Video Transcript
In this video, we will learn how to describe and explain the properties and reactions
of the element nitrogen. We will also learn how nitrogen is prepared on a small and large scale. And we’ll look at the reactions of nitrogen-based compounds.
Nitrogen has a wide range of uses, from cryopreservation of biological specimens to
food preservation to being the power source for paintball guns. It has the chemical symbol N. If we find it in the periodic table, we can see that it has an atomic number of
seven. Thus, an atom of nitrogen has seven protons, and a neutral atom will have seven
electrons. It has an atomic mass of around 14 unified atomic mass units. The vast majority of nitrogen atoms have seven neutrons. Nitrogen is in group 15 of the periodic table, so it has five valence or outer shell
electrons.
The electron configuration of a neutral atom is 1s2 2s2 2p3. To satisfy the octet rule, nitrogen needs three more outer shell electrons. If a triple bond is formed between two nitrogen atoms, then this rule is
satisfied. N2, or dinitrogen, is a neutral diatomic gas molecule. It is colorless, odorless, and tasteless. The strong covalent triple bond makes nitrogen gas unreactive. Thus, reactions involving nitrogen in its elemental form often require an electric
spark or strong heating.
Nonpolar N2 molecules don’t interact very well with polar water molecules. So N2 molecules are only slightly soluble in water. At 20 degrees C, only two milligrams of dinitrogen dissolves in 100 grams of
water. If we compare the solubility with diatomic oxygen for reference, we see that
dinitrogen is only slightly less soluble than oxygen. Carbon dioxide has a slightly higher solubility at 0.2 grams per 100 grams of
water. But ammonia has a much higher solubility at 50 grams per 100 grams of water. When nitrogen does dissolve, it forms a neutral solution. At 25 degrees C, it has a pH of approximately seven.
Pure N2 gas has a density of about 1.25 grams per liter at zero degrees C and one
atmosphere. This is slightly below the density of dry air, which is 1.29 grams per liter. The melting point of this small nonpolar molecule is low as expected, at minus 210
degrees C. This is the point at which nitrogen changes from a solid to a liquid and vice
versa. The appearance of liquid nitrogen is similar to that of water. The boiling point of nitrogen is only 14 degrees higher, at minus 196 degrees C. This is the point at which a liquid converts to a gas or vice versa. It’s also important to note that the only forces we see between molecules of nitrogen
are London dispersion forces.
Now, let’s look a little more closely at one of nitrogen’s key properties, which is
its low reactivity. Nitrogen is often referred to as an inert gas. This is generally attributed to the high strength of the nitrogen-nitrogen triple
bond. The triple bond has a bond enthalpy of 942 kilojoules per mole. The triple bond is considerably stronger than the combination of a nitrogen-nitrogen
double bond at 418 kilojoules per mole and a nitrogen-nitrogen single bond at 167
kilojoules per mole.
Since a nitrogen-nitrogen triple bond is stronger than a nitrogen-nitrogen double
bond or single bond, and for that matter a single or double bond between nitrogen
and another element, we can expect nitrogen to form the triple bond over other
combinations. We can compare this with carbon, which we know forms double and single bonds more
frequently than it forms triple bonds. Here, the bond enthalpy for a carbon-carbon double bond and single bond combined is
more than the bond enthalpy for a carbon-carbon triple bond. So just from this data, we would expect carbon to form more double and single
bonds.
Now that we know the properties of nitrogen gas, let’s look at how it’s prepared in
the lab. The easiest way to prepare nitrogen is extraction from air. Nitrogen makes up 78 percent of Earth’s atmosphere, with the other major components
of air being carbon dioxide, water, argon, and oxygen. We can remove the carbon dioxide, water, and oxygen to isolate nitrogen.
We start with a jar full of air. Water is used to displace the air, forcing it into the next section of the
apparatus. The air enters a solution of sodium hydroxide, also known as caustic soda. This is used to remove carbon dioxide. Sodium hydroxide, a base, is used as carbon dioxide is an acidic gas. The reaction equation is two NaOH aqueous plus CO2 gas react to produce Na2CO3 aqueous,
which is sodium carbonate, and H2O, liquid. The remaining gases then pass through concentrated sulfuric acid. This removes water vapor as it’s a dehydrating agent. The remaining gas is then passed through a horizontal tube containing copper turnings
over a source of heat such as a Bunsen burner. This removes the oxygen as it reacts with the copper turnings.
The reaction of copper and oxygen produces copper two oxide. The remaining gas is collected over mercury in a gas cylinder. Mercury is used rather than water as the presence of water would reintroduce water
vapor. The remaining gas is almost pure nitrogen, but it will contain a small amount of
argon. The argon is more difficult to remove, so it will remain. But it only has an abundance of around one percent by moles, and it’s unreactive. So it’s unlikely to interfere with anything we would want to use the nitrogen
for.
If we didn’t want any argon, then there is another method we could use. We can prepare nitrogen in a lab by heating a mixture of ammonium chloride solution
and sodium nitrite solution. Sodium nitrite has the chemical formula NaNO2. Note that this is different from sodium nitrate, which has the chemical formula
NaNO3. The suffix gives an indicator of the number of oxygen atoms. The word ending “-ate” is used for the ion with the larger number of oxygen atoms,
and the word ending “-ite” is used for the ion with fewer oxygen atoms. The reaction of aqueous sodium nitrite with aqueous ammonium chloride produces sodium
chloride and ammonium nitrite.
Ammonium nitrite has been created in situ, as it is highly unstable and will
decompose. When ammonium nitrite decomposes, it produces water and nitrogen gas. The nitrogen gas rises entering the tube and can be collected in a gas cylinder over
water. This method does mean that water vapor will also be collected. If we wanted water-free nitrogen, then we could pass it through concentrated sulfuric
acid and then collect it in a gas cylinder over mercury. So we have shown that the heating of ammonium chloride solution and sodium nitrite
solution produces nitrogen gas. We can represent the fact that heat was required for this reaction by using a capital
Δ. Both of from air and synthesis methods that’s mentioned are used on a reasonably
small scale.
To make tons of nitrogen, a different method is used. In industry, one of the key methods to prepare nitrogen involves fractional
distillation. During this process, air is passed through filters that remove dust and water
vapor. The filtered air is then called to approximately minus 220 degrees C. The CO2 can be removed at this stage as it freezes at minus 79 degrees C. The other major components of air condense at similar temperatures to each other,
making them more difficult to separate. But if we heat the mixture slowly to just above minus 196 degrees C, which is the
temperature at which nitrogen gas condenses, then the nitrogen will turn back into a
gas, leaving the argon and oxygen behind as liquids.
Now that we know how to isolate nitrogen, we’re going to look at the reactions of
nitrogen and its compounds. First, we’re going to look at the reaction of nitrogen with hydrogen. The reaction equation shows us that nitrogen reacts with hydrogen to produce
ammonia. For the reaction to occur, an electric spark and a temperature of 550 degrees C is
needed. In industry, vast quantities of ammonia are produced using the Haber process or
Haber–Bosch process. The products and the reactants are the same as in the previous equation. But the reaction conditions are different.
The Haber process uses catalysts containing iron and molybdenum. A high temperature and pressure are also required to give good yields of product. The Haber process was once used as a source of ammonia for the production of
explosives, but today it is mostly used to produce fertilizer. Nitrogen can also react with oxygen gas. This reaction produces nitric oxide. The reaction of nitrogen and oxygen requires very high temperatures, around 3000
degrees C. These temperatures are easily reached within an electric arc. An electric arc is like an electric spark but tends to last for a longer period of
time.
An example of an electric arc found in nature is lightning. The product, nitric oxide, can then react with oxygen further. The reaction of nitric oxide and oxygen produces nitrogen dioxide. There are lots of different oxides of nitrogen. But nitric oxide and nitrogen dioxide are the two most common.
Now, let’s look at the reaction of nitrogen with metals. Magnesium, a strong reducing agent, will react with nitrogen. This reaction requires heat and produces magnesium nitride, where a nitride ion is
N3−. The magnesium nitride, which is a greenish yellow powder, can then go on to react
with water. This reaction produces ammonia and magnesium hydroxide.
Next, we’re going to look at the reaction of nitrogen with calcium carbide. This reaction produces calcium cyanamide and solid carbon, also referred to as
soot. Calcium cyanamide can then go on to react with water. Upon reaction with water, it produces ammonia and calcium carbonate. Ammonia gas can be absorbed into water in the soil for use by plants.
We have now looked at the key reactions of nitrogen. Let’s have a look at the key reactions of some of its compounds. First, let’s look at the reaction of ammonium chloride and a strong base. We start with the mixture of ammonium chloride and a strong base in water. In this case, the base is calcium hydroxide. But sodium hydroxide is another example of a strong base. Calcium hydroxide in solid form is called slaked lime. And aqueous calcium hydroxide is called lime water and is used to test for the
presence of carbon dioxide. This reaction produces calcium chloride, water, and ammonia.
Ammonia is soluble. But as the vessel is heated, it escapes as a gas. Water vapor will also escape. A drying tube containing calcium oxide can be used to remove the water, leaving just
the ammonia gas. The tube containing the ammonia gas is fed into an upturned gas cylinder. As ammonia gas is less dense than air, it will displace the air, allowing the ammonia
gas to be collected.
We have now seen multiple reactions which produce ammonia gas, but we need to be able
to prove that we’ve made it. We can do this by reacting it with hydrogen chloride. If we add aqueous ammonia to a ball of cotton wool and we add aqueous hydrogen
chloride or hydrochloric acid to another ball of cotton wool, then the ammonia and
hydrogen chloride will evaporate. A reaction occurs between the two colorless gases. Ammonium chloride, a white crystalline solid, is produced. It initially forms as a white vapor. We have now learnt about the reactions of the element nitrogen and nitrogen-based
compounds.
Let’s look over the key points of this video. The nitrogen-nitrogen triple bond is very strong. So nitrogen is commonly found as a diatomic gas molecule. Dinitrogen composes 78 percent of the air and is often described as being inert. It is this strong covalent triple bond that causes dinitrogen to be unreactive. Dinitrogen is colorless, odorless, tasteless, poorly soluble in water, neutral, has a
melting point of minus 210 degrees C and a boiling point of minus 196 degrees C. And dinitrogen gas can be prepared by removing other gases from the air or reacting
sodium nitrite with ammonium chloride.
The nitrogen-based species ammonia can be made using the Haber process, which
involves reaction of nitrogen and hydrogen or by reaction of ammonium chloride and a
strong base, such as calcium hydroxide. We also covered that nitrogen with a lot of energy reacts with hydrogen, oxygen,
magnesium, and calcium carbide. And we looked at the test for ammonia, which involves reaction with hydrogen
chloride, producing ammonium chloride in the form of a white smoke.
