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There is a strong reason. But it is the chemical reaction, not just a better dissolution. $\ce{Mg(OH)2}$ is a base with the limited solubility, defined by $K_\mathrm{sp}=[\ce{Mg^2+}][\ce{OH-}]^2$ $\ce{NH4+}$ ion, created by $\ce{NH4Cl}$ dissolution, acts as a weak acid: $$\ce{NH4+ + H2O <<=> NH3 + H3O+}$$ with $\mathrm{p}K_\mathrm{a}=9.25$ $\... 11 It turns out that the opposite of what happens in polar solvents takes place when a non-polar solvent is used. At the same temperature, o-nitrophenol is more soluble in benzene than it's m and p isomers. Sidgwick et al.1 did a study of this and obtained the following results. (Note that the solvent they used was toluene and not benzene, but they are similar ... 8 The distinction between physical and chemical change is actually not as cut and dried as one might think. Sure, what happens to copper in nitric acid is very different from what happens to sugar in water. But what about, say, the process that occurs when a steel strip is cleaned (basically removing carbon from the surface) in an atmosphere containing ... 7 Sure they can. Even if most of the ions are in ion pairs or other complexes, they can still move by being transferred from one complex to another. All you need is to get the ionic compound to dissolve, which in most cases requires a polar solvent with which the ions will not react. A rather unexpected case involves magnesocene, whose bonding is a mixture ... 7 The nitrophenols have completely different physical behavior based on the position of nitro group: $$\begin{array}{c|ccc} \hline \text{Compound} & \text{Melting point} & \text{Boiling point} & \text{Water solubility at } \pu{25 ^\circ C}\\ \hline \text{2-Nitrophenol} & \pu{43-45 ^\circ C} & \pu{215 ^\circ C} & \pu{2 g/L} \\ \text{... 6 Glutamine in aqueous solutions degrades slowly when stored in room temperature. Therefore, we can expect that rate of degradation is faster at higher temperatures. It is evident that the hydrolysis product of glutamine is 5-pyrrolidone-2-carboxylic acid (not glutaric acid as shown for enzymatic degradation) and ammonia (Ref.1): The Ref.1 states that: The ... 6 In order for molecules to precipitate out of solution, they need to aggregate together. Amino acids that have zero net charge can aggregate together much more easily than those that are charged. Molecules that have net charge need counterions to aggregate with them to offset the charges or the electrostatic repulsion will be too great. The concept of ... 5 I think we should not conflate "solubility" with "dissociation". They are different phenomena. The former has to do with a molecule breaking away from the other solute molecules and forming interactions with the solvent. The latter is a process which occurs after dissolution, whereby the solute molecules break apart to form ions, ... 5 The mixing of two compounds is a process which requires consideration of three types of interactions: solute-solvent, solvent-solvent and solute-solute. You make a good argument for the solute-solvent interaction being stronger in mixtures of butanamide and water, compared to butylamine and water. However, as you say, this argument alone would predict a ... 5 Start with a simple thought experiment: pour 100 mL of 1 M nickel (II) sulfate solution into a beaker and very carefully layer 100 mL of 0.01 M nickel (II) sulfate solution on top of the more concentrated layer. Then, even without convection or deliberate mixing, diffusion will, sooner or later, result in the solution having concentration of 0.55 M. In what ... 4 Wikipedia (primary reference) suggests a possible reason for \ce{Fe2(CO)9} dissolving preferentially in THF versus nonpolar solvents: it reacts according to the scheme \ce{Fe2(CO)9 + THF <=> Fe(CO)5 + Fe(CO)4 \cdot THF} Such a reaction is invoked to account for the dinuclear complex giving mononuclear products with various ligands in THF. A ... 4 Sodium hypochlorite (the active ingredient in bleach) and acetone react to produce chloroform. Chloroform can slowly decompose into phosgene in air with exposure to oxygen$$\ce{2CHCl3 + O2 ->[h\nu] 2COCl2 + 2HCl}$$but not in significant quantities under the conditions you describe. You did the right thing by opening the window for ventilation and ... 4 Solubility of calcium sulphate is quite high, compared to many much less soluble salts, like barium sulphate, calcium phosphate or calcium fluoride. Most of insoluble minerals are salts. Insoluble/limited solubility salts have as crystals lower Gibbs energy than dissolved, leading to solution being thermodynamically unfavourable. That is related to ... 4 There are no simple relation between the solubility and the solubility product when a doubly charged ion is involved. The measured solubility is always much bigger than the value obtained from the solubility product. This is due to the doubly charged ion. Doubly charged ions like \ce{Zn^{2+}} are usually hydrolyzed in water, and are partly transformed ... 4 I want to extend Maurice's comment: The amount of \ce{CO2} dissolved in water is proportional to the outer pressure. At \pu{20 °C}, 1 liter water dissolves about \pu{1.7 g} \ce{CO2} at normal pressure (1 atm). If the pressure is twice as large, the amount of dissolved \ce{CO2} is twice as much, \pu{3.4 g}. To talk about solubility of gases in ... 4 This paper here which describes the isolation of Glutamine from beets, contains a note that Glutamine is rapidly hydrolysed in water at high temperatures. The paper includes procedures in water at 60C. edit: This paper here describes the degradation kinetics of L-Glutamine in aqueous solution to 5-pyrrolidine-2-carboxylic acid. 3 \ce{CaF2} has a relatively low solubility in water, about 15 mg/L, with a K_{sp} of about 4\times 10^{-11}. According to the paper you linked, the concentrations of fluoride in tea are around 5 mg/L, which corresponds to about 1.3\times 10^{-4} M. Thus, the concentration of \ce{Ca^2+} required to precipitate the fluoride out is about 2.5 mM, which ... 3 Calcium carbonate will dissolve quickly in a strong acid, like hydrochloric, even if it is dilute and therefore not so dangerous. An organic acid has a problem in that it may have a relatively low pH at first, but after it dissolves a bit of the \ce{CaCO3}, a significant amount of the acid anion is formed in the solution, which stifles the ionization of ... 3 The problem could be solved with simultaneous equations, but the following are the wrong equations.$$K_{\mathrm{sp,}\ \ce{AgCl}} = [\ce{Ag+}][\ce{Cl-}]K_{\mathrm{sp,}\ \ce{AgBr}} = [\ce{Ag+}][\ce{Br-}]$$The right equations to use would be:$$K_{\mathrm{sp,}\ \ce{AgCl}} \ge [\ce{Ag+}][\ce{Cl-}]K_{\mathrm{sp,}\ \ce{AgBr}} \ge [\ce{Ag+}][\ce{... 3 Both chloride and bromide ions are present in 10-fold excess over silver ions. That means that the chloride and bromide concentration in solution will not drop by much (they will remain major species). The solubility product of AgCl is 200-times higher than that of AgBr. If both AgCl and AgBr precipitate, the chloride solution would be 200-times higher than ... 3 This is a well known problem in qualitative analysis. When you add$\ce{NH4Cl}$to a solution containing$\ce{OH-}$ions, you produce the reaction: $$\ce{NH4+ + OH- -> NH3 + H2O}$$ The result is that the concentration of$\ce{OH-}$decreases. If this operation was done in a saturated solution of$\ce{Mg(OH)2}$, the solubility product is no more obtained. ... 3 Apart from the fact noted by Mithoron that it's soluble in polar organic solvents, we may consider what lattice binding energy the solvation forces must overcome. In a compound like sodium chloride with compact, monatomic ions the lattice forces are too strong for all but the most aggressive solvents (but lithium salts have a way around this problem). With ... 3 The mistake you made is that you assumed mass concentrations of the ions from the stoichiometric relation, which is only applicable to the amount of substance, and, as a consequence, amount concentration. You have to convert mass concentration$γ$to the molar concentration c first: $$c(\ce{Pb(OH)2}) = \frac{γ(\ce{Pb(OH)2})}{M(\ce{Pb(OH)2})} = \frac{\pu{0.... 3 There are two ways to add two miscible solvents to each other. One is the bartender skill, whereby you attempt to create two different layers of solvents and then try to keep the flask motionless to prevent mixing. The other way is to mix them all along, whether by shaking, careless pouring, or a stirring bar. I will concentrate on the second as that is, in ... 3 Hansen Solubility Parameters (HSP) - is the short answer. Your observation has very much to do with the energy density of the solvents - with the idea that lesser the difference in the energy densities better is the ability of solvent to dissolve the solute. I am very poor at explaining things and not an expert in the matter at hand (please excuse me for ... 3 First of all, as MaxW pointed out in the above comment, your math seems to be wrong. From the the K_\mathrm{sp} value given in your question:$$s = \sqrt{K_\mathrm{sp}}= \sqrt{1.0 \times 10^{-12}}= \pu{1.0 \times 10^{-6} mol/L}= \pu{1.0 \times 10^{-6} mol/L} \times \frac{\pu{190.45 g}}{\pu{1 mol}}\times \frac{\pu{1000 mg}}{\pu{1 g}} \\ = \pu{0.190 mg/L}$$... 3 This is a simple logical and semantic problem, which is not a problem at all. Your professor is right and wrong- both at the same time. He is creating a classification which does not exist and which is meaningless. Look at the word origin of electrolyte: Etymology from OED: < electro- comb. form + ancient Greek λυτός that may be dissolved, soluble (... 3 Consider a solution of$\textrm{NaCl}$in water and another solution of$\textrm{AgNO}_3$in water. If you mix these two solutions,$\textrm{AgCl}$precipitates out. If you do not consider solubility as a chemical process/change, then it is very hard to explain this. A mere physical process should be no more than juxtaposition of four different kind of ions ... 3 Mercury cyanide behaves like Mercury chloride$\ce{HgCl_2}$. Both are soluble in water, and both do not dissociate in water. These compounds are not salts. The bonds$\ce{Hg-Cl}$or$\ce{Hg-CN}\$ are more covalent than ionic. When the atom Hg is included in compounds it often behaves as a non-metal. This is due to a relativistic effect. In all elements ...