Anders Hoveland

"investigation of sodamide and of its reaction-products with phosphorus" William Phillips Winter p42-43 The article goes on to say that the reaction of sodium phosphide with water generates small portions of phosphites and hydrogen gas, in addition to the main products of sodium hydroxide and phosphine. Hydrogen sulfide is known to be able to reduce sulfur dioxide (actually sulfurous acid in aqueous solution) at room temperature. Perhaps phophine would similarly be able to reduce phosphorous acid (H3PO3)? It seems apparent that hypophosphorous acid, H3PO2, is reduced by phosphine, PH3, to elemental phosphorous. The selective oxidation of PH3 by aqueous iodine also produces H3PO2. (in a similar reaction, iodine reduces thiosulfate, S2O3[-2], to tetrathionate, S4O6[-2] ) The fact that H3PO2 even exists suggests that the chemistry of the P-O bond is more similar to the C-O bond than the S-O bond, which is to say that H3PO3 is essentially not an oxidizer at ordinary temperatures, like SO2 can be. However, one paper mentions that PH3 can reduce 1-naphthol to naphthalene: "Phosphine as a Reducing Agent" SHELDON A. BUCKLER, LOIS DOLL, FRANK K. LIND, MARTIN EPSTEIN. J. Org. Chem., 1962, 27 (3), pp 794–798 So the idea that PH3 could reduce H3PO3 is quite plausible. Whether PH3 reacts with H3PO3 in aqeous solution is uncertain. (the paper also describes the reaction of phosphine with nitrobenzene. There was no reaction at neutral conditions, but when sodium hydroxide was added then azoxybenzene was produced in high yield) It should be obvious that PH3 would not reduce H3PO4, just as H2S does not reduce aqueous H2SO4. "A dictionary of chemistry and the allied branches of other sciences, Volume 4", Henry Watts, p524 Various Other Hydrides of Phosphorous P2H4, first obtained by P. Thenard (Comptes rendus, 1844, 18, p. 652) by decomposing calcium phosphide with warm water, the products of reaction being then passed through a U tube surrounded by a freezing mixture (see also L. Gattermann, Ber., 1890, 23, p. 1174). It is a colourless liquid which boils at 57°-58° C. It is insoluble in water, but soluble in alcohol and ether. It is very unstable, being readily decomposed by heat or light. By passing the products of the decomposition of calcium phosphide with water over granular calcium chloride, the P2H4 gives a new hydride, P12H6 and phosphine, the former being an odourless, canary-yellow, amorphous powder. When heated in a vacuum it evolves phosphine, and leaves an orange-red residue of a second new hydride, P9H2 (A. Stock, W. Bottcher, and W. Lenger, Ber., 9 9, 4 39, 47, 2853). P4H2, first obtained by Le Verrier, is formed by the action of phosphorus trichloride on gaseous phosphine (Besson, Comptes rendus, 111, p. 972); by the action of water on phosphorus di-iodide and by the decomposition of calcium phosphide with hot concentrated hydrochloric acid. It is a yellow solid, which is insoluble in water. It burns when heated to about 200° C. When warmed with alcoholic potash it yields gaseous phosphine, hydrogen and a hypophosphite. Interesting Non-Reactivity of Phosphorus: Elemental phosphorous can apparently be dissolved in liquid sulfur dichloride without reaction.

This is exactly why I do not believe in free trade. Or at least not with other countries that have much lower living standards and wages. It becomes difficult to ensure good wages here when the jobs can be so easily outsourced to some third world country. If Britain does not have enough good jobs, perhaps it is not a good time to be bringing in more workers.

Governments need to be more careful before they let in more immigrants, especially from certain countries where the people are likely to comitt crime at higher rates (Pakistan, North Africa, Mexico). And they should stop trying to push ethnic diversity in the schools. Different ethnicities generally prefer to live apart. Why have people not been having enough children? Perhaps because they do not feel they can afford them? The cost of housing is just ridiculous. Many young families must delay having children if they want to give their children a decent life out of poverty. Simply bringing in more immigrants, who are willing to have many children while living in relative poverty, is just going to make the underlying problem worse. Britain is overcrowded, there is a shortage of affordable housing. More immigrants will just mean even fewer births of native of Britains. Increased population, especially desperate immigrants, will just drive down wages. Britain would be much better off with a "shortage" of workers than it would be with too many. Indeed, Britain does have too many. There is no shortage of workers. If there was, wages would be increasing and the young adults would be able to afford to live on their own!

Mass immigration has been the undoing of leftwing political parties across Europe since it erodes the shared values that are an essential prerequisite of a well-funded welfare state. Why should indigenous, working populations support the high levels of taxation necessary to sustain generous welfare payments if the beneficiaries are people unlike themselves? If they can't look at a benefit recipient and think, "There, but for the grace of God, go I", why should they continue to pay such high taxes? This problem was spelt out by David Willetts a few years ago: The basis on which you can extract large sums of money in tax and pay it out in benefits is that most people think the recipients are people like themselves, facing difficulties that they themselves could face. If values become more diverse, if lifestyles become more differentiated, then it becomes more difficult to sustain the legitimacy of a universal risk-pooling welfare state. People ask: 'Why should I pay for them when they are doing things that I wouldn't do?' This is America versus Sweden. You can have a Swedish welfare state provided that you are a homogeneous society with intensely shared values. In the United States you have a very diverse, individualistic society where people feel fewer obligations to fellow citizens. Progressives want diversity, but they thereby undermine part of the moral consensus on which a large welfare state rests. When it comes to differences between countries, social cohesion plays a major role. Broadly speaking, countries that are more ethnically or racially homogenous are more comfortable with the state seeking to mitigate inequality by transfering some of the resources fr9om the richer to poorer people frough the fiscal system. This may explain why Swedes complain less about high taxes than the inhabitants of a country of immigrants such as America. But it also suggests that even in societies with a tradition of high taxes (such as those in Scandinavia) might find that their citizens would become less willing to finance generous welfare programs were immigrants to make up a greater share of their populations.Robert Putnam studied the downsides of diversity, and found that people in more diverse communities tend to "distrust their neighbors, regardless of the color of their skin, to withdraw even from close friends, to expect the worst from their community and its leaders, to volunteer less, give less to charity and work on community projects less often, to register to vote less, to agitate for social reform more but have less faith that they can actually make a difference, and to huddle unhappily in front of the television." Michael Jonas wrote about ethnic diversity, "Birds of different feathers may sometimes flock together, but they are also less likely to look out for one another."

Acetic Anhydride the preparation of acetic anhydride, where Ac is an acetyl grop, and AcOH is acetic acid. 4AcOH + 2SCl2 --> 2AcCl + 2Ac2O + S + SO2 + 4HCl It should be obvious that this reaction cannot be conducted with more than a very small ammount of water present,and that most of the SO2 and HCl will be evolved out as a gas. Alternatively, dry sodium acetate can be used. Sodium acetate can be dried by heating to the pure compound's melting point, while stirring vigorously as the water boils out. Another reaction uses the volatile and highly poisonous PCl3 reacting with acetic acid to form acetyl chloride CH3C(=O)Cl,and phosphorous acid H3PO3. The acetyl chloride then reacts with anhydrous sodium acetate and acetic anhydride is distilled out. Acetic anhydride can also be prepared by fast pyrolysis (at 700 degC) of acetone, producing ketene CH2=C=O, which is then reacted with glacial acetic acid to form the final acetic anhydride product. sulfur dioxide forms an addition product with sodium acetate, which reacts with chlorine to produce either acetyl chloride or acetic anhydride, depending on the proportion of sodium acetate used. (German Patent 210805, year 1910) Acetic anhydride which contains a small quantity of acetic acid boils at around 138 degC. Cold acetic anhyride can remain unhydrolyzed in contact with water for a considerable period of time, but at 20 degC the hydrolysis proceeds rapidly.Acetic anhydride reacts with nitric acid at low temperature to form a compound with the formula C4H9NO7. This compound is probably (AcO)2NO2H·H2O Acetic anhydride reacts with N2O5 at low temperatures to form acetyl nitrate. (Pictet and Khotinsky, year 1907)

I think this thread's title should be changed to "Yes, you CAN make Sodium!", since there exist several different routes to preparing the element sodium, without electric current, two of which have been the subject of much successful experimentation by amateur home chemists.

I had an idea for a reaction to easily produce nitryl fluoride. The reaction has not been tested. (2)CaFeF6 + (22)HNO3 --> (4)Ca[NO3]2 + (4)Fe[NO3]3 + (22)HF + (2)NO2F + O2 Ideally the concentration of the nitric acid used in the reaction should be greater than 98%. Preparing the hexafluoroferrate H2O2 + 6HF + FeO4(-2) --> FeF6(-2) + 4H2O + O2 The H2O2 and HF must be premixed, then added to ferrate solution. A similar reaction is known to occurr for permanganate. After completion of reaction, a solution of calcium nitrate is added, which would cause CaFeF6 to precipitate out. The solution is filtered, and the precipitate is dried. The solution needs to stay alkaline before the CaFeF6 is precipitated out as a solid The fluoroferrate ion FeF6(-2) is probably only stable in alkaline solution. Addition of acid likely would cause hydrolysis, with the loss of oxygen: 4 FeF6(-2) + 2 H2O + 8 H+(aq) --> 4 FeF3 + 12 HF + O2

Will science ever come to and end? or will it continue infinately? By definition, we cannot know about what we do not know about! Science might eventually dwindle as all the "low hanging apples" currently accessable to humans run out. This is not to say that humans will have discovered all natural phenomena, simply that it may be increasingly difficult for them to make additional discoveries. Alternatively, there may be unexpected new discoveries, or ways of understanding natural phenomena, that may greatly expand the abilitity of humans to discover new things. By 1900, most scientists believed that almost all scientific principles had already been discovered and adequately described. But this was certainly not the case. Even in the realm of chemistry, countless important reactions were later described, new fields of research emerged. I can only offer my own opinion, which is that completely new and unexpected fields of research will develop that will expand the investigative power of researchers even further. There might be a plateau period of relative inactivity before that period. My guess is that no significant discoveries will be made in the next few decades, but at some point there will be a very important discovery, or a collection of advancements that will interplay with eachother to lead to a rapid advancement of technology. I would also like to note that virtually all the important scientific research in the last century has emerged from government universities. Private industry has only refined and made commercially practical such research. But the very structural importance of government research may also be the main hindrance. Several decades ago, Feynman remarked that his colleagues pursued cargo cult science: an activity which is indistinguishable from science except for its lack of useful output. Scientific research decisions have, in some ways, become overcentralised. Curiosity-driven research is now frowned upon. But at the same time, more centralisation will be required to make future discoveries that will require a high level of resources and financing. Again, in my opinion, the "information technology" age has really not been very revolutionary compared to many of the industrial advancements before (electricity and the automotive engine). The scientific world has been relatively stagnant since 1970. The progress of Science also stalled in the historical period between 500 AD to 1000 AD. The progress of science is not always a continuous steady stream of advances. Science grows in sudden clusters of small discoveries, or in great leaps and bounds of understanding.

http://woelen.homesc...uIII/index.html Copper(II) salts can be oxidized to Copper(III) complexes, with the help of a hypochlorite or persulfate oxidizer, sodium hydroxide, and sodium periodate. The complex [Cu(H3IO6)2]-2 is oxidized to [Cu(H3IO6)2(OH)2]-3, which has a very deep brown-red color The composition of these complexes can more simply be thought of as: IO4[-],Cu[+2], IO4[-],(2)OH[-],(2)H2O and IO4[-],Cu[+3],IO4[-],(4)OH[-],(2)H2O Actually, Cu(OH)2 is soluble in concentrated solutions of NaOH showing that copper(II) oxide is actually slightly amphoteric. CsCuCl3, although copper is still in its normal +2 oxidation state, is a bright reddish-orange colored compound http://woelen.homesc...uCl3/index.html

All of the following compounds exist: K2FeO4 (brownish-purple color) iron in a +6 oxidation state Na4FeO6 (green color, unstable, decomposes in sunlight) iron in a +8 oxidation state RuO4 (ruthenium is directly below iron on the periodic table, ruthenium tetroxide is explosive) Cs2CuF6 copper in a +4 oxidation state K3CuF6 copper in a +3 oxidation state colored complexes containing copper in a +3 oxidation state can be obtained by reacting copper salts with sodium periodate and sodium hypochlorite. Complexes of Copper in Unstable Oxidation States, T. V. Popova and N. V. Aksenova, Maryi State University, Ioshkar Ola, Russia AgF2 silver in a +2 oxidation state AgF3 (cannot be prepared by direct combination of the elements) silver in a +3 oxidation state Au2F10 (exists as a dimer) gold in a +5 oxidation state NiO2 (commonly referred to as nickel "peroxide") nickel in a +4 oxidation state XeO4 xenon in a +8 oxidation state ClF5 chlorine in a +5 oxidation state ClF6[+] ion (but ClF7 does not exist, since reaction of Cl6+ with F- only makes Cl5 and F2) chlorine in a +7 oxidation state KO2 potassium superoxide KO3 potassium ozonide CsF2 (this is a little researched compound, interestingly it does not seem to be very oxidizing) cesium in the +2 oxidation state, which is the only +2 oxidation state reported from within the the first column of alkali elements CsAu gold in a -1 oxidation state N5[+] pentazenium ion

The mix of concentrated nitric and hydrochloric acids in aqua regia exists in an equilibrium. The equilibrium only begins moving to the right when the acids are very concentrated. Even heating 30% conc HCl with NaNO3 will give off some brown fumes. HNO3 + (3)HCl <==> (2)H2O + NOCl + Cl2 (2)NOCl + H2O --> NO + NO2 + (2)HCl[aq] Aqueous solution of chlorine can attack gold. (2)Au + (3)Cl2 + (2)Cl[-] --> (2)AuCl4[-] An interesting way to reduce gold(III) chloride back to elemental gold is to use an alkaline solution of hydrogen peroxide. Although usually oxidizer, in some reactions H2O2 can act as a reducing agent. (2)AuCl4[-] + (3)H2O2 + (6)OH[-] --> (2)Au + (8)Cl[-] + (6)H2O + (3)O2 The gold usually separates out in a finely divided state, and appears brown by reflected light and greenish blue by transmitted light. If very dilte solutions are used, the gold sometimes separates out forming a yellowish film on the sides of the test tube.

Dichlorine heptoxide, Cl2O7, is dangerous is several unusual ways. 95-100% Concentrated Perchloric Acid Physical Properties Pure perchloric acid is unstable. If left standing at room temperature for between 10 and 30 days, at some point the temperature of the acid will spontaneously rise to 90 degC, at which point there will be an explosion. Pure perchloric acid cannot be distilled at ordinary pressures. To avoid partial decomposition (into Cl2, O2, and various chlorine oxides), distillation must be done below 200 mmHg. Reactivity "I would expect CH2Cl2 to react with HClO4 (if not sooner, then later), since HClO4 seems to react with both CHCl3 and CCl4, and CH2Cl2 is less oxidized. I would be especially careful with CH2Cl2, it should easily be at least a detonable mixture. Namely, HClO4 is entierly miscible with CHCl3, the solution discolors after a few days to yellow, and in air sheds crystals of HClO4.H2O. Commercial chloroform contains alcohol, which sheds a heavy, with CHCl3-insoluble extraordinarily explosive oil (Vorländer, v. Schilling, Lieb. Ann. 310 [1900] 374; Vorländer, Kaascht (Ber. 56 [1923] 1162). For CCl4, HClO4 is insoluble in CCl4, and gives upon shaking, a green emulsion, which discolors brown after several minutes welling up under formation of HCl and COCl2 (Vorländer, v. Schilling, Lieb. Ann. 310 [1900] 374). Preparation of solutions of Cl2O7 in CCl4 described in: F. Meyer, Keszler (Ber. 54, [1921] 569). the interaction of HClO4 with benzene has been described as follows: If one to two drops of HClO4 has 2 to 3 cm3 benzene poured over it, under heat evolution brown flakes precipitate. Mixing of equal volumes at first forms a green emulsion, which then explodes (Vorländer, v. Schilling, Lieb. Ann. 310 [1900] 374). 1 g of HClO4 solubilizes in 5g of well-cooled benzene under formation of a green solution, which when stored in a sealed tube, discolors increasingly shedding a carbon-containing material; after completion of the reaction, no free acid is found anymore (Michael, Conn, Am. chem. J. 23 [1900] 444). So mixing something like CCl4 and HClO4 can cost one their lives if not wearing protective gear, doing under fume hood,etc. It is dangerous to extrapolate so assuredly." Dichlorine Heptaoxide (perchloric anhydride) Cl2O7 Physical Properties Cl2O7 melts at (minus) -91.5 degC and boils at 80 degC. At 0 degC, it has a vapor pressure of 23.7 degC. Cl2O7 has a density of 1.86 g/mL. Cl2O7 decomposes spontaneously on standing for a few days. Cl2O7 decomposes into chlorine and oxygen under very low pressures (below 80mm Hg) or in the temperature range of 100-120 degC. Cl2O7 is soluble in benzene, slowly attacking the solvent with water to form perchloric acid. Cl2O7 detonates when heated or when subject to mechanical shock. It explodes on contact with flame or by percussion. Chlorine heptoxide is more stable than either chlorine monoxide or chlorine dioxide; however, the anhydride detonates when heated or subjected to shock. Reactivity Cl2O7 is soluble in benzene, slowly attacking the solvent with water to form perchloric acid; it also reacts with iodine to form iodine pentoxide.” Benzene reacts with Cl2O7 However this seems to contradict something else I read. "Cl2O7 is a strong oxidizer as well as an explosive that can be set off with flame or mechanical shock, or by contact with iodine. Nevertheless, it is less strongly oxidising than the other chlorine oxides, and does not attack sulfur, phosphorus, or paper when cold." Holleman, Arnold F.; Wiberg, Egon (2001). Inorganic chemistry. Translated by Mary Eagleson, William Brewer. San Diego: Academic Press. p. 464. ISBN 0123526515. ^ Byrns, A. C.; Rollefson, G. K. (1934). Journal of the American Chemical Society 56: 1250–1251. So Cl2O7 reacts with benzene, but not paper or phosphorous when cold? One explanation might be that a protective layer is formed that prevents further oxidation. P2O5 is known to form an adhesive layer on reaction with water that limits further hydration, for instance. However, failure of Cl2O7 to react with sulfur cannot be explained by a protecting layer, since nothing could form that would not be soluble or a gas. HNO3 attacks sulfur, but not benzene, so I have difficulty believing that Cl2O7 would react in an opposite way. Thus this is all quite confusing, and to make any sense of the conflicting references is a logic puzzle. Cl2O7 does not explode on contact with wood, paper or similar materials but just evaporates (Michael, Conn, ibid). They also note unreaction towards sulfur and phosphorus pieces. HClO4 on the other hand does explode violently on contact with wood and paper, and especially charcoal (Roscoe, Lieb. Ann. 121 [1862] 353). Cold, dry C6H6 solubilizes Cl2O7, then soon afterwards a reaction occurs (A. Michael, Conn, Am. chem. J. 23 [1900] 446). Cl2O7 reacts with iodine to form iodine pentoxide (I2O5). Addition of iodine to Cl2O7 can result in explosion, however, so the reaction is probably best done diluted in an appropriate solvent. Other compounds related to Cl2O7 Perchloryl fluoride (FClO3) is very stable, poisonous and reactive (Bp. -46.7C, Mp. -147.7C). Electrolysis of satd. NaClO4 in anhydrous HF yields the compound. Another way in 85-90% yield, is to warm a mixture of KClO4, HF and SbF5 at 40-50 C (Kirk Othmer). FClO3 is also stable up to 400 C, and hydrolyzes slowly. Grease and rubber tubing has caused explosions, for more reactivity see Brethericks. The German wikipedia claims alkali fluorides reacting with Cl2O7 does yield FClO3, though no exact reference is given for this. Hantzsch claimed to have made fluoronium perchlorate [FH2]ClO4 by reacting anhydrous HClO4 with anhydrous liquid HF, which under strong heat evolution was said to yield solid [FH2]ClO4 (Ber. 60 [1927] 1946), and which compound he said reacts explosively with H2O (Ber. 63 [1930] 97). Brauer and Distler (Z. anorg. u. allgem. Chem. 275 [1954], 157) tried to make this compound, but could not repeat preparation despite mixing in various ratios and temperatures. FClO3 is also made by reacting fluorine with KClO3 at -20 C in SbF5: KClO3 + F2 = KF + FClO3 Or by reacting KClO4 with HSO3F: KClO4 + HSO3F = FClO3 + KHSO4 (From: Lehrbuch der anorganischen Chemie by A.F. Holleman, E.Wiberg, N.Wiberg). On the last one, no more decent details given by Holleman et al., it could react right away or may need some warming, time to hit the more serious lit. With ammonia, FClO3 forms a perchlorylamide: FClO3 + NH3 = ClO3(NH2) + HF. This has acidic protons and they are replaceable by metal ions: K[ClO3(NH)] and K2[ClO3N] these are colorless, up to 300 C stable compounds, which explode by impact. "Dichlorine hexoxide is a dark red fuming liquid at room temperature that crystallizes as an ionic compound, chloryl perchlorate, [ClO2]+[ClO4]−. Many other reactions involving Cl2O6 reflect its ionic structure, [ClO2]+[ClO4]−, including the following: NO2F + Cl2O6 → NO2ClO4 + ClO2F "

Only useful education helps create more wealth. But it seems that most forms of education is modern society are not useful, only competitive demonstrations of competency. In this sense then, education does not help create more wealth, but rather, in a way, is a waste of time and resources. While on the scale of an individual, education leads to a higher personal income, this is not necessarily the case in society. Education often only leads to higher income because the employer would rather give the good paying jobs to those that have more education. So if other people obtain higher educational credentials, it will just mean lower incomes for the people that do not have such credentials. I believe this misunderstanding is referred to as the fallacy of composition. And the phenomena of economic disincentives caused by non-practical overeducation in a society has been given the name credentialism. A society only needs a small portion of individuals to be educated in science, medicine, or advanced mathematics. Trying to teach a greater number of people these specialised areas of knowledge will not benefit society.

That same reaction is unlikely to work with gold. AuF3 quickly hydrolyses with water to form Au2O3 and HF, which is not soluble unless the solution is very acidic. Heavy metals tend to form stronger bonds with other elements that have similarly larger atomic orbitals, such as sulfur or chlorine. The bonding has both ionic and complex coordination character. Anhydrous HNO3 by itself can actually slowly attack gold, especially if a small quantity of N2O5 (the anhydride of nitric acid) is dissolved. The nitric acid must be extremely concentrated, because the equilibrium for reacting with gold is so unfavorable. These extremely high concentrations of nitric acid, however, are generally not commercially available and rarely ever used. The main difficulty of dissolving gold is actually not oxidizing it, but rather complexing it to make it soluble. Gold is an unusual element in that gold-oxygen bonds are only metastable. Neither is it easy to make gold atoms ionize, both because of the elements high electronegetivity, and because the metal ions would not complex with water like virtually all other metal ions do. The highly corrosive and unstable red-coloured liquid, dichlorine hexoxide, can also dissolve gold, through the formation of a complex. (4)Cl2O6 + Au --> [ClO2]+[Au(ClO4)4]− + (3)ClO2 The reaction of ClO2 with ozone produces the red droplets Cl2O6, which is a hazardous explosive.

No doubt the existence of amide ions is EXTREMELY unfavorable in ammonium hydroxide. The equilibrium would be very very small, virtually negligable for most reaction kinetics. But the room temperature, uncatalysed reaction between ammonium hydroxide and hydrogen peroxide is also very slow. There is no observable reaction products after 24 hours. It takes several days before enough ammonium nitrite is created to be detectable. The very slow reaction rate is not incongruent with an amide reaction mechanism explanation. The reaction between pressurised liquid anhydrous NH3 and pure H2O2 would likely proceed at a much faster rate. As for amide getting protonated in water, I do not know the exact equilibrium constant, but I do know that sodium amide rapidly and vigorously reacts with water, so the equilibrium of amide existence in water must be very low. NH2[-] + H2O --> NH3 + OH[-]

One argument that is often made against helping the poor is that it is wrong to tax the wealthy, and that the poor have no right to any wealth. In this post, I would like to suggest that the poor may be entittled to compensation from the land and natural resources which are being withheld from them. "The deadliest form of violence is poverty" - Ghandi "The earth is the general and equal possession of all humanity and therefore cannot be the property of individuals." - Leo Tolstoy "When the 'sacredness of property' is talked of, it should always be remembered, that any such sacredness dos not belong in the same degree to landed property. No man made the land. It is the orginal inheritance of the whole species.. It is no hardship to any one to be excluded from what others have produced ... But it is some hardship to be born into the world and to find all nature's gifts previously engrossed, and no place left for the new-comer ... To me it seems almost an axiom that property in land should be interpreted strictly, and that the balance in all cases of doubt should incline against the properitor." - John Stuart Mill "Every man has a property in his own person. This nobody has a right to, but himself." - John Locke "Whenever there is in any country, uncultivated lands and unemployed poor, it is clear that the laws of property have been so far extended as to violate natural right." - Thomas Jefferson "Just as man can't exist without his body, so no rights can exist without the right to translate one's rights into reality, to think, to work and keep the results, which means: the right of property." - Ayn Rand "So long as the great majority of men are not deprived of either property or honor, they are satisfied." - Niccolo Machiavelli "Don't you know that if people could bottle the air they would? Don't you know that there would be an American Air-bottling Association? And don't you know that they would allow thousands and millions to die for want of breath, if they could not pay for air? I am not blaming anybody. I am just telling how it is." - Robert Ingersoll "Thieves respect property. They merely wish the property to become their property that they may more perfectly respect it." - G.K. Chesterton "Our houses are such unwieldy property that we are often imprisoned rather than housed by them." - Henry David Thoreau "Whenever there is a conflict between human rights and property rights, human rights must prevail." - Abraham Lincoln "The first man who, having enclosed a piece of ground, bethought himself of saying This is mine, and found people simple enough to believe him, was the real founder of civil society. From how many crimes, wars and murders, from how many horrors and misfortunes might not anyone have saved mankind, by pulling up the stakes, or filling up the ditch, and crying to his fellows, "Beware of listening to this impostor; you are undone if you once forget that the fruits of the earth belong to us all, and the earth itself to nobody." - Jean Jacques Rousseau, A Discourse on the Origin of Inequality "In my opinion, the least bad tax is the property tax on the unimproved value of land, the Henry George argument." - Milton Friedman "As soon as the land of any country has all become private property, the landlords, like all other men, love to reap where they never sowed, and demand a rent even for its natural produce." - Adam Smith, Wealth of Nations "A tax upon ground-rents would not raise the rents of houses. It would fall altogether upon the owner of the ground-rent." - Adam Smith I am not trying to use this philosophy to justify anything extreme, just to suggest that there should be a land value tax on the unimproved value of land and natural resources, and that each person (at least within the country) deserves to be given a payment for being deprived of such land and natural resources. The tax could be used both to help fund the government, and to make the payments to citizens. This would be a fair and practical way to help reduce poverty. Many economists think that income taxes create disincentives in the economy, so one solution to this potential problem would be to shift away from income tax, and start to use a land/natural resource value tax instead. If a person's land value taxes were less than his entittled compensation payment, he would not be liable to pay any land value tax. Which is to say that those who own less than the average price value of land, less than their "fair share", would not be required to pay such a tax.

Mixtures of hydrobromic acid and bromine can dissolve gold. So too can a mixture hydroiodic acid, KI, and iodine. Gold slowly corrodes in aqueous solutions of chlorine. But HF almost certainly would not work to oxidize gold. Fluorine is a very electronegetive element, and so HF is nearly impossible to oxidize in most reactions. One rare reaction, however, is: Cl2O7 + HF --> HClO4 + ClO3F

What I am most interested in is what can oxidize NH3 to oxides of nitrogen at room temperature. Supposedly, hydrogen peroxide will react with ammonia in the presence of a catalyst, but there seems to be virtually no information about this in the literature. (whether the catalyst is iron salts or acetamide?) What exactly is the chemistry between H2O2 and NH4OH ? Does acetone serve as a necessary catalyst for the reaction? When the two are mixed, there is no observable reaction, and there are many references in the literature to mixes of the two chemicals, presumably without reaction. Under what conditions can ammonia be oxidized to NH4NO2 ? Excess H2O2 would presumably oxidize the NH4NO2 to NH4NO3. Nitrite is a much more reactive reducing agent than ammonia, so to obtain any nitrite one would think that a large excess of ammonia would have to be used. Hydogen peroxide slowly decomposes in aqueous alkaline solution, so one would expect NH4OH/H2O2 solutions to gradually decompose, either with the liberation of oxygen, or possibly the oxidation of ammonia. A mixture of hydrogen peroxide and ammonium hydroxide (in a 1:3 ratio) acts as a reactive oxidizer, which can attack organic compounds and elemental carbon. The reaction rate is negligible at room temperature, but when heated to 60°C the reaction becomes vigorous and self-sustaining. Such solutions are sometimes known as "base piranha". With a 1:1:5 volume ratio of NH4OH, H2O2, and H2O, respectively, the half-life times of peroxide were 4 hours at 50°C and 40 minutes at 80°C. "Reaction of Ozone and H2O2 in NH4OH Solutions and Their Reaction with Silicon Wafers" Japanese Journal Applied Physics. 43 (2004) pp. 3335-3339. Strong solutions of H2O2 with a few drops of NH4OH or solutions of ammonium carbonate (with or without NaOH or Na2CO3) can be let to stand 24 hours without any nitrite formation occurring. But upon longer standing, even with a small amount of hydroxide then nitrite forms. Nitrite also forms when a dilute solution of H2O2 is mixed with NH4OH and a little Na2CO3 and is evaporated over pure conc. H2SO4 with a bell jar. H2O2 forms (even in very dilute solutions) nitrite very rapidly, if the H2O2 solution is mixed with a few drops of NH4OH and a little NaOH or Na2CO3, and this then boiled in a retort to a very small volume. They suggest this nitrite formation as a demonstration experiment because it is very quick to do, and then after acidification of the colorless liquid with H2SO4, the HNO2 can be nicely be proven to be present. Hoppe-Seyler "A process developed by Produits Chimiques Ugine Kuhlmana (PcUK), and practiced by Atofina (France) and Mitsubish Gas (Japan) involves the oxidation of ammonia by hydogen peroxide in the presence of butanone (MEK) and another component that apparently serves as an oxygen-transfer agent. The reaction is carried out... at at 50degC. The ratio of H2O2/MEK:NH3 used is 1:2:4. Hydrogen peroxide is activated by acetamide and disodium hydrogen phosphate (117). The mechanism of this reaction involves an activation of the ammonia and hydrogen peroxide because these compounds do not themselves react (118-121). It appears that acetamide functions as an oxygen transfer agent, possibly as the iminoperacetic acid, HOOC(=NH)CH3, which then oxidizes the transient Schiff base formed between MEK and ammonia to give give the oxaziridine, with regeneration of acetamide." (117) U.S. Pat. 3,962,878 (aug. 3, 1976), J.P. Schirmann, J. Combroux, and S. Y. Delavarenne (118) J.P. Schirmann and S. Y. Delavarenne, Tetrahedtron :ett. 635 (1972) (119) E. G. E. Hawkins, J. Chem. Soc. C, 2663 (1969) (120) E. Schmitz, Chem. Ber. 97, 2521 (1964) (121) Can. Pat. 2,017,458 (Nov. 24, 1990), J.P. Schirmann, J. P. Pleuvry, and P. Tellier (to Atochem) D. Todd, in R. Adams, ed.,Organic Reactions, Vol. 4, John Wiley & Sons, Inc., New York, 1948, Chapt. 8. "Hydrazine and its Derivitives" If I can now be allowed to speculate a little into the reaction. So it appears that H2O2 can much more readily attack imines, R2C=NH, than it can plain ammonia? I thought I came across a book that stated that there is no reaction between H2O2 and imines without the use of some acetamide, which serves as a catalyst. One wonders why nitrite forms rather than nitrate, since the H2O2 should readily oxidize the nitrite. One possibility may be that nitrite is initially the predominant product, but as nitrate begins to be formed, it alters the equilibrium, allowing formation of more nitrogen dioxide rather than nitric oxide. (2) NO2[-] + (2) H[+]aq <==> H2O + NO + NO2 NO2[-] + NO3[-] + (2) H[+]aq <==> H2O + (2)NO2 Nitrogen dioxide is known to be able to attack ammonia. (2)NO2 + (2)NO + (4)NH3 --> (2)NH4NO2 + (2)H2O + (2)N2 In any case, solutions of ammonium nitrite decompose on heating to 60 to 70degC. Perhaps the very slow spontaneous reaction of aqueous ammonia with H2O2 takes place the small equilibrium with amide anions, NH2[-]. NH2[-] + H2O2 --> NH2* + OH* + OH[-] Even in NH4OH, there is a very slight equilibrium with amide ions. (2)NH3 <==> NH2[-] + NH4[+] In the base catalysed decomposition of hydrogen peroxide, the mechanisms is presumably H2O2 + OH[-] --> HOO[-] + H2O HOO[-] + H2O2 --> HOOOH + OH[-] HOOOH --> HOOO[-] + H[+]aq HOOO[-] --> OH[-] + O2 Is the dihydrogen trioxide intermediate reactive enough to oxidize ammonia? H2O3, also known as "trioxone" quickly decomposes into water and singlet [excited] oxygen. The reaction is actually reversible, but much less favorable. (2)H2O + O2* <==> H2O + H2O3 One question I would have though is what the dominant reaction in the spontaneous decomposition of NH4OH-H2O2 solutions. Is the reaction mainly the base catalysed decomposition of H2O2 into O2 ? Or is the main reaction the formation of ammonium nitrite? What proportion of the final product contains ammonium nitrate rather than nitrite? And why does the reaction proceed so much faster when heated? Relative to the energy required to break most chemical bonds, the boiling point of the solution is not really much hoter than room temperature. Why would this make so much difference? An interesting reaction from the literature of Marcellin Berthelot, dry ammonia gas reacts with the nitrogen dioxide and nitric oxide, at room temperature, (2)NO2 + (2)NO + (4)NH3 --> (2)NH4NO2 + (2)H2O + (2)N2 Solid ammonium nitrite inside a tube explodes if heated on a water bath to between 60-70degC. And the substance gradually decomposes at room temperature, slower if cold, or faster in aqueous solutions, forming nitrogen gas.

Add 20% concentrated hydrochloric acid. It will reduce all the chlorate to chloride, giving off chlorine gas, while leaving the perchlorate alone. KClO3 + (6)HCl --> KCl + (3)H2O + (3)Cl2 Despite having one more oxygen atom, perchloric acid, at least under 70% concentrated, is suprisingly inert from an oxidative point of view. It will not even be reduced by its reaction with zinc metal, whereas nitric acid is.

There are at least three other ways to make sodium metal, that do not involve electrolysis. Aluminum metal can actually reduce sodium hydroxide to sodium metal. Now I realise that many of you chemists will immediately say this is impossible, because "sodium is a more reactive element than aluminum", so here is a link showing pictures of the reaction, but with potassium being made from magnesium, instead of sodium from aluminum. http://sites.google....allic-potassium It also works if solid NaOH is ignited with Al powder in a metal container, and a lid placed over it to prevent reaction with air. The sodium cannot be obtained in a pure state by this method, however, as it is mixed with slag. But it is still charactaristically reactive with water. (6)NaOH + (4)Al --> (2)Al2O3 + (6)Na + (3)H2 but note that AlCl3 + (3)Na --> (3)NaCl + Al It is possible to prepare sodium metal by cautiously heating sodium azide in the absence of oxygen. (2)NaN3 --> (2)Na + (3)N2 Although lithium can burn in nitrogen, both sodium and potassium nitrides are very unstable. Sodium nitride decomposes into elemental sodium, giving off nitrogen gas, at only 87°C. Comparing the decomposition of other metal nitrides Similarly, the explosive decomposition of copper azide also results in the separation of the constituent elements, but this reaction happens for very different reasons. Cu(N3)2 --> Cu + (3)N2 But the same reaction with iron (which is dangerous) will result in iron nitride. (3)Fe(N3)2 --> Fe3N2 + (8)N2 The iron nitride can be decomposed to elemental iron and nitrogen gas above 800°C. Fe3N2 --> (3)Fe + N2 The decomposition of calcium azide is similar to that of iron. Ca(N3)2 decomposes above 110degC, explosively so over 140degC. The Ca3N2 that forms only decomposes at 1600degC, at which point the elemental calcium simultaneously vaporizes out with the nitrogen. Creative Way to Make Elemental Potassium? An idea for chemical preparation of elemental potassium, which does not require electric current. It would be impractical, but very creative. Not sure if all the reactions would work. Ca3N2 + (6)KCl --> (3)CaCl2 + (3)K2 + N2 Distilling calcium nitride with potassium chloride in with steel-walled distillation may cause potassium to boil out. This proposed reaction would make use of Le Chatelier's principle. Although potassium boils at 759°C, it is possible that molten potassium could be produced below this temperature. (6)CaCl2 + Ti3N4 --> (2)Ca3N2 + (3)TiCl4 The titanium nitride (m.p. 2930°C) would be crushed into a fine powder and distilled under intense heat with calcium chloride. Titanium tetrachloride (TiCl4) is a liquid which boils at only 137 °C. (3)TiI4 + (16)NH3 --> Ti3N4 + (12)NH4I I think titanium tetraiodide (b.p. 377 °C) could be reacted with anhydrous ammonia gas to form titanium nitride and ammonium iodide. I am not sure if the NH3 could be bubbled into molten TiI4, or if the TiI4 would need to be in the vapor phase, with the intense heat required for the reaction. The reaction would be expected to procede because TiI4 is very acidic, and because the titanium-nitrogen bonds are stronger than titanium-iodide. Wikipedia claims that TiCl4 "with ammonia, titanium nitride is formed"; this is not surprising since TiCl4 reacts with water to form titanium dioxide and hydrogen chloride. Titanium tetraiodide melts at 150 °C. It can be prepared from easily obtainable materials: (3) TiO2 + (4) AlI3 --> (3)TiI4 + (2)Al2O3

A mix of nitric and hydrochloric acids, known as "aqua regia", is well known for being able to dissolve gold. There is, however, a different acid combination that can be used to dissolve gold, and the interesting thing about this reaction is that the gold will reappear after the dissolved solution of gold is diluted in water! This reaction, however, involves much more concentrated acids, and the procedure is much more dangerous. So I do not suggest you try this reaction unless you have a good chemistry background and know about the proper safety precautions. This post is more for information purposes. Dissolving Gold with Concentrated Nitric and Sulfuric Acids A hot mixture of concentrated nitric and sulfuric acids can dissolve gold, with lower oxides of nitrogen forming. Addition of water caused the gold to precipitate back out in metallic form, but if a solution of permanganate is used instead, the gold remain dissolved. Reynolds, later by Spiller Chemical engineering, Volume 2, p316 The text also mentions that even concentrated mixtures of nitric and phosphoric acid attacks gold at room temperature, although the reaction is very slow unless heated. The reaction is probably: (3)Au + (3)NO3[-] + (18)H[+] --> (3)Au[+3] + (3)NO[+] + (6)H3O[+] The nitrosyl ion, NO[+], exists in the form of nitrosyl sulfuric acid, ONOSO3H. The nitronium cations, NO2[+], which form in equilibrium in concentrated nitric acid solutions, probably initially attack the gold, creating nitrogen dioxide. Basically, Au + (3)NO2[+] --> Au[+3] + (3)NO2 The nitrogen dioxide produced would likely remain in the concentrated acid, (2)NO2 + (3)H2SO4 --> NO[+]HSO4[-] + NO2[+]HSO4[-] + H3O[+]HSO4[-] and the nitronium ions formed from the NO2 would then attack more gold. Excess sulfuric acid needs to be used. This is an equilibrium reaction, and the gold is not going to dissolve easily. The mixture needs to be extremely acidic. Even a 1:1 ratio of 70% HNO3 to 95% H2SO4 is not going to be concentrated enough. For good results, use a 1:10 rato of 70% nitric acid to 98.5% concentrated sulfuric. Essentially, there can be no water in the reaction! Even in the hot boiling mixed acids, the gold takes several minutes to dissolve. The NO[+] ion hydrolyzes (reacts with) water to form nitrous acid. NO[+] + (2)H2O --> HNO2 + H3O[+] Nitrous acid is fairly reactive, and can act as either a reducing or an oxidizing agent. It will reduce the dissolved gold (Au+3) to elemental form (Au). This explains why the gold precipitates back out when the reaction is diluted with water. (2)Au[+3] + (3)H2O (3)HNO2aq --> (2)Au + (6)H[+]aq + (3)HNO3aq (note that "aq", which stands for "aqueous", means it is dissolved in water) If fuming nitric acid is added to the reaction containing the dissolved gold, the gold will solidify out as a purple solid. The gold is probably still in its elemental form, but small particle sizes of gold are known to exhibit strong colorations, from red to purple. Nitrous acid is unstable, and only exists in the form of solutions which gradually degrade after several minutes. Solutions of nitrous acid exist in equilibrium with nitrogen dioxide and nitric oxide, the latter of which is an unstable radical which can either react with the oxygen in air to form more nitrogen dioxide, or if left on its own will disproportionate into nitrogen dioxide and nitrous oxide after several minutes. (2)HNO2 <==> H2O + NO2 + NO (3)NO --> N2O + NO2 In the reaction, (2)Au + (3)NO3[-] + (18)H[+] --> (2)Au[+3] + (3)NO[+] + (6)H3O[+], sulfate ions are not shown because they do not directly take place in the reaction. The literature even states that phosphoric acid can be used in place of the sulfuric acid. The above reaction is in ionic form. Some of you may prefer to see it in the form: (2)Au + (3)HNO3 + (15)H2SO4 --> (2)Au(SO4H)3 + (3)NOSO4H + (6)H2SO4*H2O Note that the "Au(SO4H)3" only exists in the solution, it cannot be isolated. Gold trinitrate, if it even exists, would also be nearly impossible to obtain as a pure solid. Gold trinitrate only exists in highly concentrated solutions of nitric acid. When these solutions are diluted with water, auric oxide precipitates out. Similarly, auric oxide only only dissolves in very concentrated acids, since it is only very weakly basic. Au2O3 + (9)HNO3 <==> (2)Au(NO3)3 + (3)HNO3*H2O The reaction is more interesting from a chemical perspective than a practical way to refine out gold. Nevertheless, the reaction may be useful to directly dissolve gold-silver alloys, without having to go to the trouble of inquartation, since aqua regia only dissolves such alloys with extreme difficulty. Procedure and Precautions: Yes, it is extremely dangerous. The dangers of using concentrated mixed acids are commonly taken for granted among those that frequently perform nitrations. Obviously those unfamiliar with such procedures should think twice before handling such high concentrations of acid. More details about the reaction. The concentrated acid mix that contains the dissolved gold should be gradually transfered into the larger bowl of water using a 10mL glass transfer pipette. You will also need a rubber pipette suction bulb. For those of you unfamiliar with this tool, it is basically like a turkey baster that is used to suck up a small quantity of liquid, then move it to another container. The pipette can be bought here: http://www.pelletlab.com/pipette Using the pipette to slowly add the acid mixture to the water is important for two reasons. First, safety. Water should never be added to concentrated acid, since this can result in the acid spraying up. Neither should the acid be poured into the water, because of the possibility of an accidental spill or splashing, and because it can be hard to control the rate that the liquid is poured in. Adding the acid in too fast can lead to overheating, which could result in boiling/splashing in the water. Second, it is important that each small portion of the acid quickly be diluted with as much excess water as possible. This will help prevent the gaseous nitrogen oxides (NO and NO2) from escaping. Although nitrosylsulfuric acid reacts with excess water to form a solution of nitrous acid, if not enough water is used nitrogen oxides will bubble out instead. There will inevitibly be some loses of nitrogen oxides, in the form of some bubbling and some brown gas being given off. Unfortunately, when some of the nitrogen oxides escape, there will not be enough nitrous acid to completely reduce the gold. After neutralizing, all the gold will still precipitate out, but a small portion of it will be in the form of hydrated gold oxide, Au2O3. If the gold is going to later be melted, the gold oxide should not pose any problems, as the compound decomposes to the pure metal at 160°C, giving off oxygen gas. One other note of warning, unless the gold oxide has been completely reduced, it should not be reacted with ammonia, as this will form the dangerous sensitive explosive known as "fulminating gold". In the event that the acid solution was previously boiled with ammonium sulfate to prevent precipitation of the gold, fulminating gold can result upon neutralization if too much ammonium sulfate was added. More safety information: Only use small quantities of mixed acids at a time. Be aware that with concentrated acids, even tiny drops can splash out and result in painful burns on exposed skin. To get some understanding of these dangers, try pouring cranberry into a glass, wearing a clean white long-sleaved shirt. Even with cautious pouring, you are likely to find one or two tiny little red stains on the sleeves afterwards, even though you were not aware of any splashing while the juice was being poured. If this was concentrated acid, painful burns would have been felt. You may desire to cover your shoes with a plastic bags and a rubber band, so that if any of the acid spills onto the floor, it will not seap into your shoes. Protective shoe coverings can also be purchased: http://www.labsafety.com/search/shoe%2Bcovers/ If you choose to wear rubber boots instead, it is advised that the top of the rubber be tied tight around your legs, so that if any of the acid is spilled on you, it will not drip down into the boots and collect in a puddle. If the acid is in contact with your skin for more than a few seconds, the burns will be much more severe. http://www.amazon.com/b?ie=UTF8&node=393294011 A boiling mixture of concentrated nitric and sulfuric acids is extremely dangerous, much more so than 70% concentrated sulfuric acid, for example. The chemistry of this mixture presents several unique hazards. Extremely concentrated sulfuric is a strong dehydrating agent, that will turn anything organic, such as a strip of paper or your skin, into black char immediately on contact. A note about treating concentrated nitric acid burns, after you immediately rinse the affected area with plenty of water, and neutralize with sodium bicarbonate solution, there is special recommendation for concentrated nitric acid burns. Use a swab dipped in chlorine bleach to gently scrub the affected area. Some of the yellow color from the burn should be absorbed onto the cotton swab. Continue to scrubbing with fresh swabs until no more yellow can be absorbed onto the cotton. Then rinse well in soapy water. Doing this will help remove some of the nitro compounds which have formed. These compounds act as allergens and greatly slow the healing process. In fact nitric acid burns take much longer to heal than sulfuric acid of the same concentration. The unique effects of concentrated nitric acid are due to the formation of nitronium ions, NO2[+], in equilibrium in the solution. The addition of highly concentrated sulfuric acid greatly enhances this equilibrium, and so the special burn effect of nitric acid will be greatly exaggerated by the acid mixture. In other words, it would be very important to treat the burns in the way described above, and the healing time is likely to be much longer. Further Information: One of the posters at "http://goldrefiningforum.com" stated that the "wet ash" method did indeed dissolve gold if the acids were concentrated enough, although he wrote that it was not a practical method at all. The extremely concentrated HNO3/H2SO4 mixture might be useful for dissolving gold-silver alloys, without the need for inquartation, since aqua regia only dissolves such alloys with extreme difficulty. Dissolving Gold with Manganese Dioxide Mixtures of manganese dioxide and sulfuric acid can also dissolve gold. The reaction is slower at room temperature, but rapid with heating. Permanganate and sulfuric acid after a few minutes also dissolve gold. (Allen 1872) the reaction with manganese dioxide and sulfuric acid is probably: (2)Au + (3)MnO2 + (3)H2SO4 --> Au2O3 + Mn(SO4)2 + (2)H2O where Mn(SO4)2 is manganese sulfate, and the gold oxide dissolves in the sulfuric acid. concentrated sulfuric acid still needs to be used, but it probably does not need to be quite so concentrated as required for the other reaction; a 70% concentration should be suitable. Other methods Ozone with sulfuric acid did not dissolve gold. It is known that aqueous solutions of bromine or chlorine can dissolve gold. A mixture of sulfuryl chloride and dimethylamine can dissolve gold, but not platinum. Sulfuryl chloride has the formula SO2Cl2, and can be made by reacting SO2 with dry Cl2, using an activated carbon catalyst at room temperature. It is highly corrosive and readily hydrolyses with water.