"? Talk about arrogant pricks...............

EDIT-if no one has noticed, this is BEGINNINGS

[Edited on 7-2-2014 by arkoma]

blogfast25 - 2-7-2014 at 04:13

Quote: Originally posted by arkoma

How many times does poor aga have to say,"I am a n00b, please 'dumb it down' for me so I can get a handle"? Talk about arrogant pricks...............

EDIT-if no one has noticed, this is BEGINNINGS

[Edited on 7-2-2014 by arkoma]

Totally uncalled for.

Not everyone who posts in beginnings is a beginner either.

blogfast25 - 2-7-2014 at 04:18

Quote: Originally posted by AJKOER

However, when I do quote Atomistry.com as a source, beware, as it contents, per a recent discovery of mine, are extracts taken from noted journals of the time with, however, references deleted and no supplied date of publication.

[Edited on 2-7-2014 by AJKOER]

Which is essentially an act of plagiarism.

I've also found atomistry.com demonstrably wrong in one of their entries. An interesting and useful site but take content with a pinch of salt.