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3

For the case of magnesium ion, at least, it seems a stretch to say it dissolves in ammonia solution. Assume that the ammonia solution has $\mathrm{pH}$ of $11$ and magnesium hydroxide has a solubility product of $5.61×10^{-12}$. We infer that the magnesium ion solubility at equilibrium with the hydroxide in ammonia solution is only $\pu{5.6μM}$ which is not ...


1

Your polyethylene terephthalate (PET) plastic bottles must be too thin. The thickness of the plastic and the shape of the bottle both contribute to its resistance to bursting (pressure tolerance). According to SeattlePi: Most generally used two-liter PET bottles for soda bottling begin to fail at pressures around $\pu{10.34 bars}$ (1034 kilopascals), or $\...


-5

I'm quentin, I have your answer on the solubility of Ca3(PO4)2 : Ca3(PO4)2 = 3 Ca^2+ + 2 PO4^3- Ks = [Ca2+]3 [PO4^3-]2[Ca^2] = 3s and [PO4^2-] = 2s Ks = s^5.3^3.2^2 = 108.s^5 10^-26 = 108.s^5 so s = (10-26/108)1/5 = 2,5.10-6 mol.L^-1 the solubility of tricalcium phosphate is too low


1

I recommend sulfamic acid with formula $\ce{H3NSO3}$. The reaction is irreversible, no dangerous compound: $$\ce{2 H3NSO3 + CaCO3 -> Ca(H2NSO3)2 + CO2 + H2O}$$


0

$\ce{CaCO3}$ is sedimenting when mixed with water, and therefore is not easy to remove with hot water. Nevertheless several methods exist that actually dissolve it that are better than vinegar: You could add hydrochloric acid again to convert it into soluble Calcium chloride, according to equilibrium: $$\ce{CaCO3 (s) + 2 HCl (aq) <=> CaCl2 (aq) + ...


0

The boiling point is directly dependent on the various forces of attraction that the chemical species can exert. However, for solubility, it is the relative strength of these forces of attraction. Note that higher molecular weight species, in general, boil at higher temperatures due to greater magnitudes of van der Waals interactions. Boiling Points The ...


3

As andselsik pointed out, definetley no in the general sense. Just think of some practical examples like PTFE (Teflon) - which melts at about 330 °C but is definetly not water soluble. However there is one connection I'd like to point out that allows to relate the enthalpy of melting to the solubility: Consider phase equilibrium between a soluble compound ...


1

There are basically 3 effects. The solubility is controlled by the solubility product, the constant being the multiplication of ion activities $$K_\mathrm{sp}={a_\mathrm{M}}^m\cdot {a_\mathrm{X}}^n$$ for a salt $\ce{M_{m}X_{n}}$ For low concentration of ions, it can be approximated by concentrations, related to activities via activity coefficients $\...


2

Another option could be a ion selective electrode (ISE) selective for chlorides, but it would make sense rather for regular checks, as there would be high one time investment. The principle is the same as for the potential reference electrode ... $$\ce{Ag | AgCl | KCl}$$ ... but this time with mono or polycrystallic $\ce{AgCl}$ exposed to the solution. ...


3

The only well known method to check for (the absence of) chloride ions is precipitation as AgCl (solubility 2 mg/l at 20°C). Mercury(I) chloride is similarly badly soluble, but toxic (and unstable, disproportionation) and therefore bad practice. Lead(II) chloride is already far more soluble (4 g/l) than AgCl, and practically all other simple chlorides are ...


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