nitogen group.pdf

Nitrogen and its Oxides

General properties

• 1st member of group VA

• Colourless, odourless gas

• 78% by volume in air

• Liquid nitrogen used as a coolant

• Most important use is in the manufacture of

ammonia and nitrogenous fertilizers

• Can form a large number of inorganic compounds

• A major constituent of organic compounds such as

Unreactive nature of nitrogen

N N

Strong NN bond,

Bond energy:944 kJ/mol

Reactions involving N₂ have high activation energy and

unfavourable equilibrium constant.

N₂+O₂2NO K_c =4.5x10-31

Nitrogen can be prepared from the air as shown.

Air flows into the aspirator and onto caustic soda

which dissolves carbon dioxide gas.

It is then passed through a heated combustion tube

containing heated copper turnings which remove

oxygen. Nitrogen is then collected over water.

Laboratory Preparation of

Nitrogen

Reactions of nitrogen

• With reactive metals, Li and Mg, to form

nitrides.

– 3Mg(s) + N

(g) 

Mg

N

(s), an ionic compound.

• With oxygen at very high temperature

– N

(g) + O

(g) 

2NO(g) , at very high T

–

2NO(g) + O



2NO2(g)

• With hydrogen at special conditions

Nitrous oxide:

Other methods of preparation

Heating a mixture of sodium nitrate and ammonium sulfate.

2NaNO₃ + (NH₄)₂SO₄ → Na₂SO₄ + 2N₂O+ 4H₂O.

The reaction of urea, nitric acid and sulfuric acid

2 (NH₂)₂CO + 2 HNO₃+ H₂SO₄ → 2 N₂O + 2 CO₂ + (NH₄)₂SO₄ + 2H₂O.

Direct oxidation of ammonia with a manganese dioxide-bismuth oxide catalyst: Ostwald process.

2 NH₃ + 2 O₂ → N₂O + 3 H₂O

Reacting HNO3 with SnCl2 and HCl:

2 HNO₃ + 8 HCl + 4 SnCl₂ → 5 H₂O + 4 SnCl₄ + N₂O

Hyponitrous acid decomposes to N2O and water with a half-life of 16 days at 25 °C at pH 1–3.

Structure

The bond orders have been calculated as N-N 2.73

•

Nitrous oxide is a linear

molecule. It has a boiling point

of -88 ºC, and a melting point

of -102 ºC.

•

It is colourless and has a faintly

sweet smell.

•

It is used as an anesthetic,

NITRIC OXIDE

Nitric oxide, NO, may be prepared by the action of dilute nitric acid on copper:

• In the laboratory, nitric oxide is conveniently generated by reduction of dilute nitric acid with copper:

Preparation

• In commercial settings, ·NO is produced by the oxidation of ammonia at 750–900 °C (normally at 850 °C) with platinum as catalyst:

4 NH₃ + 5 O₂ → 4 ·NO + 6 H₂O

• The uncatalyzed endothermic reaction of O₂ and N₂, which is performed at high temperature (>2000 °C) by lightning has not been developed into a practical commercial synthesis

N₂ + O₂ → 2 ·NO

by the reduction of nitrous acid in the form of sodium nitrite or potassium nitrite:

2 NaNO₂ + 2 NaI + 2 H₂SO₄ → I₂ + 4 NaHSO₄ + 2 ·NO

2 NaNO₂ + 2 FeSO₄ + 3 H₂SO₄ → Fe₂(SO₄)₃ + 2 NaHSO₄ + 2 H₂O + 2 ·NO 3 KNO_2(l) + KNO_3(l) + Cr₂O_3(s) → 2 K₂CrO_4(s) + 4 NO_(g)

Reactions

•When exposed to oxygen, ·NO is converted into nitrogen dioxide.

2 ·NO + O2 → 2 NO2

In water, ·NO reacts with oxygen and water to form

HNO2 or nitrous acid. The reaction is thought to proceed via

the following stoichiometry: 4 ·NO + O2 + 2 H2O → 4 HNO2

•·NO will react with fluorine, chlorine, and bromine to form the NOXspecies, known as the nitrosyl halides, such as nitrosyl

chloride. Nitrosyl iodide can form but is an extremely short-lived species and tends to reform I2.

2 ·NO + Cl2 → 2 NOCl

Dinitrogen trioxide or Nitrogen sesquioxide

Dinitrogen trioxide (N

O

) is formed from the

stoichiometric reaction between NO and O

or NO

and N

O

.

Thermal dissociation of N

O

, occurs above -30

°

C,

and some self-ionization of the pure liquid is

observed.

Structure

Dinitrogen trioxide, however, has N–N bond at 186 pm. The N₂O₃ molecule is planar.

Nitrogen dioxide

• The nitrogen dioxide molecule contains an unpaired electron, which is responsible for its color and paramagnetism.

• It is also responsible for the dimerization of NO₂.

• At low pressures or at high temperatures, nitrogen dioxide has a deep brown color that is due to the presence of the NO₂ molecule.

• At low temperatures, the color almost entirely disappears as dinitrogen tetraoxide, N₂O₄, forms.

Uses

NO₂ is used as an intermediate in the manufacturing of nitric acid, as a nitrating agent in manufacturing of chemical explosives, and as a flour bleaching agent.

Dinitrogen pentoxide

N₂O₅ is a rare example of a compound that adopts two

structures depending on the conditions: most commonly it is a salt, but under some conditions it is a polar molecule: [NO₂+][NO

Synthesis and properties

N

O

was first reported by Deville in 1840, who

prepared it by treating AgNO

with

Cl

A recommended

laboratory synthesis entails dehydrating HNO

with

phosphorus(V) oxide

:

P

O

+ 12 HNO

→ 4 H

PO

+ 6 N

O

In the reverse process, N

O

reacts with water

(

hydrolyses

) to produce nitric acid. Thus, nitrogen

pentoxide is the

anhydride

of nitric acid:

N

O

+ H

O → 2 HNO

N

O

exists as colourless crystals that sublime slightly

above room temperature. The salt eventually

Structure

Solid N₂O₅ is a salt, consisting of separated anions and cations. The cation is the linear nitronium ion NO₂+ and the anion is the planar nitrate NO₃− ion. Thus, the solid could be called nitronium nitrate. Both nitrogen centers have oxidation state +5.

The intact molecule O₂N–O–NO₂ exists in the gas phase (obtained by subliming N₂O₅)

Reactions and applications

Dinitrogen pentoxide, has been used as a reagent to introduce the NO₂functionality. This nitration reaction is represented as follows:

N₂O₅ + Ar–H → HNO₃ + Ar–NO₂

where Ar represents an arene moiety.

N₂O₅ is of interest for the preparation of explosives.

In the atmosphere, dinitrogen pentoxide is an important

Nitrogen oxides as precursors to

smog and acid rain

Nitrogen oxides (NO_x) emissions are estimated to be in the range of 25 - 100 megatonnes of nitrogen per year.

Natural sources are thought to make up approximately 1_/

3 of the

total. The generations of NO_x (primarily a mixture of NO₂ and NO) is the main source of smog and a significant contribution to atmospheric pollution;

Atmospheric reactions leading to acid rain

• The nitrogen dioxide reacts with the hydroxide radical, formed by the photochemical decomposition of ozone, in the presence of a non-reactive gas molecule such as

nitrogen to form nitric acid vapor.

• The conversion rate for NO_x to HNO₃ is approximately ten times faster than the equivalent reaction for sulfur dioxide.

NO

+ h

v

NO

.

+ O

.

.

.

.

At night, conversion takes place via the formation of

a nitrate radical, which subsequently forms

nitrogen pentoxide, that reacts with water on the

surface of aerosol particles to form nitric acid,

.

.

The Haber Process

In the early 1900’s a German chemist called Fritz

Haber came up with his chemical process to make

ammonia using the “free” very unreactive Nitrogen

from the air. (N

is 80% of atmosphere)

This is the reaction:

Nitrogen + Hydrogen Ammonia

Raw Materials

•N

(g) is taken from the air via a process of

fractional distillation.

•H

(g) comes from natural gas, CH

(g)

CH

(g) + H

O (g) 3H

(g) + CO (g)

•The carbon monoxide then reacts with more steam:

Raw Materials cont

Reaction Vessel H2 (g)

from methan

e

N2 (g) from

air

NH2 (g)

The Reaction

•This reaction is exothermic. We increase yield by running

the reaction at low temperatures. However at low

temperatures the reaction rate is incredibly slow.

•Compromise between rate and yield has to be reacted.

The Reaction cont

•The reversible reaction to form ammonia:

N

(g) + 3H

(g) 2NH

(g)

4 moles of gas 2 moles of gas

96 litres (4x24) 48 litres (2x24)

•If pressure is increased in reaction vessel, the reversible

reaction favours ammonia production.

The Reaction cont

•Third condition present within the reaction vessel is an Iron

catalyst.

•The catalyst is a fine mesh designed to maximise surface

area. Iron is a transition metal, and like many transition

metals it makes a good catalyst.

450oc

200 atmospheres Iron catalyst N2 from

air

H2 from

methane and steam

NH3 (g)

After the Reaction Vessel

•Coming out the reaction vessels is NH

(g) and unreacted

N

(g) and H

(g).

•First job is to isolate the NH

(g). This is done by cooling.

The NH

(g) changes state. The nitrogen and hydrogen are

Chemical properties of NH

3

• Weak alkali

• Reaction with acids

• Reaction with metal ions

• As a reducing agent

– Burning in oxygen

– Catalytic oxidation

Nitric(V) Acid

• A very strong acid.

• Turns yellow because of dissolved NO

formed from the decomposition of HNO

.

• Commonly used in making explosives,

The Ostwald Process

The Ostwald process was invented by Wilhelm Ostwald. In

the Ostwald process ammonia is oxidised to form Nitric acid.

Nitric acid is one of the largest user’s of ammonia.

The

process has 3 stages:

Stage 1

•Mixture of air & ammonia heated to

230oC

and is passed

through a metal gauze made of

platinum (90%) & Rhodium

(10%).

•Reaction produces a lot of heat energy..

Stage 1 cont

•Reaction produces nitrogen monoxide (NO) and water.

Ammonia + oxygen Nitrogen monoxide + water

4NH

(g) + 5O

(g) 4NO (g) + 6H

O (g)

Stage 2

Stage 2 cont

Nitrogen monoxide + oxygen Nitrogen dioxide

2NO (g) + O

(g) 2NO

(g)

Stage 3

•The nitrogen dioxide is then dissolved in water to produce

nitric acid.

•Nitrogen dioxide + water Nitric acid + nitrogen

monoxide

Ostwald process

• Catalytic oxidation of NH

– 4NH₃ + 5O₂ (Pt/heat)  4NO + 6H₂O

• Oxidation of NO

– 2NO + O₂  2NO₂

• Dissolving NO

in water and O

– 4NO2 + O2 + 2H2O  4HNO3

Oxidizing properties of HNO

3

• Concentrated HNO

2NO

+ 8H

+ 6e



2NO + 4H

O

• Diluted HNO

2NO

+ 4H

+ 2e



2NO

+ 2H

O

• Reacts with

– Copper

Uses of Nitric acid

Nitric acid produced is used in the manufacture of the

following:

•Artificial fertilisers – Ammonium nitrate.

•Explosives, such as 2,4,6-TNT.

•Dyes.

Nitrates(V)

• Thermal stability

– K,Na 2MNO

2MNO

+ O

Brown ring test for NO

3

-Fresh FeSO₄(aq) and NO₃-(aq) Concentrated H₂SO₄(l)

NO₃- + H₂SO₄  HNO₃ + HSO₄ -HNO₃ + 3Fe2+ _{+ 3H}+  _2H

2O + NO + 3Fe2+

Nitrogen Oxoacids and Oxoanions

• Nitric acid (HNO₃) is produced by the Ostwald process:

– The third step is 3NO₂(g) + H₂O (l) → 2HNO₃ + NO(g)

• Nitric acid is a strong oxidizing agent as well as a strong acid.

• The nitrate (NO₃-_{) also acts as an oxidizing agent.}

– All nitrate salts are water soluble.

• Nitrous acid (HNO₂) is a much weaker acid than nitric acid.

Important oxides of phosphorous.

P₄O₆ has P in its +3 oxidation state.

P₄O₁₀ has P in its +5 oxidation state.

H₃PO₃ has only two acidic H atoms; the third is bonded to the central P and does not dissociate.

The diphosphate ion and polyphosphates.

Oxoacids of Sulphur

Oxoacids are basically the acids that contain oxygen. Sulphur is known to form many oxoacids, for example H2SO4, H2SO3, etc.

In oxoacids of sulphur, sulphur exhibits a tetrahedral structure when coordinated to oxygen.

Generally, oxoacids of sulphur contain at least one S=O bond and one S-OH bond.

Sulphuric acid, H₂SO₄:

• Sulphuric acid is the most popular among the oxoacids of sulphur.

• It is a diprotic acid that is, it ionizes to give two protons.

• In sulphuric acid, one atom of sulphur is bonded to two

hydroxyl groups and the rest two oxygen atoms form pie bonds with the sulphur atom.

• Thus, sulphuric acid exhibits tetrahedral geometry.

• As the bond length of sulphur oxygen bond (S=O) is smaller in comparison to the bond length of S-OH, OH groups are

repelled by the oxygen atoms. Hence, the bond angle of O=S=O bond is larger than the HO-S-OH bond angle.

Sulphurous acid, H₂SO₃:

• Sulphurous acid is a diprotic acid that is, it ionizes two protons.

• In sulphurous acid, one atom of sulphur is bonded to

two hydroxyl groups and one oxygen atom forms a

pie bond with the sulphur atom.

• It is prepared by dissolving sulfur dioxide in water.

• Although, we don’t have any evidence of existence of sulphurous acid in solution phase, but the

Peroxodisulphuric acid, H₂S₂O₈:

Peroxodisulphuric acid contains sulphur in +6 oxidation state. Thus, it is a strong oxidizing agent and highly explosive in

nature.

It is popularly known as Marshall’s acid.

It contains one peroxide group forming a bridge between the two sulphur atoms. Each sulphur atom is connected to one hydroxyl group (S-OH bond) and two oxygen atoms (S=O bond)other than the peroxide group.

It is prepared by the reaction of chlorosulfuric acid with hydrogen peroxide: