Lesson Explainer: Properties of Nitric Acid Chemistry

In this explainer, we will learn how to describe the physical and chemical properties and uses of nitric acid.

Nitric acid (HNO3) is a colorless mineral acid found commonly in a dilute form in school laboratories. It is highly corrosive in its concentrated form, which contains around 68% nitric acid. The concentrated form is a liquid at room temperature and boils at around 120C.

The molecular structure of nitric acid can be shown through different resonance structures. The actual molecular structure will be somewhere between the two major resonance structures, shown in the diagram below.


It is possible to prepare nitric acid in the laboratory, although care should be taken due to the corrosive nature of the concentrated acids involved. The experimental setup is shown in the diagram below.

In this preparation, potassium (or sodium) nitrate is added to sulfuric acid. Once the potassium nitrate and concentrated sulfuric acid have been mixed, they are heated gently ensuring that the temperature does not exceed 100C, which could decompose the nitric acid being produced. The nitric acid, the distillate in this reaction, is then collected in a round-bottom flask and kept cool in an ice bath.

The chemical equation for this preparation of nitric acid is KNO()+HSO()HNO()+KHSO()32434saqlaq

Example 1: Identifying the Correct Chemical Formula for the Reactants and Products in the Synthesis of Nitric Acid

Nitric acid can be produced by mixing sodium nitrate with sulfuric acid and heating the mixture to 83C. Which of the following is the correct chemical equation for this reaction?

  1. NaNO+HSONaSO+HNO224323
  2. NaNO+HSONaHSO+HNO32443
  3. Na(NO)+2HSO2NaSO+HNO+H23244232
  4. NaNO+HSONaSO+HNO3443
  5. Na(NO)+HSONaSO+2HNO322443


This question is testing our knowledge of the correct chemical formulas for the reactants and products involved in the synthesis of nitric acid.

Sodium nitrate has the chemical formula NaNO3, and so answers C and E are incorrect. The sodium compound in equation A is actually sodium nitrite and not nitrate. This leaves us with options B and D.

D has an incorrect formula for sulfuric acid, which is correctly written as HSO24 in answer B, the correct answer.

In the preparation experiment above, we avoided high temperatures when dealing with nitric acid. Care should always be taken to avoid heating concentrated nitric acid as it can decompose to produce toxic brown fumes of nitrogen dioxide: 4HNO()4NO()+O()+2HO()3222lggl

This decomposition will occur to a much lesser extent even at room temperature, with concentrated nitric acid also being subject to light decomposition. For these two reasons, the acid should be kept in brown reagent bottles in a cool storeroom. The brown color of the nitrogen dioxide produced in this decomposition is what gives older samples of concentrated nitric acid a characteristic yellow color. The other gas produced in the decomposition is oxygen, along with liquid water.

Example 2: Calculating the Change in the Oxidation State of Nitrogen in the Oxidation of Graphite

Nitric acid is a powerful oxidizing agent and will even react with elemental carbon as the allotrope graphite, as shown: 3C+4HNO3CO+4NO+2HOgraphite322

By what value does the oxidation state of nitrogen change during this reaction?


Hydrogen will always have an oxidation number of +1 unless it is part of a hydride such as NaH. Oxygen is more electronegative than nitrogen, and so in the nitrate ion it will have the oxidation number of 2.

As such, the oxidation number of nitrogen in nitric acid is +5, given that the oxidation number of oxygen must be 2 and that of hydrogen must be +1.

In the products, we see nitrogen in the form of nitric oxide. Again, given that oxygen must have an oxidation number of 2, nitrogen must have an oxidation number of +2. This means that the difference between +5 and +2 is 3, which is the correct answer.

Nitric acid will react with metals. However, the products of the reaction depend very much on the metal in question and the concentration of nitric acid.

When considering metals that are more active than hydrogen, such as magnesium and zinc, very dilute nitric acid will react in a similar fashion to other mineral acids, producing a salt and hydrogen gas: Mg()+2HNO()Mg(NO)()+H()saqaqg3322

However, at greater concentrations, nitric acid can start to act as an oxidizing agent. We can observe nitric acid’s oxidizing nature in the following reaction between iron metal and more concentrated nitric acid to produce Fe3+ ions: Fe()+4HNO()Fe(NO)()+2HO()+NO()saqaqlg3332

For metals less active than hydrogen, nitric acid acts as an oxidizing agent, firstly oxidizing the metal to the oxide, which then reacts with the nitric acid to form the metal nitrate salt, nitric oxide, and water: 3Cu()+8HNO()3Cu(NO)()+4HO()+2NO()saqaqlg3322

Reactions of metals with concentrated nitric acid can produce greater oxidizing effects, such as the reaction of concentrated nitric acid with copper metal: Cu()+4HNO()Cu(NO)()+2HO()+2NO()slaqlg33222

Example 3: Ordering the Reactions of Magnesium Metal with Varying Concentrations of Nitric Acid

Magnesium can react with different concentrations of nitric acid to produce different products. Put the following reactions in order from least to most concentrated nitric acid used in the reaction: Mg+4HNOMg(NO)+2NO+2HO(1)Mg+2HNOMg(NO)+H(2)3Mg+8HNO3Mg(NO)+2NO+4HO(3)3322233223322


Although nitric acid behaves in a similar fashion to other mineral acids at low concentration, at higher concentrations it can act as an oxidizing agent.

Dilute mineral acids commonly react with metals to form a salt and hydrogen gas as can be seen in equation 2, and so the lowest concentration of nitric acid is used in equation 2.

As the concentration of nitric acid increases, the oxidizing nature can be more easily seen, and initially this results in the production of nitric oxide (NO).

As the concentration increases further, the nitric oxide is further oxidized to nitrogen dioxide (NO2). These two nitrogen species correspond to equations 3 and 1, respectively, giving us our final order of 2, 3, 1.

However, when iron, cobalt, chromium, nickel, and aluminum react with concentrated nitric acid, they form a layer of the metal oxide on the surface of the metal. This process is called passivation and protects the metal underneath from further reaction with the concentrated nitric acid.

Definition: Passivation

An active substance becomes passive through the formation of a passivating film.

How To: Performing the Brown Ring Test for Nitrates

It is possible to test for the presence of a nitrate ion (NO3) by performing the “brown ring” test.

A freshly prepared concentrated solution of iron(II) sulfate is mixed with the unidentified solution that may contain nitrate ions. After mixing in a test tube, a few drops of concentrated sulfuric acid are carefully allowed to run down the inside walls. When the concentrated sulfuric acid meets the mixture at the surface of the liquid, a brown ring appears, which then disappears when shaken or heated. There are two stages of reactions that occur in this test. Initially, nitric oxide (NO) is produced in situ: 2NaNO()+6FeSO()+4HSO()3Fe(SO)()+NaSO()+4HO()+2NO()3424243242aqaqaqaqaqlg

The nitrogen oxide produced can then react with the freshly prepared iron(II) sulfate to produce a brown solid containing the complex iron [Fe(HO)(NO)]SO254. A simplified interpretation of this reaction is as follows: FeSO()+NO()FeSONO()44aqgs

The brown ring test described above will not work in the presence of nitrite ions (NO2), and so it is useful to be able to distinguish between nitrite and nitrate ions.

Example 4: Testing for the Presence of Nitrate Ions Using the Brown Ring Test

A student wanted to detect if nitrate ions were present in a solution. They firstly added iron(II) sulfate to the solution before slowly adding concentrated sulfuric acid. They noticed the formation of two layers with a colored ring at the interface. If nitrate ions were present, what color should the ring have been?


In order to test for nitrate anions, a freshly prepared solution of iron(II) sulfate must be used. The iron sulfate solution is then mixed with the solution that may contain nitrate ions.

Concentrated sulfuric acid is carefully added to the inside of the test tube and allowed to run down the inner walls. When the acid reaches the surface of the liquid, it comes into contact with the nitric oxide produced in the reaction between iron(II) sulfate and the nitrate ions. The acid reacts with the nitric oxide, and a brown solid is formed containing a complex ion of Fe3+ and nitric oxide.

The correct answer is, therefore, brown.

Nitrate and nitrite ions can be distinguished by their reactions (or lack of) with acidified potassium permanganate. In the case of nitrite, the acidified potassium permanganate will lose its characteristic deep purple color as it is reduced to manganese(II) ions: 5KNO()+2KMnO()+3HSO()5KNO()+KSO()+2MnSO()+3HO()242432442aqaqaqaqaqaql

Nitrate ions will not reduce permanganate ions, and hence, there will be no color change.

It should be noted however that the results of these tests are not unique to nitrate and nitrite ions and that other anions may produce the same results. As such, positive results for these tests do not definitely confirm the presence of nitrate or nitrite ions.

As we have seen, nitric acid is a good oxidizing agent and is used on an industrial scale in the production of nylon and also as an oxidizer for liquid-fueled rockets. It can also be used as a cleaning agent, in the fertilizer industry, and even in the woodworking community to artificially age pine and maple!

Key Points

  • Nitric acid can be prepared in a laboratory from the reaction of potassium or sodium nitrate and sulfuric acid.
  • Nitric acid is a colorless, strong, and corrosive mineral acid.
  • Nitric acid decomposes when heated to produce NO and is an effective oxidizing agent.
  • Nitric acid reacts in a variety of ways with different metals depending on the activity of the metal and the concentration of the nitric acid.
  • Nitric acid forms metal oxide layers on the surface of some metals in a process known as passivation.
  • The presence of nitrate ions can be partially determined using the brown ring test.
  • The presence of nitrite ions can be partially determined through the decolorization of potassium permanganate.
  • Nitric acid is used as a precursor for many organic chemistry reactions, as an oxidant, and in other areas.

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