Wikipedia:Reference desk/Archives/Science/2010 November 1

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November 1

PETA video

In this video of trappers killing their fur-bearing animals, aren't the animals being either killed instantaneously or being knocked unconscious? It just seems to me much ado about nothing -- sort of like fooling people with statistics that they don't understand. Sure, it looks spooky to see these raccoons kicking a bit, but that happens after severing the carotids/trachea of a chicken, too? DRosenbach ^{(Talk | Contribs)} 01:30, 1 November 2010 (UTC)

Our article on the animal in the video is at Raccoon Dog. -- 119.31.126.67 (talk) 13:07, 1 November 2010 (UTC)

In a sense, the video is a purely objective demonstration. It's up to individuals to decide whether this activity is ethical or justifiable. In terms of scientific analysis, it is probable that the raccoon is capable of feeling pain; I think the consensus is that a quick blow to the head is among the less painful lethal techniques. See pain in animals for starters. As far as the movements - I can't tell from the videos whether we're seeing a reflex, rigor mortis, or conscious squirming - in my opinion, it is not possible to discern conclusively from these videos. Nimur (talk) 01:43, 1 November 2010 (UTC)

Just watched the first few seconds of the video. I saw one animal being struck five times over perhaps a five to ten second period. Obviously not "instantaneous." I certainly wouldn't enjoy it if i were the raccoon. HiLo48 (talk) 01:51, 1 November 2010 (UTC)

(edit conflict) (haven't watched the video btw)

PETA's position is:

"If anybody wonders 'what's this with all these reforms?', you can hear us clearly. Our goal is total animal liberation, and the day when everyone believes that animals are not ours to eat, not ours to wear, not ours to experiment [on], and not ours for entertainment or any other exploitive purpose."

according to co-founder Ingrid Newkirk.

How "well" they may be euthanized to harvest their fur does not change the fact that they are being raised in captivity and then killed solely to harvest their fur. Some consider that to be unethical, or feel a sense of empathy for the animals. Others do not. But the objection for some of those who do is not that they are being tortured to death (which presumably everyone would object to), but just that "we" are ill-using them for a clearly unnecessary and purely vanity-driven purpose. WikiDao ☯ (talk) 01:57, 1 November 2010 (UTC)

Just watched the video. I have no idea how anyone could not have a problem with that. What exactly is the point of your "question", DRosenbach? Could you please clarify exactly what it is you would like to know? Thank you. WikiDao ☯ (talk) 02:27, 1 November 2010 (UTC)

Why should he clarify his question? He states it very clearly. He wants to know if this video depicts cruelty on a still-conscious animal. I don't think there's a way to know, but personally it looks to me that the raccoon was completely 'out' after the first blow. If true, the video shows no more cruelty than a slaughterhouse, probably less cruelty than many types of fishing.

In any case, while they sometimes do good investigative work, their motives are not to educate, and I absolutely would not trust a video being distributed by PETA without third party verification. APL (talk) 03:30, 1 November 2010 (UTC)

Did you watch the entire video? If the animal was "out" after the first blow, how do you explain the squirming? Humans certainly don't move their limbs around frantically after being knocked out.

I think DRosenbach explained it in the question - personally I've seen a goat beheaded (also in China as it happens) - so it was definately dead - but it kept 'spasming' similar to what was seen in the video - we probably have an article on the phenonoma - but I don't know what it is called.Sf5xeplus (talk) 05:54, 1 November 2010 (UTC)

No. I'll admit that I only watched about half of it. Does something dramatically different happen in the second half? I don't have a reference handy, but I'm pretty sure that humans do sometimes twitch after crushing blows to the head. Hollywood doesn't portray it that way for reasons that should be obvious after watching the video. (Notable exception : "The Rock").

If you want to feel bad for those animals, feel bad that they grew up on a farm, probably in cages. Not because they died from a sudden blow to the head. APL (talk) 15:34, 1 November 2010 (UTC)

Also, in the middle of the video one person asks the cameraman, "are you recording?" He goes on to say, "make sure this video doesn't leak; that wouldn't be good for us." I highly suspect that, even in China, what these people are doing is illegal. --99.237.232.254 (talk) 05:44, 1 November 2010 (UTC)

Not necessarily. Farming animals is in many ways an unpleasant process. Anyone who's not a complete fool understands that it would be bad PR for the naive masses to be made to watch a heavily edited video of the most unpleasant aspects of your operation.

I imagine someone working in an entirely legal slaughterhouse for entirely delicious beef would have a similar reaction to a video camera. APL (talk) 15:34, 1 November 2010 (UTC)

DRosenbach, your error is attempting to apply logic to what is intended as a shock video. APL (talk) 03:33, 1 November 2010 (UTC)

• This question is asking for opinion ("seems to me it's much ado about nothing - is it?") and should be removed. (Note that my response above is the only one that even makes an attempt to link to an article or source that might be relevant, and even still does not answer the "question" because there is no way to answer the question "encyclopedically"). WikiDao ☯ (talk) 10:11, 1 November 2010 (UTC)

The question is clear and not opinion: "are the animals being either killed instantaneously or being knocked unconscious?" Yes or no? As illustration the OP refers to chickens. 92.24.189.164 (talk) 13:46, 1 November 2010 (UTC)

WikiDao -- please cool your jets. What constitutes animal cruelty itself is an opinion, so there's certainly many ways to spin my question. But my intent was not to create a windstorm of emotion (as perhaps you are suggesting I intended to do)...assuming one ascribes to a normative view of animal cruelty -- in that the killing of an animal is not inherently uncruel -- does this video depict cruelty at all, because it sure seems a lot more sensational than anything else. I didn't watch the entire video either (only the first 2/3) but it was quite boring and I just couldn't finish. The head/neck-banging of the first 2/3 doesn't seem cruel at all. DRosenbach ^{(Talk | Contribs)} 01:15, 2 November 2010 (UTC)

I'm fairly certain that the aim is to polarise opinion - ie those that are opposed will be appalled and some that were not opposed will become opposed, additionally, those that were blissfully ignorant may be appalled. It shouldn't come as a surprise that PETA characterise it as cruelty -since that's where they are coming from - so I doubt they post it for fun - but to attempt to bring what's happening into public view.

I won't comment on the cruelty or lack thereof, but will note that killing animals larger than birds is usually a two step process: first you stun them, then you kill them. Traditionally the stunning is done exactly as these Chinese guys are doing: you hit them on the head with something heavy. More modern techniques include the captive bolt pistol and electrocution. The killing is usually done be exsanguination, but these guys are doing it by inflicting massive head trauma. The style of blow they used in some cases, where they hold the animal upside down and whack it at the base of the skull, is often used in killing rabbits in one step. Having killed many things (in what I hope is a cruelty-minimizing way) I believe the animal has some brain function but nothing like its normal thoughts after a well-delivered first blow. --Sean 20:15, 1 November 2010 (UTC)

See rabbit punch. Googlemeister (talk) 13:16, 2 November 2010 (UTC)

In this case, as I noted above, there appeared to be five "first blows" right at the start of the video. Is this just evidence of incompetence? HiLo48 (talk) 20:53, 1 November 2010 (UTC)

I figured apathy; suggestive, though, of such action not being cruel as a rule. DRosenbach ^{(Talk | Contribs)} 01:59, 2 November 2010 (UTC)

I'm not sure if they have much choice - the first step, as noted above is to stun or knock out the animal - but because these animals are for fur, cutting the throat is not really an option (blood stains on fur, damage to the pelt etc) - so hitting it numerous times on the head seems a reasonable thing to do if selling pelts is something you do for a living. Probably less cruel than stunning it, then stripping the pelt and waiting for the animal to die of a massive stress related heart attack when it comes round and discovers it's got no skin anymore.. An unpleasent business alround - but I don't see that the killing of the animals was more cruel than any other obvious way of killing it. (Is there a better way?)

actually according to http://features.peta.org/ChineseFurFarms/ some are still alive when being skinned - as the video shows. It's a different video, and is about 100x worse than the first video posted - even the strong of stomach may not want to watch it.Sf5xeplus (talk) 08:00, 2 November 2010 (UTC)

I figure these guys just don't care, just like I don't care about a live oyster's opinion on sliding down my throat. I personally would care with these raccoon things, but Chinese slaughter operations are a rougher place, I suppose. --Sean 16:35, 2 November 2010 (UTC)

Also, exsanguination is important for meat quality, which is (perhaps!) not an issue here. --Sean 16:35, 2 November 2010 (UTC)

is HCl in 95% saturated salt solution still a strong acid?

I assume that a saturation limit exists for salt in water because at some point there aren't enough water molecules to separate the Na+ and Cl- ions, which are growing more acidic/basic for each other, respectively.

This makes me curious as to to acid-dissociation properties in highly saline water. Does the acidic species in solution change in highly salty water?

The pKa for HCl is given to be somewhere around -7, but hydronium is usually listed as -1.8 or something, but I assume the latter is for a dilute, distilled solution. I wonder if I can make a "more acidic hydronium species" by near-saturating the solution with neutral salts. So I have two questions:

1) Will Cl- be a stronger base in near-saturated salt solution 2) Will H3O+ be more acidic in near-saturated salt solution? Will Zundel cations and Eigen cations decompose in saturated salt solutions to form the more acidic "bare H3O+" species?

John Riemann Soong (talk) 02:09, 1 November 2010 (UTC)

Generally, the common ion effect is only a major issue in equilibrium situations. The dissociation of hydrogen chloride in water is highly extensive process. To put it bluntly, the chloride ion has such a terrible affinity for the hydrogen ion, I would doubt that you could get any discrete HCl molecules in a water solution regardless of how much you pump it full of chloride ions. --Jayron32 02:22, 1 November 2010 (UTC)

I still expect the HCl to dissolve more or less completely, but what I expect is for HCl's pKa to increase (to maybe -3) and for hydronium's pKa to decrease. Isn't hydronium pKa = HCl pKa in a solution of saturated 37% HCl? That's why as water evaporates, HCl gas evaporates from it as well right? So these are my hypotheses:

• As long as hydronium pKa > HCl pKa, you can dissolve more HCl in the water.
• If hydronium pKa < HCl pKa, then the protonated water will spontaneously reprotonate the chloride, forming HCl gas.
• I have noticed that concentrated HCl behaves a lot differently in reactions than dilute HCl, and not simply because the reaction rate increases, but because the acidic species seems to increase in strength.
• Dissolving NaCl (or perhaps some other neutral salt) in water decreases the pKa of hydronium because the hydronium has to compete with the salt for solvation stabilisation.
• At the same time the Cl- conjugate base is less-stabilised, and is so more basic.

I guess for clarity I should state that this is not an attempt to use the common ion effect but actually solvation effects. We could dissolve something else, like lithium fluoride. I notice that HCl is a weaker acid than say, perchloric acid. There are some expensive (spectroscopic, crazy test compounds) way to test this, but I wonder if there are cheaper, more elegant ways of playing with the acid strengths of what are normally considered "strong acids". Perchloric acid forms stable solid hydrates because I suspect that hydronium is never a stronger acid than perchloric acid is. However in saturated conditions I quite suspect that at some point Cl- becomes a stronger base than water. This is not simply because there is more Cl- around (we could dissolve sodium nitrate, say) -- it's because the basicity strength of Cl- actually changes. John Riemann Soong (talk) 02:44, 1 November 2010 (UTC)

I am unclear on what you are asking then; are you asking if you can create a solution which will have measurable quantities of discrete HCl molecules in a water-based solution? I suppose its possible, if you take enough water out of the solution through solvation effects with a highly soluble solute, you could do something like that. However, lithium fluoride wouldn't work as it would generate additional OH- which would throw everything off (the flouride ion is a weak base; since HF is a weak acid). --Jayron32 02:58, 1 November 2010 (UTC)

Is there such a thing as solvated HCl -- the solubility of which I suppose is very low -- before it turns into gas? What I'm trying to ask is firstly whether I can create a stronger HCl solution even though I am far from having 38% HCl say. Let's say I have 3.8% HCl and I saturate the rest of the solution with something like lithium bromide or something. Will I have the strength of 38% HCl (in terms of the pKa of the dominant acidic species in solution), just 10 times more dilute?

Because I note that 3.8% HCl (no salt) does not behave like 38% HCl that is ten times more diluted. Fischer esterification runs much more than 10 times more slowly with conc. HCl that has been 10X diluted. (Granted this is in the glacial carboxylic acid and not water, so we maybe talking about an RCOOH2+ species instead, but...) I believe 38% HCl will form carbocations (if the alkene is water-soluble -- take pyran maybe), but dilute HCl will not. I wonder if I can get the strength of concentrated HCl without needing concentrated HCl. The concentration will be much less of course, resulting in a slower reaction, but the strength will be comparable. John Riemann Soong (talk) 03:16, 1 November 2010 (UTC)

AH. But it has nothing to do with the HCl per se in these reactions, its the hydronium ion (acidity) which does the relevent work. In other words, its not the pKa, it is the pH which makes a difference here! HCl tends to be used in these cases because other strong acids have their own problems, for example H2SO4 tends to cause unwanted deydration problems (if there are any hydroxides present in the target molecules) while HNO3 can both nitrate molecules AND oxidize them. HCl has the advantage that a) it is a strong acid, so can produce very low pH's, and b) the Cl- ion is pretty much inert, and is unlikely to cause unwanted side reactions, which can occur with other strong acids. It's even better than HBr or HI, which, due to the softer nature of the Br- and I- ions, tend to be better nucleophiles, so they have their own problems with side reactions. HCl is primarily used where you need a very low pH (high hydronium/H+/acid/whatever concentration) and you don't want side reactions. --Jayron32 03:23, 1 November 2010 (UTC)

There is something more than a pH effect here I think. I read somewhere on another professor's page (confirming my suspicion) that the pH scale is really only effective down to pH=0. Lower than that, you can talk about a hydrogen ion concentration, but the H3O+ species involved becomes markedly different. At least that was what I perceived him to be saying. That is, between pKa 0 and 14, H3O+ has around the same strength (pKa = -1.8). At higher concentrations, the pKa of H3O+ seems to drop. However, no one has confirmed this for me -- in saturated HCl solution, does the pKa of molecular HCl equal the pKa of hydronium species? This would make sense to me, because if not, more molecular HCl would continue to dissociate.

My other suspicion is that there is no way to add enough HCl until solution has a pH of -7. That would be asking for a concentration of 10^7 M, which seems impossible. Thus it seems logically necessary that either the pKa of HCl must rise or the H3O+ species to drop in pKa for HCl fumes to start appearing. John Riemann Soong (talk) 03:38, 1 November 2010 (UTC)

Where did a pH of -7 come from - you dropped that in from nowhere? For HCl fumes see below about the solubility constant of HCl in water - it has practically nothing to do with pKa.Sf5xeplus (talk) 03:52, 1 November 2010 (UTC)

It is common knowledge? You can find it at hydrogen chloride. As long as HCl can spontaneously protonate H2O, HCl reformation would seem to be minimal. In order for HCl species to start appearing, H3O+ would need to reprotonate Cl-. John Riemann Soong (talk) 04:03, 1 November 2010 (UTC)

Please - note the difference between pKa and pH - conc. HCl does not have a pH of -7 , HCl has a pKa of -7 .

If the pKa of hydronium did not change to match the pKa of HCl more closely, then hydronium would have a pretty hard time protonating Cl- to form HCl gas. Thus in the "no pKa change" situation, HCl gas would start bubbling out of solution at absurdly high H+ concentrations. I assume that as long as HCl is more acidic than effective acidity of hydronium, free HCl will dissociate. Fumes start appearing when free HCl starts appearing in solution. (How much of conc. HCl is dissolved free HCl? If you lower the temperature or increase the pressure of HCl above, I don't suppose more HCl dissociates, merely dissolves?) John Riemann Soong (talk) 04:40, 1 November 2010 (UTC)

HCl gas would start bubbling out of solution at absurdly high H+ concentrations yes that's what happens

When HCl solutions are saturated the vapour pressure of HCl(g) above the solution is 1 atm (under standard conditions) you should be able to do the maths yourself (you've got pKa and the molar concentration of a conc. solution) - ask in another question if you get stuck.Sf5xeplus (talk) 04:54, 1 November 2010 (UTC)

I don't know the solubility of HCl(aq). What is the makeup of the HCl in 38% RT HCl -- 95% dissociated HCl, 5% free HCl? My question anyhow is that HCl gas would start bubbling out of solution way too late if H3O+ did not become more acidic. The pH of conc. HCl is low, but not that low. John Riemann Soong (talk) 05:04, 1 November 2010 (UTC)

conc HCl is 38% w/w, that's saturated - it doesn't really get more concentrated. Combine that with the pKa and the equilibriium reactions and you have enough information to work it out for yourself.Sf5xeplus (talk) 05:22, 1 November 2010 (UTC)

(edit conflict) I think you need to start again. There's problem at point 1 : dissolution of HCl has little to do with pKa - you can dissolve gases in liquids and them not dissociate .. in a polar solvent like water polar gases dissolve well eg NH3 and HCl. Ionisation of the dissolved gas happens, much more so in HCl - but isn't required for dissolution. ie in conc. HCl(aq) there are free HCl molecules.

The other but - if I've understood correctly is the suggestion of adding (saturated) NaCl to a HCl solution - this is a Buffer solution, and ignoring activity effects you can easily calculate the pH using the equations in the article.

As for Cl- increasing water acidity - this isn't going to happen - Cl- is a base (weak).

The meat of the question seems to be "does saturated NaCl" alter the proportion of different types of hydronium ions - in water "H3O+" doesn't really exist - all positive dipoles will be strongly solvated (H3O+ can exist in a compound as you know) - in water (H+).(H2O)(n) is a more "truthy" description where n is undefined, but greater than 1. see Hydronium#Solvation. 87.102.115.141 (talk) 03:10, 1 November 2010 (UTC)

I believe Cl- can increase water acidity not because Cl- is acting as a base but because water has to form shells to hydrate the Cl- (granted the effect is stronger when hydrating Na+). We can ignore the common ion effect altogether as that was not my intention. NaCl is simply a cheap salt in the laboratory.

But yes altering the strength of the H3O+ species is the meat of my question. As more salt is dissolved, H3O+ becomes destabilised. I suppose the Zundel and Eigen cations still exist, just more fleetingly -- and the extensive hydrogen bonding and "charge sharing" becomes more difficult. For ease of discussion I will simply talk about the "average H+ species" in water -- I am not talking about bare H3O+. But as more salt dissolves in water, I suppose the average H+ species approaches bare H3O+ that is increasingly less and less stabilised. (That is, lone pairs still solvate it, just not as well.)

My assumption with HCl was that HCl gas isn't very soluble in water if it doesn't dissociate -- after all, that is the situation with carbon dioxide! (CO2 has a fair amount of charge separation -- but does the lack of a dipole moment hurt solvation?) I suppose what I should say that is when HCl species exist in concentrated HCl soln (either as a fuming gas or solvated species), hydronium pKa = HCl pKa, yes? In dilute solution, the pKa difference is large, but as you add more and more HCl, the pKa gap closes. John Riemann Soong (talk) 03:27, 1 November 2010 (UTC)

HCl molecules are very soluble in water, just like NH3 and MeOH molecules.. Of course total lack of a dipole in CO2 affect solubility.

I can't answer the rest - your talking gibberish. Or so it seems.

The constants you want are the solubitity constant for HCl in water, and the dissociation constant for HCl. They're called constants for a reason - they're constant. For variations the concept of Activity (chemistry) is used to explain deviations. Sf5xeplus (talk) 03:37, 1 November 2010 (UTC)

In class we always talk about dilute solutions, and dilute solutions is all that we ever analysed. I am intensely curious about the dynamics of saturated solutions, since they are easily achievable in practice and they seem to result in effects that go beyond mere concentration increases.

I'm talking about the pKa gap between H3O+ (in all possible forms) and HCl. In dilute solution the difference is ~5, but in saturated HCl solution wouldn't this gap be zero?

AFAIK the dissociation constant for HCl is not actually "constant" -- there are very large concentration effects. It is constant only if the free energy change is constant. This is true in dilute conditions, but not in near-saturated conditions, yes? John Riemann Soong (talk) 03:48, 1 November 2010 (UTC)

No. pKa is a measure (-log) of the equilibrium constant which is constant - please read Activity (chemistry). It's the activity which is considered to change.

The reactions are:

HCl(g) <<<<>>>> HCl(aq) (solubility constant)

HCl(aq) <<<<>>>> H+(aq) + Cl-(aq) (acid dissociation constant)

The acid dissociation can aid the dissolution by reducing the amount of free HCl in solution.

The amount of HCl that you can get in solution depends on both these equlibriums.Sf5xeplus (talk) 04:01, 1 November 2010 (UTC)

I think you may be going wrong because you are considering Cl- ions in isolation - no such thing exists in a solution - there must be a Na+ ion for each Cl- ion when using NaCl - the counterpoint to your half of the argument is that Na+ will increase the basicity of the solution by stabilising OH- !! - does that make sense?Sf5xeplus (talk) 03:46, 1 November 2010 (UTC)

I am not considering anything in isolation. I didn't even want to talk about the common ion effect... John Riemann Soong (talk) 03:48, 1 November 2010 (UTC)

I don't really care about the identities of the specific counterions. All that matters is that it is a neutral salt, cannot form a buffer in dilute conditions, and that it requires lots of H3O+ molecules to solvate them. I am asking -- can I add lots of NaCl to a dilute HCl solution to decrease the pKa of the hydronium species? The pH will not change but will the solution become more acidic? John Riemann Soong (talk) 03:58, 1 November 2010 (UTC)

No. It's the opposite - the Cl- will suppress the ionisation of HCl (as per common ion effect and as per Le Chatelier's principle) - even though Cl- is a weak base. You can actually calculate the magnitude of the effect because HCl/NaCl is a buffer solution - albeit a weak effect. The answer is no. absolutely no. it will not work.

The pKa can vary - according to the solvent eg different pKas in different solvent... and adding NaCL changes the solvent.. but the effect is to make it less acid. If you add enough Cl- you can start to get HCl₂- which doesn't help.Sf5xeplus (talk) 04:14, 1 November 2010 (UTC)

(undent). Look, John, consider what you get when you dissolve HCl in water at any concentration. There are only a few possible species:

• Discrete HCl molecules
• The acid species (whether you consider this H+, hydronium, extra-solvated hydronium, etc.)
• The chloride ion
• The hydroxide ion (ignore this, its concentration is far too low to be useful)
• Water molecules

Forget the pKa stuff for a moment. What you need to know is which of these species are actually involved, either as reactants or as catalysts, in the chemical reaction you are trying to do. Once you know that, worry about how to tweak the relative concentration of those things. The entire pKa thing you seem to be on seems to be a completely unrelated tangent, and its steering you away from your potential solution to your problem. pKa may end up relevent later, but lets put first things first: What are you trying to do, and what species from the HCl solution participate in what you are trying to do? --Jayron32 03:56, 1 November 2010 (UTC)

I need a reagent with a more potent Hammett acidity function, essentially. I need a stronger acidic species in terms of the acidity of the H+ species in water, and not merely H+ concentration (pH). Suppose the only starting acid I have is HCl, and other weak organic acids, if need be. If I am trying to form carbocations (or hell protonate an aromatic compound that has EDGs) it would seem not only to be pH-dependent, but dependent on the maximum acidity of the H3O+ species in water. In dilute conditions this is -1.8, yes? But in concentrated conditions this might be lower.

I suppose my main interest is decreasing the pKa of H3O+ in water below -1.8. This is what my question resolves around. I believe this is quite reasonable. The harder thing to do I suppose is measure this pKa decrease.

I suppose there is literature on this topic, but I don't know how this phenomenon would be searched for. John Riemann Soong (talk) 04:11, 1 November 2010 (UTC)

If all you have is HCl, you are unlikely to make any significant changes to the Hammet acidity function with anything you can do in this way. The effects you note may actually happen, but will probably only make smallish changes to the aciditiy of HCl. What you need, if you want a stronger acid is a, um, stronger acid, like a superacid. If Conc. HCl isn't strong enough, there's really not much you can do to conc. HCl to "supe it up" enough for your purposes. --Jayron32 04:16, 1 November 2010 (UTC)

No Conc. HCl is good enough. Actually what I want to do is increase the potency of the H3O+ species in dilute HCl solution to that resembling the situation with conc. HCl, that is, to simulate the H+ species in conc. HCl. The concentration will be less, of course. I don't believe it's just a change in activity -- the change in free energy changes too, changing K. John Riemann Soong (talk) 04:23, 1 November 2010 (UTC)

(ec)You can't get the hammett acidity function significantly lower in water because if you try water acts as a base - if you add a very strong acid to water it just protonates the water (water acting as a base) and the acidity is limited by the actiity of the 'hydronium ion' - possibly you can get it a little more acidic using salts - but in general if you want to get greater protonating power you need to use a less basic solvent.

however If you really want to increase the acidity of an aqueous solution so that it has greater protonating power you need to destabilise the oxonium ion - this is best done with a salt of a strong lewis acid, and very weak lewis base - eg LiBF4 , (but please bear in mind that the pKa does not change, only the activities change) - this is how pKa is defined - thus you want to increase the activity of H+ since pKa is by mathematical definition fixed. (NaCl will not do it)Sf5xeplus (talk) 04:25, 1 November 2010 (UTC)

Won't a concentrated salt (say non-chloride) solution interfere a lot with water's stabilising basicity? It seems to me the effect is rather large as a significant stabilisation of the oxonium ion seems to come from intermolecular delocalisation of charge. AFAIK pKa is based on free energy change per mol, which tends to be constant in most situations but not always. John Riemann Soong (talk) 04:34, 1 November 2010 (UTC)

It depends on the anion an cation of the salt - specifically their acidities and basicities.Sf5xeplus (talk) 04:49, 1 November 2010 (UTC)

Question - is it clear how adding LiBF4 to water will decrease its basicity ? Sf5xeplus (talk) 04:28, 1 November 2010 (UTC)

Well Li+ is a weak Lewis acid (in the sense of vinegar weak). The BF4- anion is a weaker base/anion than Cl-, but I don't see how Li+ destabilises H3O+ any more than additional H+ ions might. John Riemann Soong (talk) 04:44, 1 November 2010 (UTC)

You said you didn't want to add additional H+, but make do with a weak HCl solution .. :) Therefor I think you've got it.Sf5xeplus (talk) 04:47, 1 November 2010 (UTC)

I understand BF4- is a weaker anion, and Na+ isn't as strong of a Lewis acid, but why wouldn't NaCl (or sodium nitrate, sodium perchlorate, etc.) achieve somewhat of the same effect? Wouldn't Na+ destabilise oxonium cations too? With conc. HCl, is the oxonium ion at its maximum destabilisation? Any further destabilisation would simply result in reprotonation of Cl-? John Riemann Soong (talk) 05:07, 1 November 2010 (UTC)

LiBF4 is overall much more of a lewis acid that NaCl right?Sf5xeplus (talk) 05:31, 1 November 2010 (UTC)

Distinguishing H3O+ activity from H3O+ acid strength

I am not really convinced that these are the same. Activity is effective concentration, perhaps in this context effective pH. There are things that dilute H3O+ would not protonate but concentrated HCl would....to me, it's not simply because there's more H3O+ to go around such that the protonated reagent is favoured in the equilibrium. A kinetic barrier appears to be crossed that seems to be related to the free energy of the H3O+ species? If we take protonating a carbonyl in the Fischer esterification, there is a large surge in reaction rate that does not appear to be solely due to the increased activity of H+. (We could take methanol and formic acid say, and watch how much methyl formate we get at a reaction run at 34C.)

AFAIK an equilibrium constant is the ratio of all the relevant forward rate constants over the reverse rate constants. Isn't it possible that due to changes in solvation structure or the type of H3O+ species in solution that you get H3O+ species with different "eases" of protonation, i.e. different rate constants for protonating "tough to protonate" reagents and getting reprotonated? Activity AFAIK is effective concentration, perhaps affected by the other solutes "stealing solvent", but would not explain a different hydrogen bonding / solvation structures adopted by H3O+ species. That is, the H3O+ species in conc. HCl is different from the H3O+ in dilute HCl solution, and isn't merely the same species at a higher activity. John Riemann Soong (talk) 06:02, 1 November 2010 (UTC)

For all that you need Activity coefficient which is the ratio of effective concentration to 'mass' concentration (ie activity:concentration)

Changes in the Activity coefficient with concentration (ie a non linear relationship between activity and concentration) are usually taken to mean that the average chemical enviroment has changed.Sf5xeplus (talk) 06:13, 1 November 2010 (UTC)

There's a table of activity coefficients of HCl(aq) here google books Corrosion of metals: physicochemical principles and current problems By Helmut Kaesche Sf5xeplus (talk) 06:30, 1 November 2010 (UTC)

According to google books:Outlines of Theoretical Chemistry By Frederick H. Getman p.452 the high measured activity coefficient of conc. HCl might be due to HCl molecules.Sf5xeplus (talk) 06:32, 1 November 2010 (UTC)

Hmm. Would there be an observable difference between hydronium pKa dropping 1 pKa unit versus the activity coefficient increasing tenfold? (Thanks so much for your effort in finding sources btw. Wikipedia chemistry articles usually do not have any concentration-dependent data...perhaps we should update them.) John Riemann Soong (talk) 07:38, 1 November 2010 (UTC)

I only have ever taken two classes that mentioned real solutions, and they always tried to avoid the concentrated solution case. Why does the activity coefficient fall into a minimum for both univalent salts and HCl? This is not the same as the free energy minimum is it? Does the activity coefficient minimum occur at the same concentration as the azeotropic concentration? John Riemann Soong (talk) 07:49, 1 November 2010 (UTC)

1. possibly in some cases (more likely for low values of Ka)- because activity is not necessarily linear with the change in equilibrium constant , whereas it is linear with change in activity coefficient (for constant concentration)

2. read activity coefficient - the coefficient usually drops to 1 as concentration tends to 0 - because 1 represents the ideal solution case, and this is best got from very low concentrations - it's in the article.

2b and c No,no. Sf5xeplus (talk) 08:04, 1 November 2010 (UTC)

The activity coefficient minimum for HCl soln is near 20% ....that's very far from zero if "saturated" is at 38%. It doesn't seem like a coincidence to me that it's near the eutectic and azeotropic compositions. John Riemann Soong (talk) 08:11, 1 November 2010 (UTC)

Oops, my bad, I'm doing the conversions wrong. (Is ionic strength == N?) There is an activity coefficient minimum < 0.78 between an ionic strength of 0.1 and 1.0, whatever that means. John Riemann Soong (talk) 08:14, 1 November 2010 (UTC)

1.0M HCl ~ 3.2% w/w Sf5xeplus (talk) 08:26, 1 November 2010 (UTC)

Hmm upon examining the data at Hydrochloric acid, it seems to me that H3O+ Cl- is optimally solvated at around 20% or near the azeotropic concentration (interestingly this is near the eutectic point as well). Does the falling activity coefficient with increasing concentration (up to around 20% w/w HCl) this actually mean the average strength of the H3O+ species decreases with concentration until the azeotropic concentration is reached?

Is this when the derivative of free energy of solvation with respect to concentration is lowest? This seems highly peculiar to me -- I would assume that maximum decrease in the free energy of solvation would be near a concentration of zero. None of this was ever mentioned in my chemistry classes. And it is also so very relevant! Why did they omit it? John Riemann Soong (talk) 08:04, 1 November 2010 (UTC)

Is this wrong now? Can you slow down a bit, cos this is starting to read like random noise. Please double check the above because it looks nearly completely wrong.Sf5xeplus (talk) 08:19, 1 November 2010 (UTC)

The relationship between free energy and activity is given at activity coefficient ie

G=G₀+RTln(a_cc) where a_c is activity coefficient at concentration, and c is concentration

differentiating with respect to concentration gives:

dG/dc=RT/c + ( RT/a_c x d(a_c)/dc )

I don't think that will have a minima.Sf5xeplus (talk) 08:37, 1 November 2010 (UTC)

clarify the original question a long time ago - if another species starts to appear at high molality eg lets say H4O²⁺ then that has a separate pKa, (because it's a different acid) the pKa of H3O⁺ or whatever doesn't change.

1. If the pKa (or activities) look odd at high conc. then this could mean that a new compound is being former at high conc.

2. But in stuff like water it gets complex due to the dynamic nature of the liquid - it may be right to take a fuzzy approach since actually cations are hard to categorise or temporal - it's easier to just say that the average 'proton' energy increases, and map that onto the activity figure

3. And finally - this means that some experimental activity figures may not be theoretically correct - this is likely for the high activity coefficients in HCl that are found when HCl is very concentrated ie when they go to 4 or 8. High figures like this suggest option 1. 87.102.115.141 (talk) 09:47, 1 November 2010 (UTC)

I note from the physical properties of concentrated HCl that the viscosity increases drastically after some threshold. But nonlinear viscosity behaviour with concentration appears to be consistent with perhaps saturated NaCl solution [1]. (This is not to even include the interesting nonlinear behaviours for viscosity, activities and K_w to be gotten at high temperatures.) Is K_w necessarily ~10^-14 at 20C in concentrated salt solution -- that is, is neutral pH for a saturated salt solution at RT necessarily 7? John Riemann Soong (talk) 19:47, 1 November 2010 (UTC)

K_water is really for pure water at standard conditions, so changes could be factored in using the changes in activitys, or activity coefficients (which are related to the changes in the free energy of the components)

Really you should call it K_{ionisation of water in saturated salt}, and yes it will be a bit different.Sf5xeplus (talk) 05:08, 2 November 2010 (UTC)

As far as I can tell the physical properties of conc. HCl are fairly smooth with concentration. except vaporpressure - which is due to the HCl solution being saturated above 38% at 1 atm.Sf5xeplus (talk) 05:11, 2 November 2010 (UTC)

By the way - in terms of chemical kinetics, the effects of ionic strength on reaction rate are described by the primary salt effect and secondary salt effect see google books Chemical kinetics and reaction dynamics Santosh K. Upadhyay pp.190-2 and 192-3 (or search - there are many examples on google).

Note that the primary salt effect depends on all the reaction materials not just the H+ : so is complex

These aren't really applicable when the reaction species changes. The secondary salt effect does cause increasing ionisation (of weak acids) as salt concentration increases .. due to increasing electrolytic strenght (dielectric strength of the solvent also increases)

For the same reason increasing salt strength could stabilise different H+(aq) ions - but the key term here is stabilise - for that reason I wouldn't expect greater protonating power to result. Sf5xeplus (talk) 07:22, 2 November 2010 (UTC)

But it occurs to me that whatever H+ protonates will be stabilised as well! ... thus, in concentrated salt solution would I see a greater favourability for SN2 reactions, or for anomeric reactions, or for acyl reactions, especially under neutral conditions? If I put methanol and acetone in concentrated salt solution for example would there be a lower barrier towards forming methyl formate (and vice versa -- a lower barrier to hydrolysis under neutral conditions)? The book says that primary salt effect doesn't appear to affect neutral reagents too much, but they seem to be talking about inorganic chemistry -- what about steps in organic chemistry that require separation of charge or where permanent dipoles already exist? So if I wanted to attack an alkyl chloride for example, in addition to the whole trick of adding potassium iodide, would adding more salt in general also help? John Riemann Soong (talk) 09:16, 2 November 2010 (UTC)

Adding salt increases the dielectric constant - assuming that is the only effect then anything with charge, or a dipole should be stabilised - both reactants, products and intermediates. How that affects the rate of a particular reaction with depend on what exists at the rate limiting step primarily - eg for anionic nucleophilic attack on an (neutral) alkyl halide the intermediate and the anionic reactant will both be stabilised - this makes the nucleophile less nucleophilic, but the intermediate will be a little more stable.

The best examples are probably when a neutral substance fissions into charged species eg SN1 nucleophilic substitution, or ionisation of a weak (non charged) acid eg acetic acid - both of these are helped by increasing dielectric constant.Sf5xeplus (talk) 09:44, 2 November 2010 (UTC)

Something like the reaction of an alcohol ROH with an acyl chloride RCOCl would also probably be speeded up, since the intermediate is a zwitterion.Sf5xeplus (talk) 09:53, 2 November 2010 (UTC)

I meant the value of the Dielectric constant when I was saying dielectric strength, not to be confused with this type of Dielectric strength.Sf5xeplus (talk) 09:59, 2 November 2010 (UTC)

Birefringence, redux

I suppose I should make my question clearer. I'm trying to write a lab report, and the article isn't at all helping identify which directions birefringence can be seen. In a uniaxial crystal, there is only two orientations (out of 6) in which birefringence can be seen, yes? (I had to guess from the hexagonal crystal structure of calcite -- the language isn't clear at all!) I couldn't figure out from the article

....if I laid the crystal shown in the images on its side and looked at the words, neither of the 4 faces would show birefringence, yes? John Riemann Soong (talk) 05:29, 1 November 2010 (UTC)

this image?

You mean this image? with 4 opposed faces. As far as I can tell there are 3 faces of one type and 1 (hexagonal) face of another.

If you pick one of the three faces, and look 'down' into it (offset from the perpendicular) you should see that as you rotate around the perpendicular the trigonal groups (CO₃^2- carbonate) go from looking flat, to side on, to flat etc - this seems to be anisotropy and I would therefor expect double refraction based on what the Birefringence article says.

The hexagonal face doesn't have such a strong effect (though there is some minor change every 120degrees) - I wouldn't expect birefringence as much. (or at all)

This isn't my subject - if that doesn't sound right I'd wait for an expert opinion.Sf5xeplus (talk) 09:22, 1 November 2010 (UTC)

KNO3 + KI

What happens when potassium nitrate and potassium iodide are heated? Is iodine produced? --Chemicalinterest (talk) 12:45, 1 November 2010 (UTC)

According to [2], the iodide is oxidized to the iodate. shoy (reactions) 13:16, 1 November 2010 (UTC)

Solar cells: Difference between dark current and diode current

Hello, I have noticed that some texts refer to diode current and dark current in PV cells interchangeably. Could someone kindly explain the difference, if there is one?

The most common rendition I have seen is:

Diode current = dark current . (exp(eV/kT) - 1)

However I have also seen basically the same equation to define dark current:

Dark current = Jo . (exp(eV/kT) - 1) where Jo is an unidentified constant.

Assuming they are *not* the same thing, what is the physical nature of how they differ? To my understanding both terms describe a forward current that flows the p-n diode under the influence of an externally applied voltage.

Many thanks. —Preceding unsigned comment added by 146.23.212.21 (talk) 13:51, 1 November 2010 (UTC)

You are looking at two different cases of "loose terminology" describing the ideal diode equation. A photoelectric cell is a photodiode, and ideally behaves according to this equation that relates current and voltage. The term "dark current" is being used to refer to the current that would flow even if no light is shining on the cell. In actuality, that is a bit loose - what this really refers to is the saturation current and is a semiconductor physics parameter. In reality, diodes are very complicated: see our article on diode modelling for an introduction to some of the mathematical idealizations of them. In any case - the parameter J₀ probably refers to the effective current density for a particular device, and depends on diffusion parameters of the semiconductor - i.e., its geometry, its chemical makeup (doping), and so on. Here is an extremely technical overview of the semiconductor physics that are relevant to this process in a P-N junction: Basic Device Behavior. (While this overview is horribly complicated, the underlying processes in a PN junction actually are complicated - so there's no sense sugar-coating it). You can accept the derivations at face-value, or you can work with idealized, empirical models of current/voltage relationships and light-intensity/power-generation measurements. Nimur (talk) 17:26, 1 November 2010 (UTC)

Thanks Nimur, glad that the problem is the complexity of the subject and not my stupidity :) (this time). Yes for my purposes it is sufficient to understand the dark current as being that which flows in the diffusion direction in a darkened cell, and balances out the drift current. As a (forward) external voltage is applied, the electric field in the space-charge region is reduced, allowing diffusion to increase, yet the drift current remains unaffected since there's still just as many electron-hole pairs being created, hence the overall diode current increases as described by the first equation given above. —Preceding unsigned comment added by 146.23.212.21 (talk) 08:21, 2 November 2010 (UTC)

knife

what kinda knife do the usa army use today? —Preceding unsigned comment added by Kj650 (talk • contribs) 22:01, 1 November 2010 (UTC)

The United States Army issues the M9 bayonet. Some services and special operation groups may use other equipment, including the Aircrew Survival Egress Knife, the United States Marine Corps' standard OKC-3S bayonet, or older M7 or M6. Here is the official page for the Bayonet at the US Army website. See also, Bayonets, knives, bayonet-knife models for a more thorough list, including historical issues. KA-BAR, while popular amongst service members for historic reasons, is not standard issue, although in some cases it may be officially issued to active duty Army units. Nimur (talk) 22:32, 1 November 2010 (UTC)

Reading M9 bayonet, I see that "some production runs of the M9 have a fuller and some do not". Blood groove redirects to Fuller (weapon), but that article currently makes no mention of that common term despite it being the title of two of its four references. Should Fuller (weapon) contain some explanation of "blood groove". -- 124.157.218.101 (talk) 01:23, 2 November 2010 (UTC)

I don't think so. Our Fuller article explains the function quite well. If anything, the referring pages should be changed to fuller. "Blood groove" is based on a misconception-- that the groove somehow increases blade performance by letting blood escape through the groove. This is false. If anyone wants to include this terminology, it should be clear that it is a misnomer.SemanticMantis (talk) 16:36, 2 November 2010 (UTC)

recycling code

what plastic recycling code do vacuum cleaners have —Preceding unsigned comment added by Kj650 (talk • contribs) 22:37, 1 November 2010 (UTC)

You asked this already, but I guess you got no answer. Ariel. (talk) 00:01, 2 November 2010 (UTC)

Usually, recyclable plastics will have a code blazoned on to them. If the vacuum cleaner doesn't specify, it's useless for us to speculate. Various brands and models use different manufacturing processes. Do you have a specific model in mind? Nimur (talk) 00:15, 2 November 2010 (UTC)

(Edit Conflict) I doubt that there can be a simple answer to this since: (a) different makes and models will surely use different plastics; (b) the same model will almost certainly use different plastics for different components; and (c) any model will certainly include significant non-plastic components (for example, metals in the motor). If particular plastic components are recyclable, they may carry an applicable symbol (probably on their inside surface). Recycling a vacuum cleaner in order to salvage and reuse component materials could probably only be done by a professional outfit specialising in such work, since some components will likely contain substances requiring special handling or disposal - some junked electrically powered equipment is exported to third-world countries where such work is profitable.

You might find the list of articles at Index of recycling topics useful. 87.81.230.195 (talk) 00:24, 2 November 2010 (UTC)