New answers tagged aqueous-solution
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I'm not sure of any conventional ways of electroplating selenium, but you could use a technique such as magnetron sputtering to spray a thin film of selenium onto another metal. Heck, you could even use glass or wood if you wanted to do it that way. And if you're going for thin, this technique can get you some of the thinnest, smoothest layers of metal out ...
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Na2CO3 is a salt of a weak acid and strong base.
As you may recall weak acids (the same applies for weak bases) have tendency to have it's equilibrium towards the extreme left of the following reaction -
$\ce{HA <=> H+ + OH-}$
The Na+ ions released from the dissociation of Na2CO3 form NaOH is in minuscule or insignificant quantities as NaOH is ...
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Sodium carbonate indeed dissociates in water : $$\ce{CaCO3 <=> Ca^2+ + CO3^2-}$$ But that is just beginning, as in solutions with carbonates, bicarbonates or dissolved carbon dioxide happen multiple chemical equilibriums:
( See a lot of info at Carbonic acid on Wikipedia):
Carbonate anion partially hydrolyzes in water: $$\ce{CO3^2- + H+ <=> HCO3-...
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The reaction system can be described as:
$\ce{H2O <=> H+ + OH-}$
$\ce{H+ + CO3^{2-} <=> HCO3-}, \mathrm{p}K_\mathrm{a2}=10.33$
Note, the net of the first two reactions imply a rise in pH in the presence of carbonate. Further, with time and carbon dioxide exposure:
$\ce{H2O + CO2 <=> HCO3- + H+}$
$\ce{H2CO3 <=> H2O + CO2}$
And, ...
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Another possibility is the grease spot photometer (aka Bunsen photometer) German Wiki page. This can be homemade, you need a piece of paper, some wax or oil, two light sources, a meter stick and for the measurement of solutions also some e.g. cardboard to shield unwanted light. Also the darker the room, the better.
The underlying idea is that light ...
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It is still possible to measure light absorbances of a colored solution without any photometer, simply by comparison.
First you prepare a reference solution with maybe 10 ppm Pb and an excess of a colorless reagent making a colored compound or complex with Pb. You prepare a series of 10 test tubes. You add 1 mL of this reference solution in the first test ...
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Perhaps you can add H2S to create PbS (as a starting step). Then, based on an extract from Wikipedia on Lead Sulfide:
Although of little commercial value, PbS is one of the oldest and most common detection element materials in various infrared detectors.[12] As an infrared detector, PbS functions as a photon detector, responding directly to the photons of ...
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Sorry to tell you a bad news. Classical methods, like titrimetry or gravimetry do not work very well at low concentrations. Failure would be a strong word. The beauty of EDTA is that it forms 1:1 complex with most metals. Pb is not an exception. If your lead solution is 0.000001 M, your EDTA solution should be 0.000001 M. Such concentrations are never ...
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Yes, do use household dilute H2O2 but, in general, do not use commercial percarbonates powders as they can contain an additive to convert the friendly percarbonate into more powerful Peracetic acid (PAA). The latter allows the otherwise weak bleaching power of the H2O2 to compete with chlorine-based bleaching. However, this comes at a price (safety). Hence, ...
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To identify a solvent:
1) If the components are in different phases, then component having same phase as that of solution is considered as solvent.
2) If both the components are in same phase, component having low b.p. is considered as solvent.
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The calculations made by Kent de los Reyes are based on the hypothesis that the solvent (water) does not produce any $OH^-$ ions. He calculates that the concentrations of $OH^-$ is about $10^{-12}$ in saturated solutions of $Tl(OH)_3$ or $Co(OH)_3$. But the real concentration of $OH^-$ is much higher, because of the water dissociation. If the saturated ...
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Water loses kinetic energy when it detaches from the surface of the ice
Water molecule in the solid phase have more hydrogen bonds than in the liquid phase. It takes energy to break these bonds, so only molecules with sufficient kinetic energy will detach (maybe because they just collided with a water molecule that transferred some kinetic energy to it). So ...
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Until now, I found the following.
Bustamante, P.; Peña, M. A.; Barra, J.: The modified extended Hansen method to determine partial solubility parameters of drugs containing a single hydrogen bonding group and their sodium derivatives: benzoic acid/Na and ibuprofen/Na. Int. J. Pharm. 194 (2000) 117-124:
benzoic acid/Na, ibuprofen/Na
Barra, J.; Peña, M. A.; ...
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There is very well-written chapter "The chemistry of monovalent copper in aqueous solutions" in Advances in inorganic chemistry, Volume 64 [1, pp. 220–223, DOI: 10.1016/B978-0-12-396462-5.00007-6] which extensively covers as to why copper(I) is unlikely to exist in aqueous solution and why nitrate is a poor ligand for the purpose of preserving monovalent ...
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As noted in the comments, $\ce{Cu+}$ is not stable in aqueous solution. It tends to disproportionate to $\ce{Cu^0}$ and $\ce{Cu^{2+}}$. You can therefore assume that all but a negligible amount turned into $\ce{Cu^{2+}}$.
Source:
Holleman/Wiberg: Lehrbuch der Anorganischen Chemie, de Gruyter, 101. edition, 1995.
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