Position of D – Block Elements in the Periodic Table

Find free online Chemistry Topics covering a broad range of concepts from research institutes around the world.

Position of D – Block Elements in the Periodic Table

We have already learnt the periodic classification of elements in XI std. the transition metals occupy from group – 3 to group – 12 of the modern periodic table.

Figure 4.1 – Position of d – block elements in the periodic table

D – Block elements composed of 3d series (4th period) Scandium to Zinc (10 elements), 4d series (5th period) Yttrium to Cadmium (10 elements) and 5d series (6th period) Lanthanum, Haffinium to mercury. As we know that the group-12 elements Zinc, Cadmium and Mercury do not have partially filled d-orbital either in their elemental state or in their normal oxidation states.

Position of D - Block Elements in the Periodic Table img 1

However they are treated as transition elements, because their properties are an extension of the properties of the respective transition elements. As per the IUPAC definition, the seventh period elements, starting from Ac, Rf to Cn also belong to transition metals. All of them are radioactive. Except Actinium; all the remaining elements are synthetically prepared and have very low half life periods.

The d-block elements are found in the middle of the period table. The d-block elements are called transition metals and have valence electrons in d orbital’s. The f-block elements,found in the two rows at the bottom of the periodic table, are called inner transition metals and have valence electrons in the f-orbital’s.

Transition elements are the elements that are found in Groups 3-12 (old groups IIA-IIB) on the periodic table (salmon-colored block in the middle of the table).

The Periodic Table provides a section for each of these groups of orbitals. The 10 electrons of the five d orbitals are filled by the elements found in the dropped central section of the table. This section is referred to as the d block elements, or the transition metals.

The d-block of the periodic table contains the elements of the groups 3-12 in which the d orbitals are progressively filled in each of the four long periods. The f-block consists of elements in which 4 f and 5 f orbitals are progressively filled. They are placed in a separate panel at the bottom of the periodic table.

The d-block elements are called transition elements because they exhibit transitional behaviour between s block and p-block elements. Their properties are transitional between highly reactive metallic elements of s block which are ionic in nature and the elements of p-block which are covalent in nature.

According to Aufbau principle , electrons first occupy the lowest energy orbital available to them and enter into higher energy orbitals only after the lower energy orbitals are filled. Therefore, 3d orbital is higher in energy than 4s. And hence electrons fill up in 4s before filling up in 3d.

The p sublevel has 3 orbitals, so can contain 6 electrons max. The d sublevel has 5 orbitals, so can contain 10 electrons max. And the 4 sublevel has 7 orbitals, so can contain 14 electrons max. In the picture below, the orbitals are represented by the boxes.

D-Block Elements:

Elements in which the last electron enters any one of the five d-orbitals of their respective penultimate shells are called d-block elements. The Importance of d-block Transition Metals. Transition metals, for the most part, are good conductors. They are also malleable, ductile, lustrous, and sliver-white in color. The d-block metals, and some of it’s key alloys, shaped the Bronze Age, Iron Age, and most importantly the steel age.

The p-block elements are found on the right side of the periodic table. They include the boron, carbon, nitrogen, oxygen and flourine families in addition to the noble gases. The noble gases have full p-orbital’s and are nonreactive.

In chemistry and atomic physics, the main group is the group of elements (sometimes called the representative elements) whose lightest members are represented by helium, lithium, beryllium, boron, carbon, nitrogen, oxygen, and fluorine as arranged in the periodic table of the elements.

Chlorine is in group 17 of periodic table, also called the halogens, and is not found as the element in nature – only as a compound. The most common of these being salt, or sodium chloride, and the potassium compounds sylvite (or potassium chloride) and carnallite (potassium magnesium chloride hexahydrate).

Group 18 (Inert gases) Elements

Find free online Chemistry Topics covering a broad range of concepts from research institutes around the world.

Group 18 (Inert gases) Elements

Occurrence:

All the noble gases occur in the atmosphere.

Physical Properties:

As we move along the noble gas elements, their atomic radius and boiling point increases from helium to radon. The first ionization energy decreases from helium to radon. Noble gases have the largest ionisation energy compared to any other elements in a given row as they have completely filled orbital in their outer most shell. They are extremely stable and have a small tendency to gain or lose electrons. The common physical properties of the group 18 elements are listed in the Table.

Physical Properties of group 18 Elements

Group 18 (Inert gases) Elements img 1

Properties of Inert Gases:

Physical Properties:

Noble gases are monoatomic, odourless, colourless, tasteless, and non-inflammable. They are highly unreactive. They are non-metallic in nature.

Chemical Properties:

Only the xenon and krypton show some chemical reactivity. Xenon fluorides are prepared by direct reaction of xenon and fluorine under different conditions as shown below.

Group 18 (Inert gases) Elements img 2

When XeF6 is heated at 50°C in a sealed quartz vessel it forms XeOF4.

Group 18 (Inert gases) Elements img 3

When the reaction is continued the following reaction takes place.

2XeOF4 + SiO2 → 2XeO2F2 + SiF4
2XeO2F2 + SiO2 → 2XeO3 + SiF4

On hydrolysis with water vapour XeF6 gives XeO3

XeF6 + 3H2O → XeO3 + 6HF

When XeF6 reacts with 2.5 M NaOH, sodium per xenate is obtained.

2XeF6 + 16NaOH → Na4XeO6 + Xe + O2 + 12NaF + 8H2O

Sodium per xenate is very much known for its strong oxidizing property. For example, it oxidises manganese (II) ion into permanganate ion even in the absence of the catalyst.

5XeO64- + 2Mn2+ + 14H+ → 2MnO4 + 5XeO3 + 7H2O

Xenon reacts with PtF6 and gave an orange yellow solid [XePtF6] and this is insoluble in CCl4.

Xenon difluoride forms addition compounds XeF2.2SbF5 and XeF2.2TaF5. Xenon
hexa fluorides forms compound with boron and alkali metals. Eg: XeF6.BF3, XeF6MF, M-alkali metals.

There is some evidence for existence of xenon dichloride XeCl2.

Krypton form krypton difluoride when an electric discharge is passed through Kr and flourine at 183°C or when gases are irradiated with SbF5 it forms KrF2.2SbF3.

Structures of Compounds of Xenon:

Compound

Hybridaisation

Shape/Structure

XeFsp3dLinear
XeF4sp3d2Square planar
XeF6sp3d3Distorted octahedron
XeOF2sp3dT Shaped
XeOF4sp3d2Square pyramidal
XeO3sp3Pyramidal

Uses of Noble Gases:

The inertness of noble gases is an important feature of their practical uses.

Helium:

  1. Helium and oxygen mixture is used by divers in place of air oxygen mixture. This prevents the painful dangerous condition called bends.
  2. Helium is used to provide inert atmosphere in electric arc welding of metals.
  3. Helium has lowest boiling point hence used in cryogenics (low temperature science)
  4. It is much less denser than air and hence used for filing air balloons.

Neon:

Neon is used in advertisement as neon sign and the brilliant red glow is caused by passing electric current through neon gas under low pressure.

Argon:

Argon prevents the oxidation of hot filament and prolongs the life in filament bulbs

Krypton:

Krypton is used in florescent bulbs, flash bulbs etc. Lamps filed with krypton are used in airports as approaching lights as they can penetrate through dense fog.

Xenon:

Xenon is used in florescent bulbs, flash bulbs and lasers. Xenon emits an intense light in discharge tubes instantly. Due to this it is used in high speed electronic flash bulbs used by photographers.

Radon:

Radon is radioactive and used as a source of gamma rays. Radon gas is sealed as small capsules and implanted in the body to destroy malignant i.e. cancer growth.

Group 17 (Halogen Group) Elements

Find free online Chemistry Topics covering a broad range of concepts from research institutes around the world.

Group 17 (Halogen Group) Elements

Chlorine

Occurrence:

The halogens are present in combined form as they are highly reactive. The main source of flourine is fluorspar or fluorite. The other ores of flourine are cryolite, fluroapatite. The main source of chlorine is sodium chloride from sea water. Bromides and iodides also occur in sea water.

Physical Properties:

The common physical properties of the group 17 elements are listed in the table.

Physical Properties of Group 17 Elements

Group 17 (Halogen Group) Elements img 1

Properties:

Chlorine is highly reactive hence it doesn’t occur free in nature. It is usually distributed as various metal chlorides. The most important chloride is sodium chloride which occurs in sea water.

Preparation:

Chlorine is prepared by the action of conc. sulphuric acid on chlorides in presence of manganese dioxide.

4NaCl + MnO2 + 4H2SO4 → Cl2 + MnCl2 + 4NaHSO4 + 2H2O

It can also be prepared by oxidising hydrochloric acid using various oxidising agents such as manganese dioxide, lead dioxide, potassium permanganate or dichromate.

PbO2 + 4HCl → PbCl2 + 2H2O + Cl2
MnO2 + 4HCl → MnCl2 + 2H2O + Cl2
2KMnO4 + 16HCl → 2KCl + 2MnCl2 + 8H2O + 5Cl2
K2Cr2O7 + 14HCl → 2KCl + 2CrCl3 + 7H2O + 3Cl2

When bleaching powder is treated with mineral acids chlorine is liberated

CaOCl2 + 2HCl → CaCl2 + H2O + Cl3
CaOCl2 + H2SO4 → CaSO4 + H2O + Cl2

Manufacture of Chlorine:

Chlorine is manufactured by the electrolysis of brine in electrolytic process or by oxidation of HCl by air in Deacon’s process.

Electrolytic Process:

When a solution of brine (NaCl) is electrolysed, Na+ and Cl ions are formed. Na+ ion reacts with
OH ions of water and forms sodium hydroxide. Hydrogen and chlorine are liberated as gases.

Group 17 (Halogen Group) Elements img 2

Deacon’s Process:

In this process a mixture of air and hydrochloric acid is passed up a chamber containing a number of shelves, pumice stones soaked in cuprous chloride are placed. Hot gases at about 723 K are passed through a jacket that surrounds the chamber.

Group 17 (Halogen Group) Elements img 3

The chlorine obtained by this method is dilute and is employed for the manufacture of bleaching powder. The catalysed reaction is given below,

Group 17 (Halogen Group) Elements img 4

Physical Properties:

Chlorine is a greenish yellow gas with a pungent irritating odour. It produces headache when inhaled even in small quantities whereas inhalation of large quantities could be fatal. It is 2.5 times heavier than air. Chlorine is soluble in water and its solution is referred as chlorine water. It deposits greenish yellow crystals of chlorine hydrate (Cl2.8H2O). It can be converted into liquid (Boiling point – 34.6°C) and yellow crystalline solid (Melting point – 102°C).

Chemical Properties:

Action with metals and non-metals:

It reacts with metals and non metals to give the corresponding chlorides.

2Na + Cl2 → 2NaCl
2Fe + 3Cl2 → 2FeCl3
2Al + 3Cl2 → 2AlCl3
Cu + Cl2 → CuCl2
H2 + Cl2 → 2HCl; ∆H = – 44kCal
2B + 3Cl2 → 2BCl3

Group 17 (Halogen Group) Elements img 5

P4 + 6Cl2 → 4PCl3
2As + 3Cl2 → 2AsCl3
2Sb + 3Cl2 → 2SbCl3

Affinity for Hydrogen:

When burnt with turpentine it forms carbon and hydrochloric acid.

C10H16 + 8Cl2 → 10C + 16HCl

It forms dioxygen when reacting with water in presence of sunlight. When chlorine in water is exposed to sunlight it loses its colour and smell as the chlorine is converted into hydrochloric acid.

2Cl2 + 2H2O → O2 + 4HCl

Chlorine reacts with ammonia to give ammonium chloride and other products as shown below:
With excess ammonia,

2NH3 + 3Cl2 → N2 + 6HCl
6HCl + 6H3 → 6NH4Cl

Overall Reaction
8NH3 + 3Cl2 → N2 + 6NH4Cl

With excess chlorine,
NH3 + 3Cl2 → NCl3 + 3HCl
3HCl + 3NH3 → 3NH4Cl

Overall Reaction
4NH3 + 3Cl2 → NCl3 + 3HCl
3HCl + 3HCl3 → 3NH4Cl

Chlorine oxidises hydrogen sulphide to sulphur and liberates bromine and iodine from iodides and bromides. However, it doesn’t oxidise fluorides.

H2S + Cl2 → 2HCl + S
Cl2 + 2KBr → 2KCl + Br2
Cl2 + 2KI → 2KCl + I2

Reaction with Alkali:

Chlorine reacts with cold dilute alkali to give chloride and hypochlorite while with hot concentrated alkali chlorides and chlorates are formed.

Cl2 + H2O → HCl + HOCl
HCl + NaOH → NaCl + H2O
HOCl + NaOH → NaOCl + H2O

Group 17 (Halogen Group) Elements img 6

(Cl2 + H2O → HCl + HOCl) × 3
(HCl + NaOH → NaCl + H2O) × 3
(HOCl + NaOH → NaOCl + H2O) × 3
3NaOCl → NaClO3 + 2NaCl

Group 17 (Halogen Group) Elements img 7

Oxidising and Bleaching Action:

Chlorine is a strong oxidising and bleaching agent because of the nascent oxygen.

Group 17 (Halogen Group) Elements img 8

Colouring matter + Nascent oxygen → Colourless oxidation product

The bleaching of chlorine is permanent. It oxidises ferrous salts to ferric, sulphites to sulphates and hydrogen sulphide to sulphur.

2FeCl2 + Cl2 → 2FeCl3
Cl2 + H2O → HCl + HOCl
2FeSO4 + H2SO4 + HOCl → Fe2(SO4)3 + HCl + H2O

Group 17 (Halogen Group) Elements img 9

Cl2 + H2O → HCl + HOCl
Na2SO3 + HOCl → Na2SO4 + HCl

Overall Reaction
Na2SO3 + H2O + Cl2 → CaOCl2 + H2O

Preparation of Bleaching Powder:

Bleaching powder is produced by passing chlorine gas through dry slaked lime (calcium hydroxide).
Ca(OH)2 + Cl2 → CaOCl2 + H2O

Displacement Redox Reactions:

Chlorine displaces bromine from bromides and iodine from iodide salts.

Cl2 + 2KBr → 2KCl + Br2
Cl2 + 2KI → 2KCl + I2

Formation of Addition Compounds:

Chlorine forms addition products with sulphur dioxide, carbon monoixde and ethylene. It forms substituted products with alkanes/arenes.

Group 17 (Halogen Group) Elements img 10

Uses of Chlorine:

It is used in

  1. Purification of drinking water
  2. Bleaching of cotton textiles, paper and rayon
  3. Extraction of gold and platinum

Hydrochloric Acid:

Laboratory Preparation:

It is prepared by the action of sodium chloride and concentrated sulphuric acid.

NaCl + H2SO4 → NaHSO4 + HCl
NaHSO4 + NaCl → Na2SO4 + HCl

Dry hydrochloric acid is obtained by passing the gas through conc. sulphuric acid.

Properties:

Hydrogen chloride is a colourless, pungent smelling gas, easily liquefied to a colourless liquid (boiling point 189K) and frozen into a white crystalline solid (melting point 159K). It is extremely soluble in water.

HCl (g) + H2O (l) → H3O+ + Cl

Chemical Properties:

Like all acids it liberates hydrogen gas from metals and carbon dioxide from carbonate and bicarbonate salts.

Zn + 2HCl → ZnCl2 + H2
Mg + 2HCl → MgCl2 + H2
Na2CO3 + 2HCl → 2NaCl + CO2 + H2
CaCO3 + 2HCl → CaCl2 + CO2 + H2
NaHCO3 + 2HCl → 2NaCl + CO2 + H2O

It liberates sulphur dioxide from sodium sulphite

Na2SO3 + 2HCl → 2NaCl + H2O + SO2

When three parts of concentrated hydrochloric acid and one part of concentrated nitric acid are mixed, Aquaregia (Royal water) is obtained. This is used for dissolving gold, platinum etc.

Au + 4H+ + NO3 + 4Cl → AuCl4 + NO + 2H2O
3Pt + 16H+ + 4NO3 + 18Cl → 3[PtCl6]2- + 4NO + 8H2O

Uses of Hydrochloric Acid:

  1. Hydrochloric acid is used for the manufacture of chlorine, ammonium chloride, glucose from corn starch etc.,
  2. It is used in the extraction of glue from bone and also for purification of bone black

Trends in Physical and Chemical Properties of Hydrogen Halides:

Preparation:

Direct combination is a useful means of preparing hydrogen chloride. The reaction between hydrogen and flourine is violent while the reaction between hydrogen and bromine or hydrogen and iodine are reversible and don’t produce pure forms.

Displacement Reactions:

Concentrated sulphuric acid displaces hydrogen chloride from ionic chlorides. At higher temperatures the hydrogen sulphate formed react with further ionic chloride. Displacement can be used for the preparation of hydrogen flourides from ionic flourides. Hydrogen bromide and hydrogen iodide are oxidised by concentrated sulphuric acid and can’t be prepared in this method.

Hydrolysis of Phosphorus Trihalides:

Gaseous hydrogen halides are produced when water is added in drops to phosphorus tri halides except phosphorus triflouride.

PX3 + 3H2O → H3PO3 + 3HX

Hydrogen bromide may be obtained by adding bromine dropwise to a paste of red phosphorous and water while hydrogen iodide is conveniently produced by adding water dropwise to a mixture of red phosphorous and iodine.

2P + 3X2 → 2PX3
2PX3 + 3H2O → H3PO3 + 3HX (where X = Br or I)

Any halogen vapours which escapes with the hydrogen halide is removed by passing the gases through a column of moist red phosphorous.

From Covalent Hydrides:

Halogens are reduced to hydrogen halides by hydrogen sulphide.

H2S + X2 → 2HX + S

Hydrogen chloride is obtained as a by-product of the reactions between hydrocarbon of halogens.

General Properties:

In line with the decreasing bond dissociation enthalpy, the thermal stability of hydrogen halides decreases from flouride to iodide. For example, Hydrogen iodide decomposes at 400° C while hydrogen flouride and hydrogen chloride are stable at this temperature.

At room temperature, hydrogen halides are gases but hydrogen flouride can be readily liquefied. The gases are colourless but, with moist air gives white fumes due to the production of droplets of hydrohalic acid. In HF, due to the presence of strong hydrogen bond it has high melting and boiling points. This effect is absent in other hydrogen halides.

HF

HClHBr

HI

Bond dissociation enthalphy (kJ mol-1)+562+431+366+299
% of ionic character4317137

Acidic Properties:

The hydrogen halides are extremely soluble in water due to the ionisation.

HX + H2O → H3O+ + X
(X – F, Cl, Br, or I)

Solutions of hydrogen halides are therefore acidic and known as hydrohalic acids. Hydrochloric, hydrobromic and hydroiodic acids are almost completely ionised and are therefore strong acids but HF is a weak acid i.e. 0.1mM solution is only 10% ionised, but in 5M and 15M solution HF is stronger acid due to the equilibrium.

HF + H2O ⇄ H3O+ + F
HF + F ⇄ HF2

At high concentration, the equilibrium involves the removal of flouride ions is important. Since it affects the dissociation of hydrogen flouride and increases and hydrogen ion concentration Several stable salts NaHF2, KHF2 and NH4HF2 are known. The other hydrogen halides do not form hydrogen dihalides.

Hydrohalic acid shows typical acidic properties. They form salts with acids, bases and reacts with metals to give hydrogen. Moist hydroflouric acid (not dry) rapidly react with silica and glass.

SiO2 + 4HF → SiF4 + 2H2O
Na2SiO3 + 6HF → Na2SiF6 + 3H2O

Oxidation:

Hydrogen iodide is readily oxidised to iodine hence it is a reducing agent.

2HI ⇄ 2H+ + I2 + 2e

Acidic solution of iodides is readily oxidised. A positive result is shown by liberation of iodine which gives a blue-black colouration with starch.

Hydrogen bromide is more difficult to oxidise than HI. HBr reduces slowly H2SO4 into SO2

2HBr + H2SO4 → 2H2O + Br2 + SO2

But hydrogen iodide and ionic iodides are rapidly reduced by H2SO4 into H2S and not into SO2.

8HI + H2SO4 → 4H2O + 4I2 + H2S

Reducing property of hydrogen iodide can be also explained by using its reaction with alcohols into ethane. It converts nitric acid into nitrous acid and dinitrogen dioxide into ammonium.

Hydrogen chloride is unaffected by concentrated sulphuric acid but affected by only strong oxidising agents like MnO2, potassium permanganate or potassium chloride.

To summarize the trend,

Property

Order

Reactivity of hydrogenDecreases from fluorine to iodine
StabilityDecreases from HF to HI
Volatility of the hydridesHF < HI < HBr < HCl
Thermal StabilityHF > HI > HBr > HCl
Boiling PointHCl < HBr < HI
Acid strengthIncreases from HF to HI

Inter Halogen Compounds:

Each halogen combines with other halogens to form a series of compounds called inter halogen compounds. In the given table of inter halogen compounds a given compound A is less electronegative than B.

Group 17 (Halogen Group) Elements img 11

Properties of Inter Halogen Compounds:

  1. The central atom will be the larger one
  2. It can be formed only between two halogen and not more than two halogens
  3. Fluorine can’t act as a central metal atom being the smallest one
  4. Due to high electronegativity with small size florine helps the central atom to attain high coordination number
  5. They can undergo the auto ionization.
  6. They are strong oxidising agents

2 ICI ⇄ I+ + ICI2
2 ICI ⇄ ICI+2 + ICI4

Reaction with Alkali:

When heated with the alkalis, larger halogen form oxyhalogens and the smaller forms halide.

Group 17 (Halogen Group) Elements img 12

Structure of Inter Halogen Compounds:

The structures of different type of interhalogen compunds can be easily explained using VSEPR theory. The details are given below.

Group 17 (Halogen Group) Elements img 13

Oxides of Halogen

Fluorine reacts readily with oxygen and forms difluorine oxide (F2O) and diflourine dioxide (F2O2) where
it has a – 1 oxidation state. Other halogens do not react with oxygen readily. But the following oxides can be prepared by some indirect methods. Except flourine all the other halogens have positive oxidation states.

Group 17 (Halogen Group) Elements img 14

Oxoacids of Halogens:

Chlorine forms four types of oxoacids namely hypochlorus acid, chlorous acid, chloric acid and perchloric acid. Bromine and iodine forms the similar acids except halous acid. However, flourine only forms hypofulric acid. The oxidizing power oxo acids follows the order:

HOX > HXO3 > HXO3 > HXO4

Group 17 (Halogen Group) Elements img 15

Oxygen Preparation and its Properties

Find free online Chemistry Topics covering a broad range of concepts from research institutes around the world.

Oxygen Preparation and its Properties

Preparation

The atmosphere and water contain 23% and 83% by mass of oxygen respectively. Most of the world’s rock contain combined oxygen. Industrially oxygen is obtained by fractional distillation of liquefied air. In the laboratory, oxygen is prepared by one of the following methods.

The decomposition of hydrogen peroxide in the presence of catalyst (MnO2) or by oxidation with potassium permanganate.

2H2O2 ⇄ 2H2O + O2
5H2O2 + 2MnO4 + 6H+ → 5O2 + 8H2O + 2Mn2+

The thermal decomposition of certain metallic oxides or oxoanions gives oxygen.

Oxygen Preparation and its Properties img 1

Properties

Under ordinary condition oxygen exists as a diatomic gas. Oxygen is paramagnetic. Like nitrogen and flourine, oxygen form strong hydrogen bonds. Oxygen exists in two allotropic forms namely dioxygen (O2) and ozone or trioxygen (O3).

Although negligible amounts of ozone occurs at sea level it is formed in the upper atmosphere by the action of ultraviolet light. In the laboratory ozone is prepared by passing electrical discharge through oxygen. At a potential of 20,000 V about 10% of oxygen is converted into ozone it gives a mixture known as ozonised oxygen. Pure ozone is obtained as a pale blue gas by the fractional distillation of liquefied ozonised oxygen.

Oxygen Preparation and its Properties img 2

Oxygen Preparation and its Properties img 3

The ozone molecule has a bent shape and symmetrical with delocalised bonding between the oxygen atoms.

Chemical Properties

The chemical properties of oxygen and ozone differ vastly. Oxygen combines with many metals and non-metals to form oxides. With some elements such as s-block elements combination of oxygen occurs at room temperature.

Some of less reactive metals react when powdered finely and made to react exothermically with oxygen at room temperature but a lump of metal is unaffected under same condition. These finely divided metals are known as pyrophoric and when set the powder on fie, heat is liberated during a reaction.

On the other hand ozone is a powerful oxidising agent and it reacts with many substances under conditions where oxygen will not react. For example, it oxidises potassium iodide to iodine. This reaction is quantitative and can be used for estimation of ozone.

O3 + 2KI + H2O → 2KOH + O2I2

Ozone is commonly used for oxidation of organic compounds. In acidic solution ozone exceeds the oxidising power of flourine and atomic oxygen. The rate of decomposition of ozone drops sharply in alkaline solution.

Uses:

  1. Oxygen is one of the essential component for the survival of living organisms.
  2. It is used in welding (oxyacetylene welding)
  3. Liquid oxygen is used as fuel in rockets etc

Allotrophic forms of Sulphur

Sulphur exists in crystalline as well as amorphous allotrophic forms. The crystalline form includes rhombic sulphur (α sulphur) and monoclinic sulphur (β sulphur). Amorphous allotropic form includes plastic sulphur (γ sulphur), milk of sulphur and colloidal sulphur.

Rhombic sulphur also known as α sulphur, is the only thermodynamically stable allotropic form at ordinary temperature and pressure. The crystals have a characteristic yellow colour and composed of S8 molecules. When heated slowly above 96°C, it converts into monoclinic sulphur.

Upon cooling below 96°C the β form converts back to α form. Monoclinic sulphur also contains S8 molecules in addition to small amount of S6 molecules. It exists as a long needle like prism and is also called as prismatic sulphur. It is stable between 96°C – 119°C and slowly changes into rhombic sulphur.

When molten sulphur is poured into cold water a yellow rubbery ribbon of plastic sulphur is produced. They are very soft and can be stretched easily. On standing (cooling slowly) it slowly becomes hard and changes to stable rhombic sulphur.

Sulphur also exists in liquid and gaseous states. At around 140°C the monoclinic sulphur melts to form mobile pale yellow liquid called λ sulphur. The vapour over the liquid sulphur consists of 90% of S8, S7 & S6 and small amount of mixture of S2, S3, S4, S5 molecules.

Sulphur Dioxide

Preparation

From Sulphur

A large-scale production of sulphur dioxide is done by burning sulphur in air. About 6-8% of sulphur is oxidised to SO3.

S + O2 → SO2
2S + 3O2 → 2SO3

From Sulphides

When sulphide ores such as galena (PbS), zinc blende (ZnS) are roasted in air, sulphur dioxide is liberated. Large amounts of sulphur dioxide required for manufacturing of sulphuric acid and other industrial purpose is prepared by this method.

Oxygen Preparation and its Properties img 4

Laboratory Preparation

Sulphur dioxide is prepared in the laboratory treating a metal or metal sulphite with sulphuric acid

Cu + 2H2SO4 → CuSO4 + SO2 + 2H2O
SO32- + 2H+ → H2O + SO2

Properties

Sulphur dioxide gas is found in volcanic eruptions. A large amount of sulphur dioxide gas is released into atmosphere from power plants using coal and oil and copper melting plants. It is a colourless gas with a suffocating odour. It is highly soluble in water and it is 2.2 times heavier than air. Sulphur dioxide can be liquefied (boiling point 263 K) at 2.5 atmospheric pressure and 288 K.

Chemical Properties

Sulphur dioxide is an acidic oxide. It dissolves in water to give sulphurous acid.

Oxygen Preparation and its Properties img 5

Reaction with Sodium Hydroxide and Sodium Carbonate

Sulphur dioxide reacts with sodium hydroxide and sodium carbonate to form sodium bisulphite and sodium sulphite respectively.

Oxygen Preparation and its Properties img 6

Oxidising Property

Sulphur dioxide, oxidises hydrogen sulphide to sulphur and magnesium to magnesium oxide.

2H2S + SO2 → 3S + 2H2O
2Mg + SO2 → 2MgO + S

Reducing Property

As it can readily be oxidised, it acts as a reducing agent. It reduces chlorine into hydrochloric acid.
SO2 + 2H2O + Cl2 → H2SO4 + 2HCl

It also reduces potassium permanganate and dichromate to Mn2+ and Cr3+ respectively.

2KMnO4 + 5SO2 + 2H2O → K2SO4 + 2MnSO4 + 2H2SO4
K2Cr2O7 + 3SO2 + H2SO4 → K2SO4 + Cr2(SO4)3 + H2O

Reaction with Oxygen

Sulphur dioxide is oxidised to sulphur trioxide upon heating with oxygen at high temperature. This reaction is used for the manufacture of sulphuric acid by contact process.

Oxygen Preparation and its Properties img 7

Bleaching Action of Sulphur Dioxide

In presence of water, sulphur dioxide bleaches coloured wool, silk, sponges and straw into colourless due to its reducing property.

Oxygen Preparation and its Properties img 8

However, the bleached product (colourless) is allowed to stand in air, it is reoxidised by atmospheric oxygen to its original colour. Hence bleaching action of sulphur dioxide is temporary.

Uses:

  1. Sulphur dioxide is used in bleaching hair, silk, wool etc.
  2. It can be used for disinfecting crops and plants in agriculture.

Structure of Sulphur Dioxide

In sulphur dioxide, sulphur atom undergoes sp2 hybridisation. A double bond arises between S and O is due to pπ – dπ overlapping.

Oxygen Preparation and its Properties img 9

Sulphuric Acid: (H2SO4)

Preparation:

Sulphuric acid can be manufactured by lead chamber process, cascade process or contact process. Here we discuss the contact process.

Manufacture of Sulphuric Acid by Contact Process:

The contact process involves the following steps.

(i) Initially sulphur dioxide is produced by burning sulphur or iron pyrites in oxygen/air.

S + O2 → SO2
4FeS2 + 11O2 → 2Fe2O3 + 8SO2

(ii) Sulphur dioxide formed is oxidised to sulphur trioxide by air in the presence of a catalyst such as V2O5 or
platinised asbestos.

(iii) The sulphur trioxide is absorbed in concentrated sulphuric acid and produces oleum (H2S2O7). The oleum is converted into sulphuric acid by diluting it with water.

Oxygen Preparation and its Properties img 10

To maximise the yield the plant is operated at 2 bar pressure and 720 K. The sulphuric acid obtained in this process is over 96% pure.

Physical Properties:

Pure sulphuric acid is a colourless, viscous liquid (Density: 1.84 g/mL at 298 K). High boiling point and viscosity of sulphuric acid is due to the association of molecules together through hydrogen bonding.

The acid freezes at 283.4 K and boils at 590 K. It is highly soluble in water and has strong affinity towards water and hence it can be used as a dehydrating agent. When dissolved in water, it forms mono (H2SO4.H2O) and dihydrates (H2SO4.2H2O) and the reaction is exothermic.

The dehydrating property can also be illustrated by its reaction with organic compounds such as sugar, oxalic acid and formic acid.

Oxygen Preparation and its Properties img 11

Chemical Properties:

Sulphuric acid is highly reactive. It can act as strong acid and an oxidising agent.

Decomposition:

Sulphuric acid is stable, however, it decomposes at high temperatures to sulphur trioxide.

H2SO4 → H2O + SO3

Acidic Nature:

It is a strong dibasic acid. Hence it forms two types of salts namely sulphates and bisulphates.

Oxygen Preparation and its Properties img 12

Oxidising Property:

Sulphuric acid is an oxidising agent as it produces nascent oxygen as shown below.

Oxygen Preparation and its Properties img 13

Sulphuric acid oxidises elements such as carbon, sulphur and phosphorus. It also oxidises bromide and iodide to bromine and iodine respectively.

C + 2H2SO4 → 2SO2 + 2H2O + CO2
S + 2H2SO4 → 3SO2 + 2H2O
P4 + 10H2SO4 → 4H3PO4 + 10SO2 + 4H2O
H2S + H2SO4 → SO2 + 2H2O + S
H2SO4 + 2HI → SO2 + 2H2O + I2
H2SO4 + 2HBr → SO2 + 2H2O + Br2

Reaction with Metals:

Sulphuric acid reacts with metals and gives different product depending on the reactants and reacting condition.

Dilute sulphuric acid reacts with metals like tin, aluminium, zinc to give corresponding sulphates.

Zn + H2SO4 → ZnSO4 + H2
2Al + 3H2SO4 → Al2(SO4)3 + 3H2

Hot concentrated sulphuric acid reacts with copper and lead to give the respective sulphates as shown below.

Cu + 2H2SO4 → CuSO4 + 2H2O + SO2
Pb + 2H2SO4 → PbSO4 + 2H2O + SO2

Sulphuric acid doesn’t react with noble metals like gold, silver and platinum.

Reaction with Salts:

It reacts with different metal salts to give metal sulphates and bisulphates.

KCl + H2SO4 → KHSO4 + HCl
KNO3 + H2SO4 → KHSO4 + HNO3
Na2CO3 + H2SO4 → Na2SO4 + H2O + CO2
2NaBr + 3H2SO4 → 2NaHSO4 + 2H2O + Br2 + SO2

Reaction with Organic Compounds:

It reacts organic compounds such as benzene to give sulphonic acids.

Oxygen Preparation and its Properties img 14

Uses of Sulphuric Acid:

  1. Sulphuric acid is used in the manufacture of fertilisers, ammonium sulphate and super phosphates and other chemicals such as hydrochloric acid, nitric acid etc.
  2. It is used as a drying agent and also used in the preparation of pigments, explosives etc.

Test for Sulphate/Sulphuric Acid:

Dilute solution of sulphuric acid/aqueous solution of sulphates gives white precipitate (barium sulphate) with barium chloride solution. It can also be detected using lead acetate solution. Here a white precipitate of lead sulphate is obtained.

Oxygen Preparation and its Properties img 15

Structure of Oxoacids of Sulphur:

Sulphur forms many oxoacids. The most important one is sulphuric acid. Some acids like sulphurous and dithionic acids are known in the form of their salts only since the free acids are unstable and cannot be isolated. Various oxo acids of sulphur with their structures are given below.

Oxygen Preparation and its Properties img 16

Oxygen Preparation and its Properties img 17

Group 15 (Nitrogen Group) Elements

Find free online Chemistry Topics covering a broad range of concepts from research institutes around the world.

Group 15 (Nitrogen Group) Elements

Occurrence:

About 78% of earth atmosphere contains dinitorgen (N2) gas. It is also present in earth crust as sodium nitrate (Chile saltpetre) and potassium nitrate (Indian saltpetre). The 11th most abundant element phosphorus, exists as phosphate (fluroapatite, chloroapatite and hydroxyapatite). The other elements arsenic, antimony and bismuth are present as sulphides and are not very abundant.

Physical Properties:

Some of the physical properties of the group 15 elements are listed below

Table 3.1 Physical Properties of Group 15 Elements

Group 15 (Nitrogen Group) Elements img 1

Nitrogen:

Preparation:

Nitrogen, the principal gas of atmosphere (78% by volume) is separated industrially from liquid air by fractional distillation

Pure nitrogen gas can be obtained by the thermal decomposition of sodium azide about 575 K

2NaN3 → 2 Na + 3Na2

It can also be obtained by oxidising ammonia using bromine water

8NH3 + 3Br2 → 6NH4Br + N2

Properties

Nitrogen gas is rather inert. Terrestrial nitrogen contains 14.5% and 0.4% of nitrogen-14 and nitrogen-15 respectively. The later is used for isotopic labelling. The chemically inert character of nitrogen is largely due to high bonding energy of the molecules 225 cal mol-1 (946 kJ mol-1).

Interestingly the triply bonded species is notable for its less reactivity in comparison with other iso-electronic triply bonded systems such as -C≡C-, C≡O, X-C≡N, X-N≡C, -C≡C-, and -C≡N. These groups can act as donor where as dinitrogen cannot. However, it can form complexes with metal (M ← N≡N) like CO to a less extent.

The only reaction of nitrogen at room temperature is with lithium forming Li3N. With other elements, nitrogen combines only at elevated temperatures. Group 2 metals and T forms ionic nitrides.

6Li + N2 → 2Li3N

Group 15 (Nitrogen Group) Elements img 2

Direct reaction with hydrogen gives ammonia. This reaction is favoured by high pressures and at optimum temperature in presence of iron catalyst. This reaction is the basis of Haber’s process for the synthesis of ammonia.

\(\frac{1}{2}\)N2 + \(\frac{3}{2}\)H2 ⇄ NH3
∆Hf = -46.2 kJ mol-1

With oxygen, nitrogen produces nitrous oxide at high temperatures. Even at 3473 K nitrous oxide yield is only 4.4%.

2N2 + O2 → 2N2O

Uses of Nitrogen

  1. Nitrogen is used for the manufacture of ammonia, nitric acid and calcium cyanamide etc.
  2. Liquid nitrogen is used for producing low temperature required in cryosurgery, and so in biological preservation.

Ammonia (NH3)

Preparation:

Ammonia is formed by the hydrolysis of urea.
NH2CONH2 + H2O → 2NH3 + CO2

Ammonia is prepared in the laboratory by heating an ammonium salt with a base.

2NH4+ + OH → 2NH3 + H2O
2NH4Cl + CaO → CaCl2 + 2NH3 + H2O

It can also be prepared by heating a metal nitrides such as magnesium nitride with water.

Mg3N2 + 6H2O → 3Mg(OH)2 + 2NH3

It is industrially manufactured by passing nitrogen and hydrogen over iron catalyst a small amount of K2O and Al2O3 is also used to increase the rate of attainment of equilibrium) at 750 K at 200 atm pressure. In the actual process the hydrogen required is obtained from water gas and nitrogen from fractional distillation of liquid air.

Properties

Ammonia is a pungent smelling gas and is lighter than air. It can readily liquefied by at about 9 atmospheric pressure. The liquid boils at – 38.4°C and freezes at – 77° C. Liquid ammonia resembles water in its physical properties. i.e. it is highly associated through strong hydrogen bonding. Ammonia is extremely soluble in water (702 Volume in 1 Volume of water) at 20°C and 760mm pressure.

At low temperatures two soluble hydrate NH3.H2O and 2NH3.H2O are isolated. In these molecules ammonia and water are linked by hydrogen bonds. In aqueous solutions also ammonia may be hydrated in a similar manner and we call the same as (NH3.H2O)

NH3 + H2O ⇄ NH4+ + OH

The dielectric constant of ammonia is considerably high to make it a fairly good ionising solvent like water.

Group 15 (Nitrogen Group) Elements img 3

Chemical Properties

Action of Heat:

Above 500°C ammonia decomposes into its elements. The decomposition may be accelerated by metallic catalysts like Nickel, Iron. Almost complete dissociation occurs on continuous sparking.

Group 15 (Nitrogen Group) Elements img 4

Reaction with Air/Oxygen:

Ammonia does not burn in air but burns freely in free oxygen with a yellowish flame to give nitrogen and steam.

4NH3 + 3O2 ⇄ N2 + 6H2O

In presence of catalyst like platinum, it burns to produce nitric oxide. This process is used for the manufacture of nitric acid and is known as ostwalds process.

4NH3 + 5O2 ⇄ 4NO + 6H2O

Reducing Property:

Ammonia acts as a reducing agent. It reduces the metal oxides to metal when passed over heated metallic oxide.

3PbO + 2NH3 → 3Pb + N2 + 3H2O

Reaction with Acids:

When treated with acids it forms ammonium salts. This reaction shows that the affinity of ammonia for proton is greater than that of water.

Reaction with Chlorine and Chlorides:

Ammonia reacts with chlorine and chlorides to give ammonium chloride as a final product. The reactions are different under different conditions as given below.

With excess ammonia

2 NH3 + 3 Cl2 → N2 + 6 HCl
6 HCl + 6 NH3 → 6 NH4Cl

Group 15 (Nitrogen Group) Elements img 5

With excess of chlorine ammonia reacts to give nitrogen trichloride, an explosive substance.

2NH3 + 6Cl2 → 2NCl3 + 6HCl
2NH3(g) + HCl(g) → NH4Cl(s)

Formation of Amides and Nitrides:

With strong electro positive metals such as sodium, ammonia forms amides while it forms nitrides with metals like magnesium.

2Na + 2NH3 → 2NaNH2 + H2
3Mg + 2NH3 → Mg3N2 + 3H2

With Metallic Salts:

Ammonia reacts with metallic salts to give metal hydroxides (in case of Fe) or forming complexes (in case Cu)

Group 15 (Nitrogen Group) Elements img 6

Formation of Amines:

Ammonia forms ammonated compounds by ion dipole attraction. Eg. [CaCl2.8NH3]. In this, the negative ends of ammonia dipole is attracted to Ca2+ ion.

It can also act as a ligand and form coordination compounds such as [Co(NH3)6]3+, [Ag(NH3)2]+. For example when excess ammonia is added to aqueous solution copper sulphate a deep blue colour compound [Cu(NH3)4]2+ is formed.

Structure of Ammonia

Ammonia molecule is pyramidal in shape N-H bond distance is 1.016 Å and H-H bond distance is 1.645 Å with a bond angle 107°. The structure of ammonia may be regarded as a tetrahedral with one lone pair of electrons in one tetrahedral position hence it has a pyramidal shape as shown in the figure.

Group 15 (Nitrogen Group) Elements img 7

Nitric Acid

Preparation

Nitric acid is prepared by heating equal amounts of potassium or sodium nitrate with concentrated sulphuric acid.

KNO3 + H2SO4 → KHSO4 + HNO3

The temperature is kept as low as possible to avoid decomposition of nitric acid. The acid condenses to a fuming liquid which is coloured brown by the presence of a little nitrogen dioxide which is formed due to the decomposition of nitric acid.

4HNO3 → 4NO2 + 2H2O + O2

Commercial Method of Preparation

Nitric acid prepared in large scales using Ostwald’s process. In this method ammonia from Haber’s process is mixed about 10 times of air. This mixture is preheated and passed into the catalyst chamber where they come in contact with platinum gauze.

The temperature rises to about 1275 K and the metallic gauze brings about the rapid catalytic oxidation of ammonia resulting in the formation of NO, which then oxidised to nitrogen dioxide.

4NH3 + 5O2 → 4NO + 6H2O + 120kJ
2NO + O2 → 2NO2

The nitrogen dioxide produced is passed through a series of adsorption towers. It reacts with water to give nitric acid. Nitric oxide formed is bleached by blowing air.

3NO2 + H2O → 2HNO3 + NO

Properties

Pure nitric acid is colourless. It boils at 86°C. The acid is completely miscible with water forming a constant boiling mixture (98% HNO3, Boiling point 120.5 °C). Fuming nitric acid contains oxides of nitrogen. It decomposes on exposure to sunlight or on being heated, into nitrogen dioxide, water and oxygen.

4HNO3 → 4NO2 + 2H2O + O2

Due to this reaction pure acid or its concentrated solution becomes yellow on standing. In most of the reactions, nitric acid acts as an oxidising agent. Hence the oxidation state changes from + 5 to a lower one. It doesn’t yield hydrogen in its reaction with metals. Nitric acid can act as an acid, an oxidizing agent and an nitrating agent.

As an Acid:

Like other acids it reacts with bases and basic oxides to form salts and water

ZnO + 2HNO3 → Zn(NO3)2 + H2O
3FeO + 10HNO3 → 3Fe(NO3)3 + NO + 5H2O

As an Oxidising Agent:

The nonmetals like carbon, sulphur, phosphorus and iodine are oxidised by nitric acid.

C + 4HNO3 → 2H2O + 4NO2 + CO2
S + 2HNO3 → H2SO4 + 2NO
P4 + 20HNO3 → 4H3PO4 + 4H2O + 20NO2
3I2 + 10HNO3 → 6HIO3 + 10NO + 2H2O
HNO3 + F2 → HF + NO3F
3H2S + 2HNO3 → 3S + 2NO + 4H2O

As an Nitrating Agent:

In organic compounds replacement of a – H atom with – NO2 is often referred as nitration. For example.

Group 15 (Nitrogen Group) Elements img 8

Nitration takes place due to the formation of nitronium ion

HNO3 + H2SO4 → NO2+ + H2O + HSO4

Action of Nitric Acid on Metals

All metals with the exception of gold, platinum, rhodium, iridium and tantalum reacts with nitric acid. Nitric acid oxidises the metals. Some metals such as aluminium, iron, cobalt, nickel and chromium are rendered passive in concentrated acid due to the formation of a layer of their oxides on the metal surface, which prevents the nitric acid from reacting with pure metal.

With weak electropositive metals like tin, arsenic, antimony, tungsten and molybdenum, nitric acid gives metal oxides in which the metal is in the higher oxidation state and the acid is reduced to a lower oxidation state. The most common products evolved when nitric acid reacts with a metal are gases NO2, NO and H2O. Occasionally N2, NH2OH and NH3 are also formed.

Group 15 (Nitrogen Group) Elements img 9

The reactions of metals with nitric acid are explained in 3 steps as follows:

Primary Reaction:
Metal nitrate is formed with the release of nascent hydrogen

M + HNO3 → MNO3 + (H)

Secondary Reaction:

Nascent hydrogen produces the reduction products of nitric acid.

Group 15 (Nitrogen Group) Elements img 10

Tertiary Reaction:

The secondary products either decompose or react to give final products

Decomposition of the Secondary:

Group 15 (Nitrogen Group) Elements img 11

Reaction of Secondary Products:

HNO2 + NH3 → N2 + 2H2O
HNO2 + NH2OH → N2O + 2H2O
HNO2 + HNO3 → 2NO2 + H2O

Examples:

Copper reacts with nitric acid in the following manner

3Cu + 6HNO3 → 3Cu(NO3)2 + 6(H)
6(H) + 3HNO3 → 3HNO2 + 3H2O
3HNO2 → HNO3 + 2NO + H2O

Overall Reaction
3Cu + 8HNO3 → 3Cu(NO3)2 + 2NO + 4H2O

The concentrated acid has a tendency to form nitrogen dioxide

Cu + 4HNO3 → Cu(NO3)2 + 2NO2 + 2H2O

Magnesium reacts with nitric acid in the following way

4Mg + 8HNO3 → 4Mg(NO3)2 + 8[H]
HNO3 + 8H → NH3 + 3H2O
HNO3 + NH3 → NH4NO3

Overall Reaction
4Mg + 10HNO → 4Mg(NO3)2 + N2O + 5H2O

If the acid is diluted we get N2O

4Mg + 10HNO3 → 4Mg(NO3)2 + N2O + 5H2O

Uses of Nitric Acid:

  1. Nitric acid is used as a oxidising agent and in the preparation of aquaregia.
  2. Salts of nitric acid are used in photography (AgNO3) and gunpowder for firearms. (NaNO3)

Oxides and Oxoacids of Nitrogen

Group 15 (Nitrogen Group) Elements img 12

Structures of Oxides of Nitrogen:

Group 15 (Nitrogen Group) Elements img 13

Structures of Oxoacids of Nitrogen:

Group 15 (Nitrogen Group) Elements img 14

Preparation of Oxoacids of Nitrogen:

Name Formula Oxidation state Preparation

Group 15 (Nitrogen Group) Elements img 15

Allotropic Forms of Phosphorus:

Phosphorus has several allotropic modification of which the three forms namely white, red and black phosphorus are most common. The freshly prepared white phosphorus is colourless but becomes pale yellow due to formation of a layer of red phosphorus upon standing.

Hence it is also known as yellow phosphorus. It is poisonous in nature and has a characteristic garlic smell. It glows in the dark due to oxidation which is called phosphorescence. Its ignition temperature is very low and hence it undergoes spontaneous combustion in air at room temperature to give P2O5.

Group 15 (Nitrogen Group) Elements img 16

The white phosphorus can be changed into red phosphorus by heating it to 420°C in the absence of air and light. Unlike white phosphorus it is not poisonous and does not show Phosphorescence. It also does not ignite at low temperatures. The red phosphorus can be converted back into white phosphorus by boiling it in an inert atmosphere and condensing the vapour under water.

The phosphorus has a layer structure and also acts as a semiconductor. The four atoms in phosphorus have polymeric structure with chains of P4 linked tetrahedrally. Unlike nitrogen P≡P is less stable than P-P single bonds. Hence, phosphorus atoms are linked through single bonds rather than triple bonds. In addition to the above two more allotropes namely scarlet and violet phosphorus are also known for phosphorus.

Group 15 (Nitrogen Group) Elements img 17

Properties of Phosphorus:

Phosphorus is highly reactive and has the following important chemical properties

Reaction with Oxygen:

Yellow phosphorus readily catches fire in air giving dense white fumes of phosphorus pentoxide. Red phosphorus also reacts with oxygen on heating to give phosphorus trioxide or phosphorus pentoxide.

Group 15 (Nitrogen Group) Elements img 18

Reaction with Chlorine:

Phosphorus reacts with chlorine to form tri and penta chloride. Yellow phosphorus reacts violently at room temperature, while red phosphorus reacts on heating.

Group 15 (Nitrogen Group) Elements img 19

Reaction with Alkali:

Yellow phosphorus reacts with alkali on boiling in an inert atmosphere liberating phosphine. Here phosphorus act as reducing agent.

Group 15 (Nitrogen Group) Elements img 20

Reaction with Nitric Acid:

When phosphorus is treated with conc. nitric acid it is oxidised to phosphoric acid. This reaction is catalysed by iodine crystals.

Group 15 (Nitrogen Group) Elements img 21

Reaction with Metals:

Phosphorus reacts with metals like Ca and Mg to give phosphides. Metals like sodium and potassium react with phosphorus vigorously.

Group 15 (Nitrogen Group) Elements img 22

Uses of Phosphorus:

  1. The red phosphorus is used in the match boxes
  2. It is also used for the production of certain alloys such as phosphor bronze

Phosphine (PH3)

Phosphine is the most important hydride of phosphorus

Preparation:

Phosphine is prepared by action of sodium hydroxide with white phosphorus in an inert atmosphere of carbon dioxide or hydrogen.

Group 15 (Nitrogen Group) Elements img 23

Phosphine is freed from phosphine dihydride (P2H4) by passing through a freezing mixture. The dihydride
condenses while phosphine does not.

Phosphine can also prepared by the hydrolysis of metallic phosphides with water or dilute mineral acids.

Group 15 (Nitrogen Group) Elements img 24

Phosphine is prepared in pure form by heating phosphorous acid.

Group 15 (Nitrogen Group) Elements img 25

A pure sample of phosphine is prepared by heating phosphonium iodide with caustic soda solution.

Group 15 (Nitrogen Group) Elements img 26

Physical Properties:

It is colourless, poisonous gas with rotten fish smell. It is slightly soluble in water and is neutral to litmus test. It condenses to a colourless liquid at 188 K and freezes to a solid at 139.5 K.

Chemical Properties:

Thermal Stability:

Phosphine decomposes into its elements when heated in absence of air at 317 K or when electric current is passed through it.

Group 15 (Nitrogen Group) Elements img 27

Combustion:

When phosphine is heated with air or oxygen it burns to give meta phosphoric acid.

Group 15 (Nitrogen Group) Elements img 28

Basic Nature:

Phosphine is weakly basic and forms phosphonium salts with halogen acids.

Group 15 (Nitrogen Group) Elements img 29

It reacts with halogens to give phosphorus penta halides.

PH3 + 4Cl2 → PCl5 + 3HCl

Reducing Property:

Phosphine precipitates some metal from their salt solutions.

3AgNO3 + PH3 → Ag3P + 3HNO3

It forms coordination compounds with lewis acids such as boron trichloride.

Group 15 (Nitrogen Group) Elements img 30

Structure:

In phosphine, phosphorus shows sp3 hybridisation. Three orbitals are occupied by bond pair and fourth corner is occupied by lone pair of electrons. Hence, bond angle is reduced to 93.5°. Phosphine has a pyramidal shape.

Uses of Phosphine:

Phosphine is used for producing smoke screen as it gives large smoke. In a ship, a pierced container with a mixture of calcium carbide and calcium phosphide, liberates phosphine and acetylene when thrown into sea. The liberated phosphine catches fire and ignites acetylene. These burning gases serves as a signal to the approaching ships. This is known as Holmes signal.

Group 15 (Nitrogen Group) Elements img 31

Phosphorous Trichloride and Pentachloride:

Phosphorous Trichloride:

Preparation:

When a slow stream of chlorine is passed over white phosphorus, phosphorous trichloride is formed. It can also be obtained by treating white phosphorus with thionyl chloride.

P4 + 8SOCl2 → 4PCl3 + 4SO2 + 2S2Cl2

Properties

When phosphorous trichloride is hydrolysed with cold water it gives phosphorous acid.

PCl3 + 3H2O → H3PO3 + 3HCl

This reaction involves the coordination of a water molecule using a vacant 3d orbital on the phosphorous atom following by elimination of HCl which is similar to hydrolysis of SiCl4.

PCl3 + H2O → PCl3.H2O → P(OH)Cl2 + HCl

This reaction is followed by two more steps to give P(OH)3 or H3PO3.

HPOCl2 + H2O → H2PO2Cl + HCl
H2PO2Cl + H2O → H2PHO3 + HCl

Similar reactions occurs with other molecules that contains alcohols and carboxylic acids.

3C2H5OH + PCl3 → 3C2H5Cl + H3PO3
3C2H5COOH + PCl3 → 3C2H5COCl + H3PO4

Uses of Phosphorus Trichloride:

Phosphorus trichloride is used as a chlorinating agent and for the preparation of H3PO3.

Group 15 (Nitrogen Group) Elements img 32

Phosphorous Pentachloride:

Preparation

When PCl3 is treated with excess chlorine, phosphorous pentachloride is obtained.

PCl3 + Cl2 → PCl5

Chemical Properties

On heating phosphorous pentachloride, it decomposes into phosphorus trichloride and chlorine.

Group 15 (Nitrogen Group) Elements img 33

Phosphorous pentachloride reacts with water to give phosphoryl chloride and orthophosphoric acid.

PCl5 + H2O → POCl3 + 2HCl
POCl3 + 3H2O → H3PO4 + 3HCl

Overall Reaction
PCl5 + 4H2O → H3PO4 + 5HCl

Phosphorous pentachloride reacts with metal to give metal chlorides. It also chlorinates organic compounds similar to phosphorus trichloride.

2Ag + PCl5 → 2 AgCl + PCl3
Sn + 2PCl5 → SnCl4 + 2PCl3
C2H5OH + PCl5 → C2H5Cl + HCl + POCl3
C2H5COOH + PCl5 → C2H5COCl + HCl + POCl3

Uses of Phosphorus Pentachloride

Phosphorous pentachloride is a chlorinating agent and is useful for replacing hydroxyl groups by chlorine atom.

Structure of Oxides and Oxoacids of Phosphorus

Phosphorous forms phosphorous trioxide, phosphorous tetra oxide and phosphorous pentaoxides. In phosphorous trioxide four phosphorous atoms lie at the corners of a tetrahedron and six oxygen atoms along the edges. The P-O bond distance is 165.6 pm which is shorter than the single bond distance of P-O (184 pm) due to pπ-dπ bonding and results in considerable double bond character.

Group 15 (Nitrogen Group) Elements img 34

In P4O10 each P atoms form three bonds to oxygen atom and also an additional coordinate bond with an
oxygen atom.

Terminal P-O bond length is 143 pm, which is less than the expected single bond distance. This may be due to lateral overlap of filled p orbitals of an oxygen atom with empty d orbital on phosphorous.

Group 15 (Nitrogen Group) Elements img 35

Oxoacids of Phosphorous-Structure:

Group 15 (Nitrogen Group) Elements img 36

Oxoacids of Phosphorus-Preparation:

Group 15 (Nitrogen Group) Elements img 37

Group 16 (Oxygen group) Elements:

Occurrence:

Elements belonging group 16 are called chalgogens or ore forming elements as most of the ores are oxides or sulphides. First element oxygen, the most abundant element, exists in both as dioxygen in air (above 20% by weight as well as volume) and in combined form as oxides.

Oxygen and sulphur makes up about 46.6% & 0.034 & of earth crust by weight respectively. Sulphur exists as sulphates (gypsum, epsom etc…) and sulphide (galena, Zinc blende etc…). It is also present in the volcanic ashes. The other elements of this groups are scarce and are often found as selenides, tellurides etc along with
sulphide ores.

Physical Properties:

The common physical properties of the group 16 elements are listed in the Table.

Table 3.2 Physical Properties of Group 16 Elements

Group 15 (Nitrogen Group) Elements img 38