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4

Besides steric factors related to the small size of the transition metal core, we could be seeing an electronic effect described in this answer. Iron is fairly early in the transition metal series, so when pushed beyond the $+3$ oxidation state it has few $d$ electrons in the central core. As explained in the referenced answer, this makes the iron strongly ...

3

Following the @Nilay Ghosh recommendation, I found an SDS sheet for $\ce{Cu2[HgI4]}$ and there is SDS for $\ce{HgI2}$. For $\ce{Cu2[HgI4]}$: Acute toxicity estimate Oral (Expert judgment): 5,1 mg/kg Acute toxicity estimate Inhalation (Expert judgment) 4 h: 0,051 mg/l Acute toxicity estimate Dermal (Expert judgment): 5,1 mg/kg For $\ce{HgI2}$: LD50 oral ...

1

,Assuming glycinate is in its anionic form ($\ce{H2N-CH2-CO2^-}$), one thing to consider is it may not act just as a simple anionic ligand like chloride or hydroxide. You can count off four atoms from the nitrogen to either carboxyl-ate oxygen and those atoms both have available electron pairs. So you add the metal atom and you get a five-membered ring ...

0

In solutions more basic than pH 9.6 glycine mainly exists as the carboxylate anion you describe. Between 9.6 and 2.34 it exists mainly in the zwitterion form. titration curve from here

1

The anion $\ce{AuCl4−}$ is the most known complex of gold(III). It is very stable complex and have very high formation constant ($K_f$). That is the reason the metal chloride complex have a lower reduction potential than metal aqua ion. As pointed out in the other answer, this is a consequence of Nernst's law. Let's look at the two redox equations we ...

1

From your question, I suppose you do not know about crystal field theory. To properly understand how and why certain complexes are octahedral and other ones are tetrahedral you need to fully understand this theory. You can find information in Greenwood and Earnshaw, Chemistry of the Elements. A more detailed info about the theory is in Theory of Groups in ...

8

The authoritative source Nomenclature of Inorganic Chemistry, IUPAC Recommendations 2005 (Red Book) lists $\ce{H[AuCl4]}$ as an example of a salt in the subsection IR-4.4.3.4 Generalized salt formulae [1, pp. 61–62]. Further, introduction to the section IR-8 Inorganic Acids and Derivatives underlines that IUPAC nomenclature is established from composition ...

-1

This difference of redox potentials is a consequence of Nernst's law. The second equation is nothing else as the first one, where the concentration of the $\ce{Au^{3+}}$ in a solution of $\ce{AuCl4^{-}}$ is extremely low. This sort of comparison could be used to calculate the equilibrium constant of the equilibrium $$\ce{Au^{3+} + 4 Cl- <-> AuCl4^{-}}$$...

3

What is meant by a "very strong" Be−O bond? If berylium's tendency to hold on to water ligands is unusually strong, is it due to its small ionic size? You got that absolutely right. $\ce{Be\bond{-}O}$ is a strong bond because of the small size of Be. Smaller cation size means a stronger pull on the $\ce{O}$ electrons, thus reducing the bond length ...

2

I'm sure OP get the answer by reading all comments and the answer elsewhere. However, I'd like to point out few things on this spectrophotometric determination of $\ce{Fe^2+}$ in drinking water (it is not distilled water): Iron can exist in natural waters where it originates from the dissolution of minerals. Even though most iron containing minerals are not ...

1

The o-phenanthrolin is extraordinarily sensitive to the smallest amount of $\ce{Fe^{2+}}$. A concentration as low as of $2$ microgramms $\ce{Fe^{2+}}$ per liter is still colored and its concentration can be measurable by optical absorption. As a comparison, a concentration of $\pu{30g/L}$ $\ce{Fe^{2+}}$ is the minimum measurable by Beer's law in water.

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