In this video, we will investigate nitric acid, how it is made in the lab, its physical and chemical properties and some of its reactions, as well as its uses. Let’s start by having a look at some of the commercial and industrial applications and uses of nitric acid.
Nitric acid, or HNO3, is a highly useful and common starting substance for many products in industry. It is made from the H+ or hydrogen ion bonded to the NO3− or nitrate ion. An old Latin name for this acid is aqua fortis, which means strong water. Although this compound is not water, it can dissolve in water and is colorless like water in its very pure form. The first recorded synthesis of nitric acid was approximately 800 AD by an alchemist. Since then, mankind has found many uses for this acid.
It is used in the production of dyes, which we use to change the color of fabrics; in the production of drugs and medicines; to purify and treat gold and silver as well as carbon nanotubes; to etch metal surfaces, which is the treatment of certain parts of a metal surface with a strong acid to create a design, for example, in some medals; as an oxidizing agent in rocket fuel. Nitric acid is used as a reagent in elemental analysis in the laboratory. Elemental analysis is when a scientist determines which elements are present in a sample and in what quantities they are present.
Nitric acid is also used in the manufacture of explosives. Perhaps you’ve heard of TNT or trinitrotoluene. This substance, along with nitroglycerin, are the main components of dynamite. Perhaps one of the most important uses of nitric acid is in the production of fertilizers. Many fertilizers are salts of nitric acid. This equation shows the production of the fertilizer component ammonium nitrate. Ammonia and concentrated nitric acid undergo a combination reaction to produce this fertilizer product.
Now we know how useful nitric acid is, let’s have a look at how it is made in the lab. Here is the experimental setup used to make nitric acid. Concentrated sulfuric acid and a dry nitrate salt, for example, sodium nitrate, are added together in a glass retort vessel. The mixture is heated with a Bunsen burner. The following reaction occurs. The products are sodium bisulfate and nitric acid. This nitric acid is given off as brown fumes in the retort vessel. The nitric acid is distilled off, cooled, and collected. Now the nitric acid product is highly concentrated and fuming.
Pure nitric acid is liquid, colorless, and clear. However, it can decompose into a yellow form. The yellow color of nitric acid can come from the nitrogen dioxide, which is one of the decomposition products. Heat or light can cause this decomposition, and that is why nitric acid is stored in brown bottles.
Let’s now turn our attention to the physical properties of nitric acid. We have seen that nitric acid in its pure form is a colorless liquid at room temperature, although commercially available nitric acid is yellow brown because of partial decomposition to nitrogen dioxide gas. Its density is 1.51 grams per centimeter cubed. This is for the anhydrous or water-free version of nitric acid at 20 degrees Celsius. Commercially purchased concentrated nitric acid is not anhydrous but is about 68 percent by mass of nitric acid dissolved in water. Nitric acid is highly water-soluble and completely miscible in any proportions. This is because of the ionic nature of the components of nitric acid and the polar nature of water.
The melting point of pure nitric acid is negative 42 degrees Celsius and the boiling point 83 degrees Celsius. In its pure solid form, nitric acid is white. Note, however, that commercially available nitric acid has different melting and boiling points because of the presence of water. Nitric acid has a pungent suffocating odor, especially in its pure form, when it is fuming. It is toxic to inhale and corrosive and oxidizing in nature, causing burns and damages to biological tissues as well as damage to metals. We will investigate the oxidizing nature of nitric acid a bit further later on.
Lastly, nitric acid is a strong acid. Its oxidizing nature as well as its strength are really chemical properties. Let’s have a look now at some chemical properties of this acid.
We said nitric acid is a strong acid. This means it dissociates or ionizes into its ions completely in water. We’ve already seen how it breaks apart into the hydrogen ion and the nitrate ion. And, and the hydrogen ion can react with water to produce the hydronium ion H3O+. The hydronium ions cause blue litmus paper to turn red, confirming that this is indeed an acid. We’ve also seen that nitric acid decomposes in the presence of heat or sunlight. We saw this equation earlier. And we learned that this decomposition transforms colorless liquid nitric acid into a yellow-brown color because of the presence of nitrogen dioxide gas. When this decomposition occurs as the result of heat, we call it a thermal decomposition.
Now, because nitric acid is an acid, it can react with a base. The general equation of the reaction between an acid and a base is acid plus base gives salt plus water. In the case of nitric acid, it can react with a base, specifically an alkali base — for example, sodium hydroxide — to produce the salt sodium nitrate and water. Nitric acid can also react with a base that is a metal oxide or a basic oxide, for example, potassium oxide, producing the salt potassium nitrate and water. Many acids, including nitric acid, can react with a metal carbonate to produce salt, water, and carbon dioxide. Here is an example of this reaction using nitric acid and calcium carbonate. The salt product is calcium nitrate, and carbon dioxide is given off as a gas.
Now there are many types of reactions that nitric acid can undergo. We’ve only looked at a few here. Let’s look at one last type of reaction in some depth. We’ll look at the oxidizing nature of nitric acid on metals. Nitric acid is a powerful oxidizing agent. It can even react explosively with some nonmetal compounds. The products produced between the reaction of nitric acid and a metal depends on two things: the concentration of nitric acid and the nature of the metal involved in the reaction.
Now, most acids can react with metals to produce a salt and hydrogen. Nitric acid can undergo these types of reactions when it is dilute. For example, dilute nitric acid can react with magnesium metal to form the salt magnesium nitrate and hydrogen, which is given off as a gas. Let’s look at how the metal has been oxidized. Magnesium initially had an oxidation state of zero in its elemental form. And after the reaction, it has been oxidized to an oxidation state of plus two. Hydrogen initially had an oxidation state of plus one in the nitric acid. And it was reduced to an oxidation state of zero.
This reaction that we have seen here will compete in the reaction vessel with another reaction between nitric acid and magnesium. Here is the equation for the competing reaction. Again, we have dilute nitric acid reacting with magnesium and forming the same salt, magnesium nitrate. However, there are two different products other than hydrogen in this case: NO, which is nitrogen monoxide, and water. Nitrogen monoxide is commonly referred to as nitric oxide. It is given off as a gas, and this gas is colorless. Magnesium in this reaction undergoes the same oxidation as in the previous reaction.
The nitrogen monoxide, or nitric oxide, can react further with oxygen to produce the brown-yellow nitrogen dioxide gas. But the important thing to remember is that dilute nitric acid produces nitrogen monoxide gas. So for now, we’ll remove the extra equation with the oxygen. Let’s write a general equation for this reaction. Here is the general equation. Dilute nitric acid plus metal can give salt, nitric oxide or nitrogen monoxide, plus water. And the competing equation has also been boxed in green, so you can compare the two.
These are both for dilute nitric acid. What about concentrated nitric acid? The equation for the reaction of concentrated nitric acid with a metal is very similar to the bottom equation. So let’s erase the top equation for some space. In the case of concentrated nitric acid reacting with magnesium, everything is exactly the same as for dilute nitric acid and magnesium except one of the products is nitrogen dioxide gas, which we know is a yellow-brown gas. Can you see that this gas is formed directly in the reaction and not as an afterwards reaction, as in the case of dilute nitric acid?
Again, we have the same salt being formed, magnesium nitrate. And water is also a product. Balancing the two equations will obviously be different because of the two different gaseous products, but everything else is the same. Let’s write a general equation for the reaction of concentrated nitric acid with a metal. The general equation will be concentrated nitric acid plus metal gives us salt, nitrogen dioxide, plus water. And again, the magnesium is oxidized in the same manner as in the other reactions.
What do you think the difference will be in the observations of these two reactions? Well, the only difference is the gas that is liberated. One is brown yellow, and the other is colorless. What if we replace magnesium with another metal, for example, copper? The reaction equations will be identical to those of magnesium if the metal also forms a two plus charge. However, in this case, the product is copper nitrate and not magnesium nitrate. The copper(II) nitrate solution will be a deep-blue color.
Will all metals react with nitric acid? The answer is no. Gold, platinum, and other platinum group metals do not react with nitric acid. However, if the nitric acid is mixed with another acid like hydrochloric acid, these together can dissolve these metals.
So far, we have learnt about the uses of nitric acid, its production, its physical properties, and some of its chemical reactions. But how can we test for the presence of nitric acid or another nitrate compound?
If we have a solution which we suspect contains nitrate NO3− ions, we can add to it a solution of iron(II) sulfate, FeSO4, also known as ferrous sulfate. These two are mixed together gently, and then a smaller amount of concentrated sulfuric acid is added dropwise down the sides of the test tube. The careful addition of the denser sulfuric acid will cause it to sink to the bottom of the test tube. Two separate layers will be visible. A brown ring will appear at the interface of the acid and reactant solutions if there are nitrate ions present. This ring will disappear by heat or shaking.
Here is the equation for the reaction that occurs at the interface of the two solutions. The iron in the ferrous sulfate with a plus two oxidation state is oxidized to plus three. Fe3+ is a brown-yellow ion. This forms part of the brown ring. Also, the nitrogen monoxide or nitric oxide that is produced reacts further to produce other complex compounds. And these together form the brown ring, confirming the presence of nitrate ions in the reactant.
In this example, we used sodium nitrate. We’re going to look at one more test. And that is for the nitrite ion. If we take a solution of a salt, which we suspect contains nitrite ions, NO2−, we can add acidified potassium permanganate, which is KMnO4 plus a few drops of concentrated sulfuric acid. The two solutions together will be a pale-purple color from the potassium permanganate. If there are indeed nitrite ions present, the purple color will disappear. Here is the equation for the reaction, again using sodium nitrite as the example.
The purple permanganate ion is converted to the colorless manganese two plus ion. This is also a redox reaction, but we won’t go into the details here. This test for nitrite ions is not unique to these ions. Potassium permanganate can be decolored by other ions too, so this test is not proof enough of the presence of nitrite ions. We would need to do another test to confirm that there are indeed nitrite ions present in the solution.
We’ve learnt a lot about nitric acid. Let’s summarize some of the more important points. We learnt that there are many industrial and commercial uses of nitric acid. It is a common starting material in the manufacture of explosives, fertilizers, dyes, and many other products. We learnt that in the lab it can be prepared from sodium nitrate and concentrated sulfuric acid.
We investigated some of its physical properties, which include its pungent aroma, its toxicity, the fact that it is corrosive and highly water-soluble, and that in its pure form it is a colorless liquid at room temperature. We saw that it decomposes in heat or sunlight to produce the yellow-brown gas nitrogen dioxide and that often nitric acid appears yellow brown rather than colorless because of this decomposition product. And commercially available nitric acid is yellow brown and a 68 percent by mass nitric acid in water.
We investigated some of the reactions of nitric acid. We learnt that it is a strong acid and a strong oxidizer. And interestingly, we learnt that the concentration of nitric acid when it reacts with a metal can influence the products formed. Nitrogen monoxide is produced with a dilute acid and nitrogen dioxide with a concentrated acid. And finally, we looked at the brown ring test for nitrate ions and the potassium permanganate test for nitrite ions, although the potassium permanganate test is not conclusive enough.