10

We can. But I see few reasons why it is not used: Iron is much cheaper than zinc. There can be remaining residue of iron/zinc, coated by copper, or just being excessive. While copper can be melted away and iron stays, zinc would melt together with copper, causing unwanted impurity (unless wanted for making brass alloys) If we remove copper for ...


8

Energy storage depends on the electromotive potential (i.e. difference between species in the electromotive series) and on the number of electrons available. Li, for example, has an oxidation potential of ~3.04 V relative to hydrogen, but Al has one of 1.66 V, so Li has the greater potential. On the other hand, Li has only one freely available outer ...


7

An atom is only neutral when viewed as a single object from large enough distance. But as an electron comes closer to the atom, it "notices" the electron cloud first. This cloud also "notices" the electron and deforms—the atom polarizes—so as to keep the atomic electrons farther on average from the extra electron, since like charges repel. But this ...


6

Let's write the half-cell equations in standard form as half-cell reductions. In a table of standard electrode potentials it would silly to write both the reduction and oxidation reactions since that would needlessly double the size of the table. $$ \begin{align} \ce{A+ + e- &-> A} &\quad V_\ce{A} \\ \ce{B^{2+} + 2e- &-> B} &\quad ...


6

Serial usage For the serial usage, one must avoid combination of cells of different kind and health. They should be as identical as possible. The best is using the same chemistry the same format the same vendor the same brand the same lot the same aging Otherwise the cells will have the different capacity and different discharge profiles under the load. ...


6

After some searching, I found this 1956 Thesis(Reference), which describes the production of peroxydisulfuric by electrolyzing concentrated sulfuric acid. The yield varied according to the concentration of the acid, temperature, current, and the nature of the electrode. In the Thesis, the chemist noted that discharged $\ce{HSO4-}$ united together to form ...


6

Who was responsible for this naming system and how can we change it? Michael Faraday was responsible for the terms anode and cathode more than hundred years ago. All the confusion regarding the nomenclature will vanish if you do not associate electrostatic signs with these two terms. One should identify the electrode labels with the redox processes rather ...


6

Can going to non-standard conditions reverse the cell potential of a voltaic cell? I have already answered that question, in the affirmative, for a simple tin and lead galvanic cell, here: https://chemistry.stackexchange.com/a/116734/79678. What follows is an elaboration with some specific values and illustrative figures. In fig. 1, the tin and lead ...


6

My understanding is that a redox couple is an unordered pair of two conjugate species This is conceptually perfect and there is no problem when we talk about electrode potentials of half cells because as I had mentioned in your earlier queries, the electrode potential value and its associated sign do not know nor care how you write the half cell. What is ...


5

I have done a significant amount of research over the past ten years to trace to origins of these electrochemical conventions and luckily got a chance to discuss these with some top electrochemists. I have been planning to write an article on this issue since it is a perpetual confusion. Basically, the origin of these "signs" issues originated in Germany and ...


5

For the acidic electrolysis, use the reactions where $\ce{H+}$ occurs. As $\ce{OH-}$ is not available in considerable amount there as a reagent, neither it is created as a product. Generally, for a reaction choice, apply the principle of availability and stability, allowing for a reagent to exist in (relative) abundance. $\ce{OH-}$ or anions of ...


5

This is pure scam! All he is doing is using an electrode which is corroding during electrolysis. The flocculated material you see in tap water is just a metallic hydroxide. If tap water had this much of metals, nobody will survive. Such a fraudulent salesman should be shooed away!


5

The typical gas sensor marketed to hobbyists and enthusiasts (except for $\ce{CO2}$) are of the metal oxide semiconductor type, not to be confused with MOS as in MOSFET. How do these work? In these detectors, a semiconductor material made of the oxide of a metal, usually tin. I'll quote from this question: ...sintered composite based on the ...


5

According to the definition used by IUPAC, the standard electrode potential $E^\circ$ of the standard hydrogen electrode is zero at all temperatures. For solutions in protic solvents, the universal reference electrode for which, under standard conditions, the standard electrode potential ($\ce{H+}/\ce{H2}$) is zero at all temperatures. The absolute ...


5

Moist air, rich in an electrolyte (salt particles) or human contact, providing both NaCl and H+ may supply the reagents needed for galvanic corrosion, with dissimilar metals in direct contact. This usually proceeds, albeit slowly, over time. Note, exposure to fruit juices could be especially problematic, resulting in a matter of days of continuous contact ...


4

The main thing to point out is that your question doesn't actually make sense the way you want it to. What does it mean if a reduction potential is zero? Or positive? Or negative? Is that favorable or not? It turns out we don't specify this at all! As an example, consider the standard hydrogen electrode and it's reduction potential: $$\ce{2H+ + 2e- -> ...


4

The voltage you measure between the terminals of a voltaic cell will depend on two factors: The intrinsic maximum voltage $(V_\mathrm{max} = E_\mathrm{cell})$ that the cell could produce, depending on the $E^o_{red}$ of each half cell, the ion concentrations and the temperature. This is calculated from the Nernst equation: $$E_\mathrm{cell} = E^⦵_\mathrm{...


4

According to my notes and many sources on the internet, electrons and cations both travel from the anode (A in the image) to the cathode (B in the image). The idea of the salt bridge is to prevent electrolytes mixing while providing ion flow. When you have a high concentration of inert ions in the salt bridge, cations in the salt bridge will flow into B, ...


4

Note that the actual potential for a particular redox reaction is not a fixed value, but depends on concentrations ( more exactly activities ) of reagents. The standard redox potentials are potentials with activities equal to 1, If we consider reactions $$\begin{align}\ce{ 2 H+ + 2e- &<=> H2 \\ 2 H2O + 2 e- &<=> 2 OH- + H2 \\ }\...


4

Generally, the electrode reactions are both based on lead in different oxidation states. At the negative electrode, $\ce{Pb}$ is oxidized to $\ce{Pb^2+}$ during discharge. $$\ce{Pb <=> Pb^2+ + 2 e-}$$ At the positive electrode, $\ce{Pb^4+}$ is reduced to $\ce{Pb^2+}$. $$\ce{Pb^4+ + 2 e+ <=> Pb^2+}$$ For a classical lead–acid battery, the overall ...


4

Many recommended procedures such as this one call for connecting (+) to (+), then (-) on the "good" battery to a metal component in the "dead" car engine. The negative terminal of the battery is grounded to the metal components, so we may think of the last connection as (-) on the dead car. If a spark were to form when the circuit is established on the ...


4

In electrochemistry, the rate of electrolysis depends on rate of charge entering the cell. Second most important point is that one can either control potential or current but not both during electrolysis i.e. you cannot have both values set. If you are fixing potential at 2.5 V, the value of current is not in your hand. I don't think Ohm's law remains valid ...


4

Because in the original question there is no elemental chlorine present. It says that you start from a solution of chloride and iodide, so both ions that can be oxidized to the corresponding halogen. And the rest of the answer is already given in your textbook example. Chlorine is a much stronger oxidizing agent. This means that it oxidizes others very well. ...


4

Your question is a valid question, and ignore downvotes. They don't mean anything. Your understanding is very good and that you realized that the electrode potential is a property of the electrode and it really does not care how the reaction is written. However, a equation is $needed$ to keep track of the electrons lost or gained in the Nernst equation. So ...


4

The electrochemical stability window is most important when considering components of an electrochemical system that you do not want to be oxidized or reduced. This refers most often to the electrolyte or protective coatings. For example, in lithium ion batteries it is highly desirable that the electrolyte does not react/change/degrade in any way as a ...


4

From your comments it seems that you are looking for an approximation of a log. I wish you clarified that in the main question without mentioning calculators. It seemed you just wanted to avoid a calculator for some unknown reasons. As Poutnik states, anyone who can post here, will certainly have access to computers and hence the ability to calculate logs. ...


3

I am sorry, this is not an answer but a comment, too long to be edited in the comments section. I don't agree with "straightforward" in the sentence : " It is relatively straightforward to understand that the more the diffusion layer grows, the shallower the concentration gradient gets and therefore that the current decreases. " This would be an ...


3

I still teach switching the sign. I find it easier to remember adding up reduction potential and oxidation potential. The half reactions are written as reduction in a table of reduction potentials, so it makes sense that you have to treat the oxidation half reaction differently. If the cell potential is calculated from reduction potential of the cathode ...


3

what actually happens inside the battery once it is fully discharged? Alkaline batteries use the exchange of electrons from zinc to manganese dioxide to produce electricity $$\begin{array}{rcl}\\ \ce{Zn (s) +2 OH- (aq)}& \ce{->}& \ce{ZnO (s) + H2O (l) + 2 e-}\qquad \text{(anode)}\\ \text{(high potential anode)}\quad \ce{2e-}&\ce{->[load][\...


3

You are right with copper does not react with aluminium ions. But as a side reaction, copper may get slowly oxidized by oxygen and dissolve in mildly acidic solution of aluminium salt. Such a thing may happen, if you decalcify copper heating spiral by vinegar and let it stay overnight. $$\ce{ 2 Cu + 4 CH3COOH + O2 -> 2 (CH3COO)2Cu + 2H2O }$$ ...


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