25

Saturating a liquid with one solute does not mean that the liquid will no longer dissolve another solute. However, you can expect the solubility of the second solute to be different, generally lower, than in the neat solvent. One relevant concept here (though not specifically applicable to sucrose), in the case of ionic solutes, is the common-ion effect. ...


22

Honey is indeed a complex mixture containing more than hundred compounds. As for Wikipedia and depending on the point of view it is a supersatured liquid solution a viscous supercooled liquid (in the sense that it can get so viscous as to appear solid, without affecting its status of being a supersatured solution, and undergoes glass transition). https://...


21

In the comment to my previous answer, you asked for a theoretical reason for the solubilities, not considering energy data. Since I know from energy considerations that the issue is not the solvation of the anions, I can present a reason based on the strength of the ionic bond in the two compounds. This reference (as well as others) states the bonding in $\...


21

As it happens, the enthalpy of solution of $\ce{NaCl}$ in water (that is, the energy change associated with the dissolution of sodium chloride crystals in water) at standard conditions is very slightly positive, i.e., it is an endothermic process. At a constant temperature and pressure, these kinds of thermodynamic processes are dictated by the change in ...


21

You got me curious, so I poked around a bit on this. First there is insoluble and then there is insoluble. The Wikipedia article on capsaicin lists its solubility as $\pu{0.0013 g}/\pu{100 mL}$ which is $13$ parts per million. So capsaicin is "relatively insoluble", but not wholly so. Second the Wikipedia article also points out that capsaicin itself is ...


20

In short, for this group 2 hydroxide crystal, the dissolution and solvation process is an exothermic equilibrium. Increasing the temperature drives the reaction in the reverse, favouring crystallisation! However, it would be nice to rationalise some of these terms, so I will lay out a very brief overview of general solubility factors so you can understand ...


20

No, the correct way of putting it is $$\mathrm{Almost~all~of~the~\mathbf{inorganic}~nitrate~salts~are~soluble~in~water. }$$ The families of organic nitrate salts are typically nitrates of azoles and imidazoles. Some bright examples are (R) & (S)-miconazole nitrates, isoconazole nitrate and econazol nitrate. Econazol nitrate (Other names: Spectazole, EN) ...


19

Okay, for this scenario, the two concepts that play a key role are the lattice enthalpy and the hydration enthalpy. Lattice enthalpy is the energy released when one mole of a compound is formed from its constituent gaseous ions under standard conditions. Lattice enthalpy is dependent on the length of the ionic bond between the cation and the anion in the ...


17

Generally, the melting point of para isomer is quite higher than that of ortho or meta isomers. This is due to the fact that it has a symmetrical structure and therefore, its molecules can easily pack closely in the crystal lattice. As a result, intermolecular forces of attraction are stronger and therefore, greater energy is required to break its lattice ...


16

The reason for this is mainly because of a factor called Nucleation. Although it is not exactly obvious what is the source of the site for the nucleation to occur, what is clear is that there are sites present on ice-cream that are not present on ice alone. The carbon dioxide (CO$_2$) in the drink nucleates and forms bubbles (sometimes even a frothy foam) on ...


16

Soda contains $\ce{CO2}$. Over time the gaseous $\ce{CO2}$ comes to the surface. The ice cream increases the rate of this $\ce{CO2}$ from the soda causing increased fizzing. This increased rate is a result of the solvation of ice cream particles in the liquid. The particles serve as nucleation sites for the growth of gas bubbles. The bubbles get large quiet ...


16

As someone said here, this: The teacher stated that the ionic compounds dissolve in water except some carbonates. Is indeed an oversimplification. First of all, the distinction between an "ionic compound" to other compounds isn't too defined. What your teacher probably said, or didn't say but wanted to, is that some ionic compounds easily dissolve in ...


16

Ethanol and acetone are not non-polar organic solvents. Each one has a slight dipole moment; due to the difference of electronegativity between $\ce{H}$ and $\ce{O}$ in ethanol and between $\ce{C}$ and $\ce{O}$ in acetone. Wax is composed of heavy, long-chain alkanes. And as "Like dissolves like" try to dissolve your wax in toluene or in xylene.


16

Well, I went and searched for the solubility of oxygen in water and oil, and found this summary paper on the NIST web site: "The Solubility of Oxygen and Ozone in Liquids" by Battino, Rettich and Tominaga, J. Phys. Chem.Ref. Data., vol 12, no. 2, 1983. Conveniently, the paper gives solubility data for oxygen in both water and olive oil. The solubility is ...


16

The short answer is: yes, this is possible. Unfortunately, solubility is a fairly complex phenomenon to explain simply. Let's start with some examples where solubility is higher in a binary mixture than either solvent alone. For a solid-liquid-liquid example: phenanthrene-cyclohexane-diiodomethane.[1] For a liquid-liquid-liquid example like the one you were ...


16

Melting and dissolving are all the same when you look at mixtures close to saturation. You can say water lowers the melting point of the sugar, or that the solubility of sugar increases with temperature. Different description, same fact. What makes this seem different from e.g. a salt water solution is that the molten (i.e. non-crystalline) sugar is fully ...


15

For an ionic substance to dissolve in water, there are two competing factors that determine the enthalpy of solution $\Delta{H}_\mathrm{sol}$ which is the enthalpy (energy) change when a solute is dissolved in a solvent: The lattice energy (LE), the energy of formation of the crystal between infinitely separated ions. As LE is proportional to the charges ...


15

Apart from the methods, Ringo already described, you can do a few other tests. Aluminium This is loosely translated from the German chemgapedia.de. Look at the pretty pictures. Probably the easiest test you can do is reacting it with Morin in ethanoic acidic medium. It forms a yellow-green chelate complex, which has strong fluorescence under UV light. The ...


15

$\ce{SiCl_4}$ does not quite dissolve in water; rather, it reacts with water. So does $\ce{CCl_4}$, albeit extremely slowly, so for most purposes one may safely assume it doesn't. The difference is mostly due to the atomic size of C and Si. Smaller С atom is completely blocked off by four bulky chlorine atoms, so the water molecule can't reach it. Larger Si ...


14

Fluorocarbons are compounds that are hydrophobe and lipophobe. It's a special property of perfluorinated compounds: They are non-polar and thus hydrophobic but in addition as the high electronegativity of fluorine reduces the polarizability of the atom, fluorocarbons are only weakly susceptible to the fleeting dipoles that form the basis of the London ...


14

The enthalpy of solution going from the ideal gas state to the solution at infinite dilution is exothermic. However, more importantly, the Gibbs energy of solution is positive meaning that dissolution of methane into water is not a favored process. Most of this comes from the relatively large negative entropy change associated with the dissolution at room ...


14

I'm afraid this is rather a non-answer (or why is it so difficult to answer this)... Pretty much all nitrates are soluble. This is often explained by the exceptionally good delocalization of the negative charge. On the other hand, silver salts in general aren't well soluble (I recall only fluoride, nitrate and perchlorate as soluble. Sulfate, carbonate, ...


14

What is Sand? Scope of this Answer In order to give a good answer, we first have to decide on the type of sand we want to talk about, because its composition is highly dependent on that. Is it beach sand? If so, where from? All these questions should have been addressed in the question. I will limit the scope of this answer to pure silicium dioxide $\ce{...


13

Teacher stated that the ionic compounds dissolve in water except some carbonates. That is, at best, an oversimplification. Other ionic compounds such as silver sulfide are sparingly soluble in water. Note that this isn't a carbonate. And sodium hydrogen carbonate - $\ce{NaHCO3}$ - is soluble in water. Sodium carbonate ($\ce{Na2CO3}$) is also soluble in ...


12

In order to dissolve a salt, you have to break apart the ions and hydrate them in solution. You can use the enthalpies of hydration of the ions, and the crystal lattice energy of the solid, to predict which compounds will dissolve. I found data that the crystal lattice energy of $\ce{AgCl}$ is -916.3 kJ/mol (experimental), while the lattice energy of $\ce{...


12

Fluorous solvents have the odd property of being both hydrophobic and lipophobic (Ref) and thus are not miscible with either aqueous or many organic solvents as you've noted. Fluorine, as the most electronegative element, does odd things to a molecule (and is often used in pharmaceutical compounds just because of some of these odd properties). In the ...


12

Being a chemist, I know that paraffin would require a lipophilic solvent as pointed out by others. It occurred to me that ordinary vegetable oil is a cheap and of course innocuous lipophilic substance. I spilled about 2 oz molten candle wax on a marble surface and after it solidified removed a good portion by scraping with a plastic spoon. This was followed ...


12

Very simply, you explain the reason for this solubility rule by taking in consideration the energy requirements for the breaking of intermolecular forces between the molecules in the solute and the solvent. Note: this is only a simplified explanation as it also depends on other factors such as change in entropy Here is some background information on ...


12

This is not just some vapor pressure. This is the equilibrium vapor pressure. Thermodynamics is all about equilibrium, you know. And equilibrium, roughly speaking, is what takes place in a closed container after a billion years. Immiscible as they are, the liquids still have some solubility in each other (maybe extremely low, but anyway). Over the course of ...


11

Well, $\ce{Ag2O}$ is just a basic oxide. As such, it would dissolve in suitable acids ($\ce{HNO3}$ would do), but I guess that's not quite what you want. Well, some metal oxides ($\ce{ZnO}$, for instance) are amphoteric and thus would dissolve in $\ce{KOH}$ as well, via formation of hydroxo complexes. Sadly, this is not the case with $\ce{Ag2O}$. Then we ...


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