22

It turns out that 2 identical snowflakes have been observed, but... Two Identical Snowflakes Although when we think "snowflake" we usually picture an object with 6-fold radial symmetry, snowflakes actually come in many different shapes (reference 1). Many sites report that in 1988 an NCAR researcher found 2 identical snowflakes of the hollow columnar ...


14

Short story Maybe it will help to think first about a ridiculous case. Say you have a large hydrophilic solid, and a small drop of water is added on top. We would not expect the entire solid to disappear into the liquid. Yes, there will be some adsorption and solubility here, yet most of the solid will remain undisturbed. In some sense, the solid 'does not ...


12

One of the most difficult parts of chemistry is learning to recognize a type of reaction based solely on its reactants. This has to be done before you can apply a reaction pattern to the problem, and so it is critical to get this step right in the beginning. In this case, you are looking at the reaction: $$ \ce{AgNO3 + HCl -> AgCl + HNO3} $$ You ...


12

Carbon dioxide from the atmosphere (or from decaying matter in the ocean) reacts with water to form carbonic acid: $$ \ce{ CO2 + H2O -> H_2CO_3 }$$ and this reacts with calcium ions to form calcium carbonate: $$ \ce{H2CO3 + Ca^{2+} -> CaCO3 + 2H+ }$$ The solubility of calcium carbonate is about $13~\mathrm{mg \over L}$, so if the concentration of ...


11

Here's one easy way : Make a solution of salt ($\ce{NaCl}$) in water, and another solution of Silver Nitrate ($\ce{AgNO_3}$), which you can buy easily) in water. Upon mixing the two, $\ce{AgCl}$ (Silver chloride, a white compound) will instantly precipitate.


10

Why Do Carbonates, Oxides, and Pure Metals Precipitate Before Metal Sulfides? They don't. I cannot get to the particular paper you read, but I can provide both a logical and a quantitative argument that the metal sulfides precipitate first. The thermodynamics of this situation do not involve reduction potentials so much as simple solubilities. Metal ...


9

The two words are often used interchangeably (perhaps incorrectly), though there is a slight nuance. IUPAC defines the following: Crystallisation: The formation of a crystalline solid from a solution, melt vapour or a different solid phase, generally by the lowering of the temperature or by evaporation of a solvent. Precipitation: The sedimentation ...


8

Although you ask specifically about forming precipitates your question is also about the solubility of one substance in another. The simplest answer as to why precipitates form is because the free energy $G$ is minimised when some solid, its solute and solvent exist in equilibrium compared to the situation of only solute in a solvent or only solid and ...


7

Actually, the textbook should remove this question. You should discuss this with your teacher. It will be a service to future students. The heavier elements of group II elements form insoluble sulfates, phosphates and carbonates. Now carbonates display another interesting phenomenon. If you keep bubbling carbon dioxide in into a group II carbonate solution (...


6

The rule of thumb for hydroxide salts is actually that all hydroxides are insoluble save those of the alkali metals and the heavier alkaline earth metals (namely $\ce{Ca^{2+}}$, $\ce{Sr}^{2+}$, and $\ce{Ba}^{2+}$). There are many resources you can look up to check the solubility rules of certain ions, this one, for example. Before researching this, I ...


6

Is it possible to obtain pure, precipitated iron with no oxidation by some chemical process? Very pure iron can be obtained by pyrolysis of iron pentacarbonyl. Iron metal obtained in this manner is in fact called carbonyl iron and is commercially available. $$\ce{Fe(CO)5 -> Fe(s) + 5 CO}$$ Doing this yourself is very dangerous, however, because iron ...


6

I don't have time to do the calculations manually, but inputing your problem data into a chemical equilibrium software (here Dozzaqueux) reveals that the precipitate you observe is Pb(OH)2, while no PbI2 should form. The former is white, while the later is yellow; can you tell us what color was your precipitate?


6

It's probably an experimental error. The 'barely' may be important here. For one, you may not be sure how accurate your initial molarities are. A mistake in one of them ruins it all. Titrate them properly and check. Titration usually gives two significant digits. Also, since its a cube overall, a small mistake can get much larger. You may want more ...


6

Barium nitrate has a water solubility of $\pu{10.5g/100mL}$ at $\pu{25^oC}$. It isn't specified in the question what concentration of sulfate you suspect might be present, but given that you are trying to check for sulfate by precipitating with barium, barium nitrate should be the way to go.


6

Your choices are restrained as the precipitation of $\ce{SO4^{2-}}$ in $\ce{BaSO4}$ is the classical way to quantify the former and an electrochemical determination (in aqueous solution) is not practical. Electing $\ce{Ba(OH)2}$ may lead to the formation of silver hydroxyde, equally poorly soluble in water. $\ce{Ba(PO3)2}$ itself is very poorly soluble, as ...


6

This must have been a very good question for a geochemist. They'd give you a complete answer with details of why $\ce{Fe2(CO3)3}$ does not but $\ce{Fe2O3}$ exist in our surrounding. Yet, I would like to give only a simple answer for the question: No, $\ce{Fe2(CO3)3}$ does not form (in sort) when you added $\ce{Na2(CO3)3}$ to an aqueous solution of $\ce{FeCl3}...


6

$\ce{AgCl}$ ppts before $\ce{Ag2CrO4}$. But in order to visually detect the $\ce{Ag2CrO4}$ you must add a slight excess of $\ce{Ag^+}$. Hence you use a "blank" with no $\ce{Cl^-}$ to account for the amount of $\ce{Ag^+}$ needed to get enough of the $\ce{Ag2CrO4}$ ppt to form so that it can be detected visually.


6

No, absolutely not. Precipitation reactions can be either endothermic and exothermic. Table 1. Thermodynamic data of precipitation for some salts \begin{array}{cccccc} \hline \text{Salt} & \Delta G_\text{ppt}^\circ & \Delta H_\mathrm{ppt}^\circ & -T\Delta S_\mathrm{ppt}^\circ(\pu{25 °C}) \\ \hline \ce{Be(OH)2} & -121 & -31 & -90 \\ \...


5

Precipitation is usually a single or double displacement reaction, and refers to the formation of an insoluble, solid salt from soluble ions. For instance, silver bromide, $\ce{AgBr}$, is insoluble in water. However, it is made up of $\ce{Ag+}$ and $\ce{Br-}$ ions, which we can introduce in the form of soluble salts: $$\ce{LiBr(aq) + AgNO3(aq) -> LiNO3(...


5

The ions combine while in solution, and since the resulting compound is not soluble, it precipitates. This happens so quickly because as the precipitate forms, it takes those individual ions out of solution, thus pulling the reaction toward the products side, according to Le Chatelier's principle.


5

Let's speculate a bit ;) Colestyramine is a twodimensional polymer network, build from styrene and and 1,4-divinylbenzene. The polymer has a lot of benzyltrimethylammonium groups. That means it can act as an anion exchanger. It is used to strongly bind bile acids Neither the polymer itself, nor the adduct with bile acid are supposed to degrade in the bowel....


5

With $\ce{(NH4)2HPO4}$, $\ce{Mg^2+}$ forms a white, crystalline precipitate of magnesium ammonium phosphate hexahydrate: $$\ce{Mg^2+ + HPO4^2- + NH4^+ + OH- + 5H2O \rightleftharpoons Mg(NH4)PO4 \cdot 6H2O\downarrow}$$ This is a very sensitive test for $\ce{Mg^2+}$; however, calcium, strontium, barium and other heavy metal ions also form precipitates of ...


5

No, recall the definition of a precipitate "insoluble ionic compound". Take this reaction into consideration $$\ce{ NaOH~(s) + HCl~(l) <=> H2O~(l) + NaCl~(aq)}$$ none of the products are insoluble. You'll learn more when you go over gas evolution reactions.


5

This procedure is done in our laboratory class, but first you have to add sodium hydroxide until the lemon juice is somewhat alkaline. Then add the calcium chloride (use $1~\mathrm{ml}$ for every $2~\mathrm{ml}$ of the lemon juice + $\ce{NaOH}$ solution). I think the calcium chloride would first react with the hydroxide, producing calcium hydroxide, which ...


5

As noted in your question and in one of the comments, copper forms many different complexes having a variety of colors from red to green to blue to black and probably more. Of course the situation here is even more "complex" (sorry) as you can end up with mixtures of the different copper complexes, as appears to be the case over the course of your ...


5

Sulfuric acid is an excellent dessicant and is also very non-volatile itself (it boils at 337 °C). In other words it would likely still have been there in yet another month. Your instructor was correct to suggest that you rinse the sulfuric acid with some cold water. Benzoic acid does have some non-negligible solubility in water (1.7 g/L at ...


4

The question is most likely badly worded and should have been What is the maximum mass of $\ce{AgCl}$ that can be dissolved? As we all know, $\ce{AgCl}$ is damn insouble in aquaeous solutions — and due to the equilibrium constant it doesn’t get better at all if you add chloride ions (reducing the silver ion concentration). So without having done the ...


4

Different solutes in the same solution indeed can affect each others' solubility. One example where this effect can be used to advantage is the "salting out" of a protein from water solution. Adding a salt (ammonium sulfate is especially good for this purpose) reduces the solubility of the protein in water, causing it to precipitate or crystallize out. ...


4

Sure... here's a an article on it and a patent for the process. In brief, "[Donald Sadoway, the John F. Elliott Professor of Materials Chemistry at MIT] found that a process called molten oxide electrolysis could... [produce] steel as a byproduct... Sadoway's method used an iridium anode... But after more research... the MIT team identified an inexpensive ...


Only top voted, non community-wiki answers of a minimum length are eligible