New answers tagged acid-base
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After washing by water, get rid of the residual water with methanol or acetone.
The residue of solvent will finally evaporate, leaving much less water than without this step. Similar procedure is used in preparation chemistry to remove water from crystaline products.
Ethanol is usable too, but there is usually some 4-5% of water.
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A rough calculation for the pH of 0.2-ᴍ formic acid gives:
$$ \mathrm{pH} = 1/2 (\mathrm{p}K_\mathrm{a} - \log(0.2)) = 1/2 (3.74 + 0.70) = 2.22 $$
You can check by calculating the equilibrium constant from the concentrations of all the species, and the estimate is pretty good.
The pH of the buffer is 3.44 using the Henderson Hasselbalch expression (see other ...
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Basically, you are on the right track and answered the question. First, NaNH$_2$ is used as a base to deprotonate the terminal alkyne (see this Reagent Friday). Second, the generated nucleophile opens the epoxide in an S$_\mathrm{N}$2 fashion, usually adding to the lesser substituted carbon which in your case yields a quaternary alcohol after acidic workup (...
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Hydrogen and chlorine atoms of $\ce{HCl}$ in gaseous state are covalently bound and is termed hydrogen chloride. When this gas is bubbled into water, it ionizes completely to give $\ce{H3O+}$ (free proton + water molecule) and $\ce{Cl-}$ ions and becomes an acid solution which is termed hydrochloric acid. Even in gaseous $\ce{HCl}$, the charge is not ...
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No bond is completely ionic or completely covalent, Fajan's rule. Also $HCl$ is predominantly a covalent compound rather than ionic.
Shouldn't, H+Cl- be a salt since hydrogen is positive and Cl is negative?
Sure, $\ce{HCl}$ does dissociate into $\ce{H+}$ and $\ce{Cl-}$ in a polar solvent, but that doesn't mean $\ce{HCl}$ is simply $\ce{H+}$ and $\ce{Cl-}$. ...
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Interesting question. You successfully found the components but forgot the solvent(s)? Do you remember that pH is defined for aqueous solutions only. All these indicator molecules are large relatively hydrophobic organic molecules so they are not water soluble. Thymol blue and phenolphthalein are such examples. One needs to add alcohol to dissolve these ...
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For a reaction to happen, it is sufficient if 1 in a 1000 thiamine structures are deprotonated. For a species with a $\mathrm{p}K_\mathrm{a}$ of 18 to be 50% deprotonated, you would need an aqueous solution with a pH of 18, which is not possible as far as I know.
So even if thiamine is present in aqueous solution and not (as mentioned in Rafael L's answer) ...
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An alternative to @Waylander 's explanation. First, we brominate cinnamic acid to form cinnamic acid dibromide and because of the base $\ce{Na2CO3}$ the cinnamic acid is deprotonated, and thereby both CO2 and Br are eliminated leading to the formation of bromostyrene and then KOH does its work.
Note: The answer may seem similar as of user55119 but here the ...
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It is the molarity of total acetic acid/acetate content, so the 3rd option.
It is quite general principle applied in the pH buffer context. You can have a wide rage of $\ce{pH}$ values with the particular buffer type, with the same total molarity of "active substance", like citrate, phosphate or mixed buffers, being adjusted by variable amount of ...
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The pKa of meta-cresol suggests that it is slightly less acidic than phenol. As commented, the difference is quite subtle. You could invoke the greater electronegativity of the carbon in methyl as compared to hydrogen, you can neglect any hyperconjugation, and you can simply dismiss it. What, really, is the significance of that tiny difference in pKa? You ...
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Put simple: these are two mutually independent definitions.
The Bronsted concept is based on the possibility to donate a proton (i.e., an acid), or to accept a proton (i.e., a base). But it isn't only the presence/absence of a proton in a compound, it equally depends on the reaction partner if a molecule may act as an acid, or as a base.
The Lewis concept ...
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