Preparation And Properties Of Nitric Acid (1959)

[Music] in this film we will examine three different phases of the chemistry of nitric acid its production its activity as an acid and its activity as an oxidizing agent pure nitric acid is a colorless liquid boiling at 86° Centigrade with some decomposition into Brown nitrogen dioxide gas water and oxygen this relatively low boiling point makes it practical to prepare nitric acid in the laboratory by means of the reaction between sodium nitrate and sulfuric acid to this retort we have added about 20 G of sodium nitrate sulfuric acid will be added through the funnel and the mixture heated a retort is used for this experiment because it is an all glass piece of apparatus concentrated nitric acid when hot rapidly attacks both Rubber and cork so that the use of a conventional flask with Stoppers and connections is unwise as the mixture is heated nitric acid will be formed will distill over through the neck of the retor into this flask where the cold water in this Beaker will cause the vapors to condense back into liquid nitric acid after a few minutes of heating the retort has this appearance I believe you can see the light brownish color of the decomposition products in the upper part of the retort the bubbles of nitric acid Vapor Rising from the mass of sodium nitrate and sulfuric acid at the end of the retoric tube drops of liquid nitric acid fall slowly into the erand Meer flat the two react reactions of interest in this experiment are first that for the production of nitric acid from sodium nitrate and sulfuric acid and secondly the equation for the decomposition of nitric acid during distillation the first equation uh is not a complicated reaction since it's not an oxidation reduction equation note that the product is nitric acid which distills and so sodium hydrogen sulfate in order to use the second acid hydrogen in the sulfuric acid uh we would need to go to such a high temperature that the nitric acid uh would be uh long since decomposed the pure nitric acid during distillation decomposes to some extent into Nitric o nitrogen dioxide water and oxygen and this nitrogen dioxide being Brown accounts for the brownish color of the nitric acid formed in this reaction the ordinary concentrated nitric acid used in industry or uh in the laboratory is really a solution containing about 68% nitric acid molecules by weight and about 32% water a solution of this material freshly prepared is in the bottle on on the left this material is comparatively stable although when it uh stands around for some time decomposition into NO2 the brown gas occurs and accounts for the yellow color which can be seen in the bottle on the right we will now prepare some very dilute nitric acid in order that we may study the properties of nitric acid in this form in this graduated cylinder we have 95 mL of water I will add this water to the beaker in this smaller graduate we have 5 milliliters of concentrated nitric acid adding that acid to the water in the beaker reduces the concentration of the acid by a factor of 20 the acid is now about 1/ 12th of its former strength now we're going to test this acid uh in its reactions with indicators and also in its reactions with active metals in this test tube we have a few milliliters of distilled water and to this I'm going to add a few drops of methyl red indicator the indicator has assumed its red form in the distilled water and we're now going to cause it to change to it the color that it has in alkaline solution or basic solution by adding a drop or two of ammonium hydroxide solution this yellow color is characteristic of the indicator in alkaline solution we will now add a few drops of our very dilute nitric acid to the test tube containing the methyl red indic here the color instantly changes back to the reddish purple form characteristic of the indicator in acid solution in this reaction then delute nitric acid has behaved as a typical acid another reaction typical of most acids is the reaction of the acid with an active metal to produce hydrogen in this Beaker we have our very dilute nitric acid which we prepared a few moments before and I'm going to pour 40 millit of this acid into the graduated cylinder in this bottle we have some very finely divided magnesium metal magnesium turnings We'll add the acid to the magnesium and then immediately put the glass plate in place we now permit the acid and the Magnesium to react for just a few moments f a splint and test the gas in the bottle and the sharp bark characteristic of the explosion of hydrogen is obtained so hydrogen is a product of the reaction of very dilute nitric acid with magnesium when a piece of filter paper is moistened with a concentrated solution of ammonium hydroxide and then held to the mouth of the bottle of nitric acid which we just made if you minutes ago you'll notice the formation of dense white clouds of ammonium nitrate salt the formation of this salt indicates that nitric acid fumes are being given off by the concentrated nitric acid in the flask these fumes are reacting with the ammonia Vapors from the filter paper in the first uh experiment we saw that dilute nitric acid very dilute uh acts principally as an acid therefore its reaction with magnesium which we demonstrated to produce hydrogen may be written in this fashion uh magnesium ion and hydrogen gas being formed next we studi the reaction of concentrated nitric acid Vapors with ammonia Vapors in each case the vapors were liberated from concentrated Solutions of the corresponding solute white fumes were formed in this reaction these white fumes are ammonium nitrate salt which has this formula when we examine the oxidizing activity of nitric acid we need to do so at several different concentrations since the reduction product of the nitric acid varies with the concentration first we've placed a few copper turnings in the test tube and we'll add to them a few milliliters of the concentrated or 16 normal nitric acid the reaction takes place rapidly the reduction product of the nitric acid is this brown gas nitrogen dioxide NO2 and note that even after the reaction has been continuing for a few moments the brown NO2 fills the tube right down to the very bottom indicating that this gas is the direct product of the reaction the blue solution of course is copper nitrate a similar reaction will be observed when we treat zinc of which we placed a few pieces in the test tub with concentrated or 16 normal nitric acid of course this time the solution will not be blue the reaction is very vigorous and again note that the brown fumes of NO2 completely fill the tube and are the initial product of the reaction when iron is treated with concentrated nitric acid however no reaction takes place or very little because the surface of the iron becomes passive this is probably due to the formation of a coating of oxide in the test tube we've placed several iron Nails We'll add the concentrated nitric acid a little if anything happens the solution becomes slightly colored due to the solution of some of the iron oxide on the surface of the nails in the acid but uh the violent reaction that we saw before is simply not happening in this sequence of experiments we have Illustrated several chemical properties of concentrated nitric acid as it reacts as an oxidizing [Applause] agent first we examine the reaction of concentrated nitric acid with copper nitric acid even in concentrated solution is a strong acid and extensively ionized so the active oxidizing ingredient is the nitrate ion we saw that the reduction product was the brown gas nitrogen dioxide in order to balance this partial equation we must add two hydrogen ions to this side and form one molecule of water on the right this side now has a net charge of one plus this side is neutral so one electron must be added here copper was oxidized to the ceric ion and yielded two electrons this is equation then must be multiplied by two for electrical balance and the two partials combined giving us copper plus four hydrogen ions plus two nitrate ions producing cupric ion two molecules of NO2 and two molecules of water when zinc was substituted for copper much the same results were obtained so that the partial for nitric acid Remains the Same but now a new partial for zinc namely zinc gives zinc ion plus two electrons must be substituted for the copper partial in this equation I don't think it's necessary to combine these two partials since the results will be entirely similar uh to the reaction with copper however when iron was used in the form of the nails we observed that no reaction took place uh we turn now to a consideration of nitric acid as an oxidizing agent uh in more dilute concentration in each test tube we have a few pieces of Mossy zinc in the bottle on this side we have some six normal nitric acid and in this bottle some 16 normal nitrc now when six normal nitric acid performs as an oxidizing agent the initial product of the reaction is the colorless gas nitric oxide n o this reacts rapidly with oxygen in the air forming NO2 however notice that the the color of the gas in this test tube is not very dark brown it consists mostly of the colorless and all for purposes of comparison We'll add some 16 normal Nitric to the zinc in this test tube you should observe the much deeper color of the NO2 formed when the concentrated nitric ax on the zinc and the almost colorless n formed by the nitric acid when dilute or six normal nitric acid was used as an oxidizing agent we observed that the gas produced was the colorless nitric oxide with the formula n o this is a different product and requires a different partial equation to illustrate its formation in this case one nitrate ion reacts with four hydrogen ions require three electrons and produces nitric oxide and water the partial equation for the reducing agent zinc is the same as before multiplying by the appropriate factors and combining the partials we obtain this equation for the overall reaction between zinc and six normal nitric acid nitric oxide is colorless and the gas originally formed in the tube had a brown color this was due to the reaction of the nitric oxide produced by the reaction with oxygen in the air this reaction occurs spontaneously at room temperature and forms the brown nitrogen dioxide this equation accounts for the color of the gas originally present in the test tube finally we will examine the oxidizing character istics of very dilute or about one normal nitric acid when it reacts with an active metal such as zinc in the test tube we have some granulated zinc metal divided into very fine pieces in the graduated cylinder we've prepared a solution of nitric acid one normal in strength the reaction in this case is slow so after adding the nitric acid to the zc a few bubbles are seen to rise and we will now wait about 5 minutes for the reaction to proceed we'll now decant the solution in the test tube from the remaining zinc which was present in excess into this Beaker and make the solution basic with sodium hydroxide solution the white precipitate which forms is zinc hydroxide and need not concern us we're more interested though in any Vapors or gas that might be given off by this solution and we can test those with this piece of red litness paper which you see turns blue as we hold it in the air above the solution in the beaker this indicates that ammonia gas is being evolved from the solution and that the ammonium ion was a product of the reduction of the dilute nitric acid by the zinc the reaction of a very dilute or about one normal nitric acid with zinc an active metal uh proceeds in yet a different fashion in this case the nitrate ion is reduced to the ammonium ion the part IAL equation requiring 10 hydrogen and eight electrons on this side and three molecules of water on this side for balancing the partial equation for zinc is as we have seen previously multiplying this partial by four we can obtain an overall equation for the reaction of zinc and very dilute nitric acid with the production of zinc ions and the ammonium ions we then made the the solution basing with sodium hydroxide obtaining zinc hydroxide which is of no concern and also obtaining the reaction between the hydroxide ion from the sodium hydroxide and the ammonium ion from the reduction of the nitrate ion which produced ammonia gas which was evolved from the solution we then tested the ammonia with litmus paper and obtained the characteristic color change from red to blue uh a test for the ammonium ion we can now test the very concentrated nitric acid which we prepared a few minutes ago from the sodium nitrate and the sulfuric acid as an oxidizing agent on substances generally difficult to oxidize We'll add some of this acid to this test tube which contains a small Cork and from to this test which contains a few pieces of sulfur and then we'll warm the tubes gently with the bunson burn now let's pause for a few minutes to permit these reactions uh to proceed to completion after about 5 minutes you can see the large quantity of brown fumes here which indicate that an oxidation reduction reaction has taken place uh the poor old cork is pretty much disappeared and is uh has been reduced to sort of a gooey mass of material the sulfur and the sulfuric acid are still bubbling slowly here and the oxidation product of the sulfur should be sulfuric acid or the sulfate ion uh in order to determine whether or not this is the case we'll decant the solution from The sulfur into this tube containing some water and then make the traditional barium chloride test for the presence of the sulfate ion you'll recall that sulfate Ion with barium chloride gives a fight precipitate obviously then in this reaction the reduction product of the nitric acid was the brown gas NO2 and the oxidation product of the sulfur was the sulfate ion the equation for the reaction between concentrated nitric acid and the sulfur may be derived as follows the nitric acid partial is the same as the one previously used for concentrated nitric acid the sulfur partial was demonstrated to produce sulfate ion and uh may be derived in this fashion using water as a source of oxygen combining the two partials gives us this for the overall equation we then test for the presence of the sulfate Ion with the barium ion obtaining a precipitate of berium sulfate in solution nitrates are ionized to give a metal ion if it's a salt a nitrate or hydrogen ion from nitric acid and the nitrate ion it's necessary to have a good test for the presence of the nitrate ion in solution and this test will be the next thing to be demonstrated this test tube contains one drop of six normal nitric acid as a source of nitrate ion now to this tube I'm going to add a few milliliters of a solution of feros sulfate which has been freshly prepared and this solution will then be agitated for mixing next we take some concentrated sulfuric acid and we tilt the tube so that it's possible to pour the sulfuric acid carefully down the side of the tube sulfuric acid is much denser than feros sufate solution and passes beneath the feros sufate solution to the bottom of the tube where it forms a layer of sulfuric acid where the two solutions meat a brown ring is formed and this brown ring constitutes the test for the nitrate ion the equations for the reactions involved in the brown ring test for the nitrate ion are on the board we had a dilute solution of nitrate iion made acid by the sulfuric acid and when we added the feros sulfate the feros ion reduced the nitrate ion to nitric oxide and was itself oxidized to the feric ion the complete equation for this reaction uh appears here the nitric oxide formed then reacted with additional fer o i present in the solution to give a feros nitric oxide side complex which probably has the formula Fe n o two plus charges remaining from the feros and which has a brown color this color appears at the junction between the sulfuric acid layer and the feros sulfate layer in this film we have demonstrated a laboratory method for the preparation of nitric acid we have studied its activity as an acid and as an oxidizing agent in various concentrations and we have developed a test for the presence of the nitrate ion the equations for the reactions involved have also been discussed [Music]