# Tag Info

6

It is true that acid-base indicators ($\mathrm{pH}$-indicators) are either weak acid or weak base. What I understood reading your question is that you have wrong impression about bases. To clear your view, not all bases contain $\ce{OH-}$ ions. Specially, most organic bases are weak and contain electronegative ion (e.g., $\ce{N}$ or $\ce{O}$) with at least ...

5

$$\ce{NaOH + HCl -> NaCl + H2O}$$ Let $x$ Litres be the amount of $\ce{NaOH}$ needed. $$-\log[\ce{H+}] = 2$$ $$-\log\left(\frac{(0.5 \times 0.02 - x \times 0.5)}{(0.02 + x)}\right) = 2$$ $$x = \frac{49}{2550}$$ $$x \approx 0.0192 \; \text{(3 sig. fig.)}$$ Thus, $\pu{0.0192 L}$ of $\ce{NaOH}$ is needed.

5

A basic indicator is usually a Broensted-Lawry base ( accepting protons ) rather than Arrhenius base ( releasing hydroxide ions ): $$\ce{B + H2O <=> BH+ + OH-}$$ or $$\ce{BH+ + H2O <=> B + H3O+}$$ depending on if the indicator is used in its base form or conjugate acid form (the latter is usually more stable and more soluble). Regardless of ...

4

You got the solubility part reversed. The solubility of $\ce{AgCl}$ is lower than the solubility of $\ce{Ag2CrO4}:$ $$s(\ce{AgCl}) = \sqrt{K_\mathrm{sp}(\ce{AgCl})} = \sqrt{\pu{1.8E-10 mol2 L-2}} = \pu{1.34E-5 mol L-1}$$ $$s(\ce{Ag2CrO4}) = \sqrt[3]{\frac{K_\mathrm{sp}(\ce{Ag2CrO4})}{4}} = \sqrt[3]{\frac{\pu{1.1E-12 mol3 L-3}}{4}} = \pu{6.50E-5 mol L-1}$$ ...

4

Let's suppose we want to titrate a solution containing an unknown monoprotic and weak acid. We use a strong base, such as $\ce{NaOH}$. When the number (and moles) of hydroxide ions is equal to the amount of hydronium ions, here we have the equivalence point. The equivalence point is, when the molar amount of the spent hydroxide is equal the molar amount ...

4

Your question: Given a sample of an alcoholic beverage (be it beer, wine, etc..), what are possible methods to find out the ABV (alcohol by volume) of it with moderate accuracy? There are a few method used in winery industry as Karl mentioned a comment in elsewhere. I assume you want to find out better methods to find the alcohol content: The best ...

4

Your title and the main body of the post are two different questions. Anyway, it is important to understand the difference in each case. Let us talk about the analyte. Your analyte can come in two forms: (i) either as a solid or (ii) as a solution. In the case of a solid analyte you would quantitatively transfer a known exact weight of the sample to a ...

4

My question is, since are looking at the equivalence point in a titration, why can't we use $N_\text{acid} \times \text{Volume}_\text{acid} = N_\text{base} \times \text{Volume}_\text{base}$, since equivalents of acids and bases are the same at the equivalence point? If we can, how would we do it? First of all, I would suggest that you stick to the original ...

4

To understand what is happening with the two sets of titration curves, start by considering the titration of acetic acid ("HOAc" for short) with sodium hydroxide solution. The titration curves are shown below: In the figure, the mutual concentrations are simply the concentrations of the reactants. Thus, 0.1 M HOAc is titrated with 0.1 M NaOH, 0.01 M HOAc is ...

3

Citric acid has $\mathrm{p}K_\mathrm{a}$-values of 3.1, 4.7, and 6.4, while those of tartaric acid are 3.0 and 4.3. If you adjust the $\mathrm{pH}$ to 6.4, tartaric acid would be roughly 100% deprotonated, while 50% of citric acid still has one proton to give off. If you titrate this solution with $\ce{NaOH},$ you could estimate the buffer capacity, giving ...

3

I don't think your memory is serving you right. That is why we write everything in the notebook, especially color changes. I think you are doing distribution experiments where iodine is distributed between aqueous layer and an organic layer. When we add indicator for titration, it is not a solid starch but starch which is boiled in water. So when you added ...

3

Given: HClO4= 100.5, d HClO4=1.8 (ACO)2O=102, d (ACO)2O=1.1 H2O=18, d H2O=1 A) Preparation of 100 ml perchloric acid 0.1 Molar aqueous solution. I had to do some poking around to understand the density number. Without a temperature the density of liquids is nebulous. Wikipedia lists a density of 1.768 g/ml which seems to be for the 72.5% azeotrope. So ...

3

This practice is done in analytical chemistry in order to minimize the relative weighing error on the balance. Preferring a larger formula weight for a primary standard has nothing to do with impurity levels. We have to start with the highest purity standard. For example, you wish to prepare a 0.010 M solution of oxalic acid dihydrate in 1 L flask. Its ...

3

In this case, cerium(IV) sulfate is a strong oxidising agent, hence, using $\ce{HCl}$ would likely be oxidised to form $\ce{Cl}$ of higher oxidation state. Therefore, the titre would be vastly higher than expected. $\ce{HNO3},$ on the other hand, is an oxidising agent and would oxidise $\ce{Fe}$ to a higher oxidation state than expected (of $\ce{Fe^3+}$). ...

3

This is not an attempt to answer the question(s). However, it is seemingly clear that OP is doing a lab practical, and I'd like to clear his assumption about density of $\ce{HClO4}$ acid. The series of density ($d$) of $\ce{HClO4}$ acid by percentage at $\pu{25 ^\circ C}$ have been published in 1941 (Ref.1): $$\begin{array}{cccc} \% \ \ce{HClO4} \ (w/w) ... 3 It is due to the fact that at half equivalence point, the pH of the solution is equal to the \ce{pK_a} value of the weak acid. And this pH does not depend on the initial concentration of the acid. You should take into account something that does not appear on your diagrams. The concentration of the strong base (used on the abscissa) is not given ! They are ... 3 It is not quite correct of the statement that "ascorbic acid cannot be determined by Beer Lambert's law, because it does not absorb visible light." Although the second part of the statement is true, the highlighted part is incorrect since Beer Lambert's law applies for ultra-violet range of the spectrum. According to Ref.1, ascorbic acid (Vitamin C)... 2 In simplistic terms, Arrhenius acids are molecules while Arrhenius bases are ionic compounds. Much like \ce{BaCl2} – another ionic compound – it is a good first rule of thumb to assume it completely dissociates into its respective ions: one \ce{Ba^2+} and two single-charge anions. You could say that barium were never really connected to either hydroxide; ... 2 You cannot do simple volumetric redox titration to determine iron content because the amount of iron is very very small in a single fruit (on the order of fraction of a milligram). Imagine what would be the buret reading? Classical methods are good for large concentrations >> 1% wt/wt With such small quantities, UV-Vis absorption spectroscopy or atomic ... 2 It is an interesting question. Daniel Harris is revising his book with my former mentor. Hope he clarifies this section in the revised version. Your point number 1 is misleading. The reason is that before the titration, theoretically there is no Fe(III). So Nernst equation should not be used- electrode potential is infinite (log 0 is undefined). When you ... 2 One way to determine the amount of vitamin C in food is to use a redox titration. The redox reaction is better than an acid-base titration since there are additional acids in a juice, but few of them interfere with the oxidation of ascorbic acid by iodine. Iodine is relatively insoluble, but this can be improved by complexing the iodine with iodide to form ... 2 From the comments: [OP] Okay Thank you all I finally got it. so K is just thrown out the window. and we are left with the equation (F- + H2O <-> FH + OH-) So i take the Ka to make Kb by (Kw/Ka) and my concentration of F- and OH- was (0.02/0.3) which I plugged into sqrt((Kb)*(M of OH- or F-)) took the -log of that got the pOH and then converted it back ... 2 There are a few reasons. Some of the most important are: -Oxalic acid dihydrate isn't hygroscopic (it doesn't absorb water from the atmosphere, whereas sulfuric acid does). If a substance can absorb water from the atmosphere, then when you weigh it out, you won't be able to accurately calculate the number of moles used because you don't know how much is ... 2 The suitability of indicators is given not only by their range, but also by the direction of the transition. Generally, the transition is always much better from the light colour to the more intense colour, as the equivalence point and it's approaching is much easier to track. So for MO (red 3.1- yellow 4.4), it is toward acidic, for PHPH (8.2 clear -10.0 ... 2 The n- factor must be obviously wrong because two electrons are involved in the reduction of peroxide to oxygen. Write the half cell yourself. Indian chemistry textbooks should have pity on the world and stop teaching these equivalents to young students. This is a thing of the 18th century. Those who invented this concept moved on. 2 Thanks, Martin, for pointing out the relevant paper with all the answers! Basic procedure This titration was first described by Mohamed Zakibarakat, Mohamed Fathy Abd El-Wahab, and Mohamed Mahmoudel-Sadrin in 1954 (doi.org/10.1021/ac60100a013) as an alternative to using 2,6-dichlorophenol-indophenol as titrant. Here is the procedure for a solution of pure ... 2 An indicator does not affect a particular reaction. But in your solution, you have two successive reactions when adding HCl to this solution : First$$\ce{CO_3^{2-} + H^+ -> HCO_3^-} and a given indicator must be added to determine the end of this reaction. If you don't, all you see is a colorless solution being transformed into another colorless ...

2

This is a note to M. Farooq's answer: This titration consists of two back to back redox reactions. As Farooq pointed out, your first reductive half reaction is actually $\ce{IO3- <=> I2}$ (or $\ce{IO3- <=> I3-}$). In either way, generated $\ce{I2}$ in the presence of extra $\ce{I-}$ would become $\ce{I3-}$. The presence of $\ce{I3-}$ can be ...

2

The indicators are never in the form "InOH" with a covalence between "In" and OH, producing $\ce{OH^-}$ ions in water. If a molecule contains one OH group attached to a Carbon atom, it would be an acid, an alcohol, en enol or a phenol. These sorts of molecules are never releasing $\ce{OH^-}$ ions in water. On the contrary, they are weak acids.

1

Measure the volume or mass (depending on the titration approach you are using, you can do it by mass) of undiluted honey that you are titrating. After that, provided that you don't splash any out or loose some by transferring to another container, you can add as much distilled water as you like. It will make no difference to the titration result. Yes, the pH ...

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