Nitrogen

I

Introduction

Nitrogen, symbol N, gaseous element that makes up the largest portion of the Earth's atmosphere. The atomic number of nitrogen is 7. Nitrogen is in group 15 (or Va) of the periodic table.

Nitrogen was isolated by the British doctor Daniel Rutherford in 1772 and recognized as an elemental gas by the French chemist Antoine Laurent Lavoisier about 1776.

II

Properties

Nitrogen is a colourless, odourless, tasteless, non-toxic gas. It can be condensed into a colourless liquid, which can in turn be compressed into a colourless, crystalline solid. Nitrogen exists in two natural isotopic forms, and four radioactive isotopes have been artificially prepared. Nitrogen melts at -210.01° C (-346.02° F), boils at -195.79° C (-320.42° F), and has a density of 1.251 g/litre at 0° C (32° F) and 1 atmosphere pressure. The atomic weight of nitrogen is 14.007.

Nitrogen is obtained from the atmosphere by passing air over heated copper or iron. The oxygen is removed from the air, leaving nitrogen mixed with inert gases. Pure nitrogen is obtained by fractional distillation of liquid air; because liquid nitrogen has a lower boiling point than liquid oxygen, the nitrogen distills off first and can be collected.

Nitrogen composes about four-fifths (78.03 per cent) by volume of the atmosphere. Nitrogen is inert and serves as a diluter for oxygen in burning and respiration processes. It is an important element in plant nutrition; certain bacteria in the soil convert atmospheric nitrogen into a form, such as nitrate, that can be absorbed by plants, a process called nitrogen fixation. Nitrogen in the form of protein is an important constituent of animal tissue. The element occurs in the combined state in minerals, of which saltpetre (KNO₃) and Chile saltpetre (NaNO₃) are commercially important products.

Nitrogen combines with other elements only at very high temperatures or pressures. It is converted to an active form by passing through an electric discharge at low pressure. The nitrogen so produced is very active, combining with alkali metals to form azides; with the vapour of zinc, mercury, cadmium, and arsenic to form nitrides; and with many hydrocarbons to form hydrocyanic acid and cyanides, also known as nitriles. Activated nitrogen returns to ordinary nitrogen in about one minute.

In the combined state nitrogen takes part in many reactions; it forms so many compounds that a systematic scheme of compounds containing nitrogen in place of oxygen was created by the American chemist Edward Franklin. In compounds nitrogen exists in all the valence states between -3 and +5. Ammonia, hydrazine, and hydroxylamine represent compounds in which the valence of nitrogen is -3, -2, and -1, respectively. Oxides of nitrogen represent nitrogen in all the positive valence states.

Burning fossil fuels in cars and power stations, for example, causes some of the nitrogen in air to be converted into nitrogen oxides. These nitrogen oxides (NO_x) react to produce nitrates and nitric acid that cause smog and acid rain, and they also catalyse the formation of harmful ground-level ozone.

III

Uses

Most of the nitrogen used in the chemical industry is obtained by the fractional distillation of liquid air. It is then used to synthesize ammonia. From ammonia produced in this manner, a wide variety of important chemical products are prepared, including fertilizers, nitric acid, urea, hydrazine, and amines. In addition, an ammonia compound is used in the preparation of nitrous oxide (N ₂O) a colourless gas popularly known as laughing gas. Mixed with oxygen, nitrous oxide is used as an anaesthetic for some types of surgery.

Used as a coolant, liquid nitrogen has found widespread application in the field of cryogenics. With the recent advent of ceramic materials that become superconductive at the boiling point of nitrogen, the use of nitrogen as a coolant is increasing (see Superconductivity).