The Nitrogen Group: Physical Properties and the Inertness of Nitrogen

1Surrounded by nitrogen, short of nitrogen

Every breath you take is about 78% nitrogen, and the air above one hectare of a bean field holds tens of thousands of tonnes of it. Yet crops are so often short of nitrogen that farmers buy fertilisers such as urea to supply it. Plants cannot use the nitrogen in the air because the ${\mathrm{N}}_2\backslash mathrm\{N\_2\}$ molecule is one of the most unreactive substances known. Only a few bacteria, such as those in the root nodules of beans, can split it. Phosphorus, directly below nitrogen in Group 15, is the opposite: white phosphorus catches fire in air at about 30 °C and has to be stored under water. This lesson looks at how the physical properties change down Group 15 and why nitrogen is so inert.

2Where they occur

Table 10.1: Sources of the Group 15 elements

3Physical properties down the group

Table 10.2: The Group 15 elements

• Ionisation energy and electronegativity decrease down the group as the atoms grow (Figure 10.1). Nitrogen (3.04) is far more electronegative than the rest of the group; only fluorine, oxygen and chlorine are higher. That is why it forms hydrogen bonds (Lesson 2).
• Metallic character increases: nitrogen and phosphorus are non-metals, arsenic and antimony metalloids, bismuth a (poor) metal.
• Melting points rise from the small molecules ${\mathrm{N}}_2\backslash mathrm\{N\_2\}$ and ${\mathrm{P}}_4\backslash mathrm\{P\_4\}$ to the giant layered structures of arsenic and antimony, then fall for bismuth, whose bonding is more metallic (Figure 10.2).

Figure 10.1: First ionisation energies down Group 15.

Figure 10.2: Melting points down Group 15. Arsenic does not melt at atmospheric pressure; the bar shows its sublimation temperature.

4Why nitrogen is so unreactive

Nitrogen and phosphorus both have three unpaired p electrons, so each atom can form three bonds. They use them in completely different ways (Figure 10.3).

• Nitrogen is small, so its 2p orbitals can overlap sideways with those of a second nitrogen atom. Two N atoms form one σ and two π bonds: the triple bond N≡N, only 110 pm long, with a bond energy of 945 kJ mol⁻¹, among the strongest of all chemical bonds.
• Phosphorus is larger, so its 3p orbitals overlap poorly sideways: the P≡P bond (490 kJ mol⁻¹) is only about half as strong. Each phosphorus atom instead forms three single bonds to three different atoms, giving ${\mathrm{P}}_4\backslash mathrm\{P\_4\}$ tetrahedra (white phosphorus) or long chains (red phosphorus).

Figure 10.3: ${\mathrm{N}}_2\backslash mathrm\{N\_2\}$ and ${\mathrm{P}}_4\backslash mathrm\{P\_4\}$.

Nitrogen's inertness has three causes:

1. The very strong triple bond. Almost every reaction of ${\mathrm{N}}_2\backslash mathrm\{N\_2\}$ must start by breaking or weakening it, so the activation energy is very high and the reactions are extremely slow without a catalyst or very high temperature.
2. No polarity. ${\mathrm{N}}_2\backslash mathrm\{N\_2\}$ is a symmetrical, non-polar molecule, so it does not attract charged or polar reagents.
3. Tightly held electrons. Its high ionisation energy and the large energy gap between its filled and empty orbitals make it a poor electron donor and a poor electron acceptor.

White phosphorus, in contrast, has P–P–P bond angles of only 60°, far smaller than the normal angle for p orbitals, so its bonds are strained and break easily; it ignites in air at about 30 °C. Red phosphorus, a polymer without this strain, does not ignite below about 240 °C.

Using the inertness. Nitrogen gas is pumped into packets of crisps and milk powder and into food stores to keep oxygen out, and into tyres and electrical equipment. Liquid nitrogen (−196 °C) freezes biological samples and is used to remove warts.