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15

The best known conducting aqueous solutions are that of strong acids in water because the hydronium ion (=protonated water) has the highest electrical conductivity known today. The infinite dilution conductivity of of hydronium ion are compared below from the Book Chapter "Proton Transfer Reactions and Kinetics in Water" by Stillingerenter. The ...


12

You are most likely getting an inaccurate value. As anticipated in Poutnik's comment above is difficult to reliably measure a $\mathrm{pH}$ of very acidic solution but the $\mathrm{pH}$ scale is indeed open, and negative $\mathrm{pH}$ are a real thing. It is easy and more meaningful to state a very high concentration or activity rather than reporting ...


11

"Hydronium" content can be easily determined by simple acid base titration. Keep in mind that there is nothing fundamental about pH. It is just a convenient scale for expressing hydrogen ion concentration. For example, a kilogram is a recent invention for expressing mass measurement, it does not mean that ancient people did not weigh anything ...


9

Because the acid and conjugate bases are equimolar - and because the equilibrium constant is small, obviating the need to solve a quadratic if you approached this in another way - the Henderson-Hasselbalch equation is: $$\pu{pH} = {\rm p}K_\mathrm{a} + {\rm log}_{10}\left({[A^-]\over[HA]}\right)$$ $$\pu{pH} = 4.74 + {\rm log}_{10}\left({0.2\ {\rm M}\over 0.2\...


9

The conductivity might be viewed relative to other electrolytes and by concentration. By concentration: Up to about 25% $\ce{H2SO4:H2O}$, conductivity increases fairly linearly with concentration, reaches a maximum at ~35%, and then drops precipitately (no pun intended)! And sulfuric acid has such "affinity" for water that oleum, or fuming ...


7

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 ...


7

The strength of permanganic acid that you quote, combined with that of potassium hydroxide as a base, would guarantee that pure potassium permanganate is neutral in aqueous solution. But commercially prepared potassium permanganate is made in the presence of alkali, the use of potassium instead of sodium arising from the fact that the reaction scheme does ...


6

The main component of egg white, ovalbumin, contains 93 negatively charged residues (Asp + Glu) and 68 positively charged residues (Arg + Lys), as well as 14 His. If you make the solvent less polar by adding ethanol to water, they will tend to form the neutral species (Arg+ and Lys+ will lose a proton, and Glu- and Asp- will pick up a proton). If you had the ...


5

The purpose of this $c^\circ$ is to ensure dimensional consistency. Keep in mind that Wikipedia can be edited by anyone, although the content is often of very good quality, sometimes the volunteer writer misses some points and assumes that the reader might be aware of his/her symbols. In your link, the writer does not explicitly define $c^\circ$. The German ...


5

Most proteins are polyelectrolytes because the constituent aminoacids may have ionizable groups, typically carboxylic and basic (imino, guanidino) groups on the sidechains, in addition to the terminating amino and carboxylic groups (when not modified). The degree of ionization depends on the fold of the protein and the pH of the solution. Denaturation of a ...


5

Reused my answer to how-to-numerically-model-a-phosphoric-acid-titration-curve, edited for diprotic cases: First, calculate the common denominator $CD$: $$a_1 = [\ce{H+}]^2$$ $$a_2 = [\ce{H+}] \cdot K_\mathrm{a1} = a_1 \cdot \frac {K_\mathrm{a1}}{[\ce{H+}]}$$ $$a_3 = K_\mathrm{a1} \cdot K_\mathrm{a2} = a_2 \cdot \frac {K_\mathrm{a2}}{[\ce{H+}]}$$ $$CD = a_1 +...


4

Take four test tubes with sample, add bromothymol blue in two of them, methyl red in the rest. Add a few drops of HCl to one of the BTB tubes, and a few drops of NaOH in one of those with methyl red. One of the tubes will have a noticable change of colour.


4

Let's consider an aqueous solution, the concentration of which is $\pu{1 M}$ in $\ce{NH3}$ and $\pu{1 M}$ in $\ce{NaOH}$. Thus, following equilibrium would be taken place: $$\ce{NH3 (aq) + H2O <=> NH4+ (aq) + OH- (aq)}\tag1$$ $$\ce{NaOH (aq) -> Na+ (aq) + OH- (aq)}\tag2$$ The $\mathrm{p}K_\mathrm{b}$ of equilibrium $(1)$ is 4.75, thus $K_\mathrm{b} =...


4

There is an easy way to do what OP wants, assuming OP wants to prepare solutions at $\pu{25 ^\circ C}$. So, OP can prepare saturated solution of $\ce{Mg(OH)2}$ solution: $$\ce{Mg(OH)2_{(s)} <=> Mg^2+_{(aq)} + 2OH-_{(aq)}}$$ Since $K_\mathrm{sp}$ of $\ce{Mg(OH)2}$ is $\pu{5.61 \times 10^{-12} M3}$, you can find the solubility of $\ce{Mg(OH)2}$ at $\pu{...


4

If $c$ is the concentration of $\ce{Mg^2+}$ in the $\ce{Mg(OH)2}$ solution, the concentration $[\ce{OH-}] = 2c.$ At $\mathrm{pH}~9,$ $[\ce{OH-}] = \pu{1E-5 M},$ then $c = \pu{5E-6 M}.$ So you have to dissolve $\pu{1.25 μmol}$ $\ce{Mg(OH)2}$ in $\pu{250 mL}$ water. This is $\pu{71.3 μg}$ of $\ce{Mg(OH)2}.$ This is difficult to do in practice, as such a small ...


4

Most people find alkaline taste bitter and "bad tasting". There are some exceptions; shellfish, particularly conch, may have a pH of more than 8.0. This is also the approximate pH of bile, which tastes particularly foul when retching. Crackers made with baking powder or baking soda also may be somewhat alkaline, as is egg yolk. However, your ...


4

Aeration does not seem to be the way to go. If there are no other ions in the water, electrolysis (Kolbe reaction) would convert the $\ce{HOAc}$ to ethane and carbon dioxide: $$\ce{2CH3CO2^- -> 2e^- + 2CO2 + H3C-CH3}$$ This process could probably be arranged to be done in a pipe with electrodes on the sides, so that the water was treated as it was used, ...


4

I like what Maurice said "When nobody understands a scientific phenomena, we give it a name. Here the name is overpotential. Here the hydrogen is said to have a big "over potential" on mercury cathode. That is a bright and remarkable way to hide our ignorance." This is little on the extreme side but people have spent their life on ...


4

Recall that $\mathrm p\ce{OH} = -\log_{10}\space [\ce{OH-}]$ Where $\ce{[OH-]}$ is the concentration of hydroxide ions. Suppose a solution has $10^{-7}$ moles of $\ce{[OH-]}$ ions. Then this implies $\mathrm p\ce{OH} = -\log_{10}\space [\ce{OH-}]$ and $\mathrm p\ce{OH} = -\log_{10}\space [10^{-7}]$ which is equal to $7$. Suppose a solution has $10^{-4}$ ...


4

$ \begin{align} (n_{\ce{NaOH}})_i= \pu{0.2 mmol}\\ (n_{\ce{HCl}})_i= \pu{0.1 mmol}\\ (n_{\ce{NaH2PO4}})_i= \pu{0.1 mmol}\\ (n_{\ce{Na3PO4}})_i= \pu{0.05 mmol} \end{align} $ At first the $\ce{NaOH}$ will react with $\ce{HCl}$ as per the following reaction: $$ \begin{align} \begin{array}{ccccc} \ce{NaOH} & + & \ce{HCl} & \ce{->} & \ce{NaCl} ...


3

Poutnik has given an excellent answer to your question: Why $\ce{NaHCO3}$ is able to lower the $\mathrm{pH}$ and what is the reaction that occur when I add it? It is also important that your curiosity shown in the title: Why is sodium bicarbonate added to lower the $\mathrm{pH}$? Since I did not know of your synthetic procedure, I guess this is for your ...


3

This answer addresses the original interpretation of the question "Why $\ce{NaHCO3}$ ?“ rather than "Why to lower pH ?“. For the latter, see the Mathew's answer. The reason of using $\ce{NaHCO3}$ is the reaction: $$\ce{HCO3- + OH- <=> CO3^2- + H2O}$$ First, near all hydroxide is converted to carbonate, and then an excess of bicarbonate ...


3

Acetic acid does not form an azeotrop with water and is less volatile than water. So aeration of water would do the opposite - enriching of water by acetic acid due preferred evaporation of water. Reverse osmosis should help, perhaps after neutralization to be mostly in acetate form. The question is, if it is worthy the troubles. The cheaper, easier and ...


3

You could simply calculate the $\mathrm{pH}$ of the solution. $K_\mathrm{a{_1}}$ of $\ce{H2CO3}$ is $\pu{4.30E-7}$ and $K_\mathrm{a{_2}}$ of $\ce{H2CO3}$ is $\pu{4.72E-11}$.The approximate $\mathrm{pH}$ of the solution can be given by: $$\mathrm{pH}=\frac{(\mathrm{p}K_\mathrm{a{_1}} + \mathrm{p}K_\mathrm{a{_2}})}{2}$$ This yields $\mathrm{pH}= 8.347$, which ...


3

$\ce{Mg(OH)2}$ is a strong base since it is ionic in nature; it usually dissociates completely and so it's degree of dissociation is one. For weaker salts, the concentration values you assigned for the ions, $x$ and $2x$ respectively, do depend on the molarity of the $\ce{Mg(OH)2}.$ You have to define $x$ in the terms of its degree of dissociation $(\alpha),$...


3

By using a indicator, you cannot achieve $\mathrm{pH} \ \pm 0.1$ accuracy. To get that accuracy, as Karl suggested, you must buy a precision $\mathrm{pH}$ meter (you still has to calibrate it before use). The best indicator to get fairly accurate reading within the $\mathrm{pH}$ you are interested is bromothymol blue indicator. The following is a display of ...


3

You cannot "predict" any error by thinking that if we calibrate with two buffers the error will be x, and if we calibrate with three buffers the error will be y. Ideally, in each case, the pH of the sample should be identical. Good quality pH meters tell you how close they are to the theoretical slope in terms of percentage. It should be close to ...


3

The equation $$\mathrm{p}K_\mathrm{a} + \mathrm{p}K_\mathrm{b} = \mathrm{p}K_\mathrm{w}\tag{1}$$ has two degrees of freedom, so two values are independent on each other and the third one depends on the other two. Reaction equilibrium constant $K_\mathrm{a}$ for $$\ce{HA + H2O <=> H3O+ + A-}\tag{R1}$$ is chemically independent on $K_\mathrm{w}$ for ...


3

The background of the topic is, that for the function $$y = \frac {a(T)}{x}$$ $y$ does depend on $T$, but $x$ does not. $\mathrm{pH}$ of basic solutions is temperature dependent via $\mathrm{p}K_\mathrm{w}$ temperature dependance. Additionally, $\mathrm{pH}$ of weak acid/base solutions is temperature dependent via $\mathrm{p}K_\mathrm{a}$ or $\mathrm{p}K_\...


3

You did not mention what kind of waste water are you talking about. Is it from a textile industry, paper mill, leather factory? Waste water is not a "compound", it can be any junk and its pH can vary over several orders of magnitude! As I have stated before, there is nothing fundamental about pH. It is a scale of convenience. It can be negative and ...


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