[Edited on 7-3-2012 by disulfideprotein]bfesser - 5-3-2012 at 15:54
It's not a bad start, but it could use some editing work. Personally, I'm a fan of the two-column format used in many modern journals. A smaller
font would make reading a bit easier. Also, wrapping text around images isn't very professional in appearance. Maybe just put an imaginary box
around the images, and a figure label (with caption if necessary) below each.disulfideprotein - 6-3-2012 at 08:37
Okay I will work on it, thanks!Hexavalent - 6-3-2012 at 08:49
Nice, although I agree with bfesser. Additionally, you might want to add mole amounts of your chemicals too, and sort out your references. It would be
nice to see more primary links, and NEVER use Wikipedia directly as a listed source.
If you don't know how to calculate the moles, then review the formula triangle relating mass, molarity and Mr.
Also, consider cutting down on the image size on the first page, and calculate your percent yield.disulfideprotein - 6-3-2012 at 09:49
Thanks, I will put in the mole amounts (I put in the moles in grams!) and conclude that this is reversible reaction. Also wikipedia was JUST for the
image. Yes I know how to do the calculations. I will put that in and update it. Also I will put more direct links and add a few more notes.
[Edited on 6-3-2012 by disulfideprotein]bbartlog - 6-3-2012 at 10:54
I was under the impression that this subforum was for descriptions of actual experimental work, not pretty writeups of how something works in theory.disulfideprotein - 6-3-2012 at 11:17
I did do this experiment! IT DID WORK! This is an explanation and instructions to the process and comparison of different methods. IT IS NOT THEORY, I
POSTED THIS SO OTHER PEOPLE COULD LEARN ABOUT IT AND TEST IT TOO.bfesser - 6-3-2012 at 13:19
disulfideprotein, few more suggestions, but by no means a comprehensive edit:
<ul type="circle">
<li>copper(II) hydroxide instead of cupric hydroxide, although any experienced chemist knows cupric vs. cuprous</li>
<li>break into paragraphs as shown, ". . . done method of making Schweizers reagent. ¶ Why? Schweizers reagent is used in . .
."</li>
<li>" . . . to just use the 2 step process." → " . . . to just use the two-step process."</li>
<li>Avoid the use of "I" and "you" in technical writing. <em>e.g.</em> "Now, as you might know Schwiezer. . ."</li>
<li>put each reaction equation on a separate line and label them as equation 1 and 2</li>
<li>"C6H10O5" → "C<sub>6</sub>H<sub>10</sub>O<sub>5</sub>", "NH4OH" →
"NH<sub>4</sub>OH", "Cu(OH)2" → "Cu(OH)<sub>2</sub>"</li></ul>
<hr width="300" />
Hexavalent, you catch on rather quickly.
Nice, although I agree with bfesser. Additionally, you might want to add mole amounts of your chemicals too, and sort out your references. It would be
nice to see more primary links, and NEVER use Wikipedia directly as a listed source.
Be careful not to plagiarize the work of others. It would be nice to see some more primary sources in your writeup--literature references, patents,
etc. Wikipedia should </strong>never</strong> be listed as a reference. All Wikipedia articles have a list of references from which
material and facts have been drawn.
kavu - 6-3-2012 at 13:52
Reading some articles from journals would help you see what the commonly accepted format of presenting results is. Few general points: Using sections
will make the writeup easier to read and also experimental is to be written in passive voice. Numbered lists are not generally used. An example of an
experimental (synthetic) procedure would go something like this (synthesis of phthalimide that I used to make phthalide):
In a 100 ml RBF powdered phthalic acid anhydride (1,0 g; 6,75 mmol) was mixed with urea (0,2 g; 3,33 mmol). Mixture was heated with a heat gun until
molten. A vigorous reaction producing white fumes proceeded and the oil solidified. After cooling to RT 20 ml of water was added and solid filtered.
Crude phthalimide separated as white powder (0,8 g; 80 %; m.p. 233 °C lit. 238 °C).bbartlog - 6-3-2012 at 16:00
I did do this experiment! IT DID WORK! This is an explanation and instructions to the process and comparison of different methods. IT IS NOT THEORY, I
POSTED THIS SO OTHER PEOPLE COULD LEARN ABOUT IT AND TEST IT TOO.
Your instructions frankly don't sound like they were written by someone who has done the experiment, or previously written up experimental results.
I'm open to the possibility that you may actually have done this and are just really bad at describing what you did, but:
1) You used a full mole of anhydrous CuSO4? Really?
2) You put it (dry?) in a beaker and then added 'NH4OH', by which I have to assume you mean an ammonia solution of some strength you don't specify,
and which was therefore measured in some way you don't describe.
3) All your weights are given to sub-milligram precision... why?
4) You don't offer any visual description of what happened. Or what size beakers you used (improbably large ones by the sound of it). Or how much
water was ultimately involved. Or how much toilet paper was dissolved. Or how many syringes full of HCl you ultimately needed to neutralize the mole
of reagent you mixed up.
I can't speak for others but I personally believe that your report is probably a fabrication, or at best a rewrite of something you did in a quite
different way.disulfideprotein - 6-3-2012 at 16:32
bbartlog, I used molar amounts but one mole, .1 mole it would be a waste to use so much CuSO4. They are given in sub-milligram precision because if
some one has a better balance they may be able to get better measurements. My report is not a fabrication! I am uploading a new version with
corrections. The size of the beakers and other insterments are based on the fractions of the moles the person would like use. I said to put the HCl
into the solution until precipitation rayon stopped. The next version will have pictures.bbartlog - 6-3-2012 at 17:11
OK, in that case I apologize for accusing you. But yah know, if *you* did it using a tenth of a mole (which sounds a lot more plausible, now we're in
the land of 500ml beakers rather than 3 liter ones), just write it up that way! You'll be better served by just offering an account of your
experiment, including the precision of your own measurements and so on, than by what looks like an attempt to turn it into a much more generic set of
instructions.disulfideprotein - 6-3-2012 at 18:27
Thanks, now since I don't want to reuse more of my chemicals I am posting newly synthesised micro examples. (less that 50 ml). I have posted a new
version. Please give me more feedback. I am trying to correct everything that I can. -Please note that I am currently doing other corrections at this
time-
[Edited on 7-3-2012 by disulfideprotein]bbartlog - 7-3-2012 at 23:01
Much improved. There is a likely error in your ammonia stoichiometry, though. You calculate the required amount based on the molecular weight of
'NH4OH' (so 0.2 mole is about 7g) and from there end up using ~70g of 10% ammonia solution. However, as I understand it, the percentage strengths
given for ammonia solution are based on the weight of NH3. 70g of 10% ammonia solution contains over 0.4 moles of ammonia, *not* 0.2.
This may explain why in your experimental you only end up adding 70g of ammonia solution twice, rather than 70g and then 140g as your calculations
would seem to require.
You might want to explicitly specify that your CuSO4 was anhydrous and give some alternate instructions for those who only have the more commonly
available pentahydrate at hand.
You should also indicate the strength of your HCl. If it was the 31% commercial stuff then filling a 500ml beaker for neutralization purposes was
massive overkill (that's like 5 moles).
Finally there are still some minor typos ('was be formed', 'cuperic' for example).barley81 - 8-3-2012 at 04:04
<a href="http://www.chem.ox.ac.uk/vrchemistry/complex/default.html">This site</a> shows that the precipitate made by adding NaOH to
CuSO<sub>4</sub> solution is actually basic copper sulfate. I never knew that. You might want to include something about that in the
write-up (probably the same thing happens when ammonia is used in place of NaOH). The new version is much better.
[Edited on 8-3-2012 by barley81]disulfideprotein - 8-3-2012 at 08:13
<a href="http://www.chem.ox.ac.uk/vrchemistry/complex/default.html">This site</a> shows that the precipitate made by adding NaOH to
CuSO<sub>4</sub> solution is actually basic copper sulfate. I never knew that. You might want to include something about that in the
write-up (probably the same thing happens when ammonia is used in place of NaOH). The new version is much better.
[Edited on 8-3-2012 by barley81]
I do not believe from cupric sulphate we get cupric sulphate. When the ammonium hydroxide is dissolved in the solution the exchange of Ions is like I
said. Directed at the last comment: Thank you, that is why I published this method: It is easier and faster, and you dont get inefficient percipients
that you need to wash out.
Bbartlog, I will get on to that. Thank you for spotting that error.
