General Chemistry/Chemistries of Various Elements/Group 15

The Nitrogen Family

Group 15 (VA) contains nitrogen, phosphorous, arsenic, antimony, and bismuth. Elements in Group 15 have five valence electrons. Because the elements can either gain three electrons or lose five to gain a stable configuration, they more often form covalent compounds unless bonded to an active metal. Their electron affinities are not very large. Metallic properties increase markedly from gaseous nitrogen to barely-metallic bismuth with increasing size and mass. Nitrogen and phosphorus are non-metallic, and arsenic and antimony are metalloids.

Elements of this group are known as pnicogens and their compounds as pniconides. The name is derived from from the Greek word pnicomigs meaning suffocation.

These elements are much less reactive than the elements of Group 16, and their chemistries are more complicated. Most of the chemistry of these elements is in the +3 or +5 oxidation states, although they form gaseous compounds with hydrogen in the -3 oxidation state: ammonia NH₃, phosphine PH₃, arsine AsH₃, stibine AsH₃, and bismuthine BiH₃; these all burn in oxygen to give oxides or the free element (in the case of nitrogen).

$4{\hbox{NH}}_{3}+3{\hbox{O}}_{2}\to 6{\hbox{H}}_{2}{\hbox{O}}+2{\hbox{N}}_{2}$

All form oxides—nitrogen with difficulty, the others with ease. Most of the oxides are acidic, exceptions being nitrous oxide N₂O, nitric oxide NO, and bismuth oxide Bi₂O₃. Except for nitrogen, typical oxides are in the +3 or +5 (bismuth excluded) oxidation states. All form halides—nitrogen with difficulty, but those of phosphorus, arsenic, and antimony fully hydrolyze in water.

Nitrogen and phosphorus form important acids in the +5 oxidation state. Nitrogen forms nitric acid HNO₃, a substance used to create medicines and explosives (but this acid is corrosive and dangerous, so don't touch it or even spill it on anything), and salts known as nitrates such as potassium nitrate KNO₃, an important fertilizer. Phosphorus forms phosphoric acid, H₃PO₄; phosphates are salts of phosphoric acid. Some phosphates are essential to respiration and thus life itself. General covalency of nitrogen family is 3.

Nitrogen

The nitrogen cycle shows how nitrogen is passed along organisms and the atmosphere.

Nitrogen occurs naturally as the diatomic gas N₂. It comprises about 78% of the air we breathe. The bond holding the two nitrogen atoms together is triple covalent, so it is very strong. Because of that, nitrogen is very unreactive. It is used in many places when an inert gas is needed. However, nitrogen will react with some substances:

$6{\hbox{Li}}_{(s)}+{\hbox{N}}_{2(g)}\to 2{\hbox{Li}}_{3}{\hbox{N}}_{(s)}$ $3{\hbox{Mg}}_{(s)}+{\hbox{N}}_{2(g)}\to {\hbox{Mg}}_{3}{\hbox{N}}_{2(s)}$	Although nitrogen gas is usually considered inert, it does react with some elements by burning.
${\hbox{Li}}_{3}{\hbox{N}}_{(s)}+3{\hbox{H}}_{2}{\hbox{O}}\to 3{\hbox{LiOH}}_{(aq)}+{\hbox{NH}}_{3(g)}$ ${\hbox{Mg}}_{3}{\hbox{N}}_{2(s)}+6{\hbox{H}}_{2}{\hbox{O}}\to 3{\hbox{Mg}}({\hbox{OH}})_{2(aq)}+2{\hbox{NH}}_{3(g)}$	The nitrides react violently with water to form ammonia gas and a basic solution.

In its pure form, nitrogen is not very useful, and it will suffocate any animal that breathes pure nitrogen. It is much more important when it is a component of ammonia, nitrate, oxides, or biomolecules like protein. Due it is very unreactive nature, it is difficult to get nitrogen to react and form these useful substances. Any process that can convert elemental nitrogen into a nitrogen compound is called nitrogen fixation. Nitrogen fixation is biologically important because amino acids, proteins, and enzymes contain nitrogen. It is commercially important because it is used in explosives, rocket fuels, and fertilizers.

There are many nitrogen fixation reactions:

${\hbox{N}}_{2}+8{\hbox{H}}^{+}+8{\hbox{e}}^{-}+energy\to 2{\hbox{NH}}_{3}+{\hbox{H}}_{2}$	This occurs in bacterial enzymes. The ammonia (NH3) quickly becomes ammonium (NH4+). The nitrogen in the bacteria enters the soil where plants can absorb it. Humans and animals that eat those plants can get the nitrogen in useful compounds.
${\hbox{N}}_{2}+3{\hbox{H}}_{2}\leftrightarrow 2{\hbox{NH}}_{3}$	The Haber process is used for commercially producing ammonia. This reaction only occurs at very high pressures and temperatures (around 20 MPa and 500 °C) and in the presence of an iron catalyst. Also, the reaction occurs in somewhat complex equipment that must input pure reactants and extract the ammonia.

In the cold, dense, hydrogen-rich atmospheres of Jupiter, Saturn, Uranus, and Neptune, nitrogen ordinarily exists combined with hydrogen as ammonia.

Keep in mind that ammonia is a gas at STP. The household product called "ammonia" is actually an aqueous solution of ammonium hydroxide (NH₄OH) that forms when ammonia gas is dissolved in water. Ammonia, quite unlike hydrogen compounds of Groups 16 and 17, is a base in its reactions, forming salts with weak and strong acids alike. Such a substance as ammonium chloride (NH₄Cl) is a soluble, strongly ionic salt.

Nitrogen compounds are often extremely unstable because nitrogen atoms in nitrogen compounds tend to seek each other to recombine as nitrogen gas. Many nitrogen compounds are literal explosives, including TNT and nitroglycerin. These explosives are in common use in construction projects for the demolition of buildings and other obstacles to new construction, or to get access to minerals in mining operations.

Phosphorus

White phosphorus atomic structure

Phosphorus has two common allotropes: red phosphorus and white phosphorus. White phosphorus (P₄) has a waxy appearance and turns yellow when exposed to light. When exposed to oxygen in the dark, it glows pale green.

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White phosphorus ignites under all but the most delicate conditions. The combustion of white phosphorus produces phosphorus(V) oxide:

${\hbox{P}}_{4}+5{\hbox{O}}_{2}\to {\hbox{P}}_{4}{\hbox{O}}_{10}$

One of its most common uses is in military weapons that cause severe burning of the object hit by the weapon.

Red phosphorus is an amorphous solid. It is more stable and explodes at temperatures higher than those of white phosphorus. It is still, however, dangerously reactive. Both forms of phosphorus are insoluble in water and can be interconverted with various applications of heat, pressure, and light.

There also exist black phosphorus and violet phosphorus. Unlike nitrogen, phosphorus will not readily form a diatomic molecule with a triple bond. Diphosphorus does exist, but only between a temperature range of 1200 °C and 2000 °C.

Phosphorus pentoxide strongly reacts with water to form phosphoric acid, a substance that removes rust from iron, especially on ships; it is often known as "naval jelly". But phosphoric acid is corrosive to flesh and not to be touched.

Phosphorus is essential to life in the form of phosphates in bones and in substances known as ADP and ATP that transform food into useful energy in cells.

Others

A crystal of bismuth, showing its colorful iridescent tarnish.

Arsenic is similar to phosphorus. It has three allotropes: grey arsenic, yellow arsenic, and black arsenic. Grey arsenic is the most common form. Its structure is similar to graphite.

Antimony does not have physical properties of a metal, but behaves chemically as a non-metal.

Bismuth is a brittle, silvery metal. Bismuth is actually radioactive, decaying into thallium-205. Because its half-life is 19 x 10¹⁸ years, about a million times the age of the universe, bismuth is usually considered stable.

Bismuth is much less radioactive than the nearly-harmless and unavoidable radioactive isotopes of carbon and potassium in living things. Unlike arsenic and antimony, its compounds aren't toxic unless something else in the compound is itself toxic or the substance is very acidic or alkaline. In fact, a bismuth compound is very common in a heavily-used stomach medication that requires no prescription.