p Block Elements Class 12

p Block Elements Class 12

p Block Elements Class 12

We learn about s block elements, p block elements and d and f block elements in inorganic chemistry.

p Block Elements Definition

The p-block elements are defined as the elements in the periodic table in which the last electron enters in p-orbital of the atom.

p Block Elements Classification

The p-block elements are classified into six groups. These 6 group elements are as follows-
  • Group-13: B, Al, Ga, In, Ti (ns2np1)
  • Group-14: C,Si, Ge, Sn, Pb (ns²np²)
  • Group-15: N, P, Ar, Sb, Bi (ns²np³)
  • Group-16: O, S, Se, Te, Po (ns²np⁴)
  • Group-17: F, Cl, Br, I, At (ns²np5)
  • Group-18: He, Ne, Ar, Kr, Xe, Rn (ns²np6)


The group 15 elements includes nitrogen (N), phosphorus (P), arsenic (As), antimony (Sb), bismuth (Bi).
As we move down the group in periodic table metallic character increases with increase in atomic number. N and P are non-metal. Ar and Sb are metalloids, while Bi is a metal.


  • Nitrogen gas (N2) is occupied by the 78% of total atmosphere. It is main component of proteins, which plays many critical roles in our body. It occurs as sodium nitrate in the earth’s crust.
  • Phosphorus occurs as minerals of the apatite family, Ca9[PO4]6.CaX2, (X=F, Cl or OH). It is present in the form of proteins in milk and eggs. It is also present in bones and living cell.
  • Arsenic, antimony and bismuth are generally associated with sulphide minerals. The main ores of bismuth are Bi2Oand sulphide Bi2S3.


The general electronic configuration of group 15 elements is ns²np³. The outermost three electrons are unpaired in p-sub shell.
Since d and f sub shell are absent in N, P and present in As, Sb and Bi, that makes N and P behave differently from As, Sb, and Bi.
ElementsAtomic numberCondensed electronic configuration
Nitrogen (N)7[He] 2s²2p³
Phosphorus (P)15[Ne] 3s²3p³
Bismuth (Bi)83[Xe]4f¹⁴5d¹06s²6p³


  • As we move down the group from N to Bi covalent, ionic radii and metallic character increases and value of ionization enthalpy and electronegativity decrease. N is the most electronegative element in the group 15.
  • N is a gas and exists as diatomic molecule with a triple bond between the atoms. P, As and Sb probably exist as tetra atomic molecules containing single bonds between each pair of atom. Bi is mono-atomic metal.
  • The density, atomic volume and boiling points increase down the group of group 15 elements.

Allotropy of group 15 elements

Allotropy is the existence of two or more different physical forms of a Chemical element.
  • Nitrogen in its solid state exists in two form cubic crystal structure and hexagonal crystalline structure.
  • Phosphorus exists in three forms namely white, red and black phosphorus. White phosphorus is waxy solid of P4 tetrahedral and it is also available in gaseous state. Red phosphorus has complex chain structure while black phosphorus is thermodynamically stable having layer structure.
  • Arsenic and antimony also exists in three allotropic forms. Bismuth does not show allotropy.

Chemical Properties

  • Action of air: With oxygen all the elements of group 15 to corresponding oxides on heating.
N2 + O→ 2NO
  • Action of acids: There is no action of hot concentrated nitric acid (HNO3) on nitrogen but with phosphorus and arsenic they forms oxyacids. While Sb and Bi forms oxide and nitrate respectively.
P4 + 20 HNO3 → 4H3PO4 + 20NO2 + 4H2O
  • Action of alkalies (NaOH): nitrogen have no action of alkalies, but phosphorus form phosphine and hypophosphite
  • Action of metal: nitrogen forms metal nitrides.
          6Li + N→ 2Li3N
Phosphorus forms metal phosphides.
6Mg + PO→ 2Mg3P2
Arsenic and antimony forms covalent arsenides and antimonides respectively.
  • Reactivity towards hydrogen: in group 15 all elements forms hydrides with formula AH3, where A= N, P, As, Sb, Bi. e.g.NH3, PH3, etc. Basicity and thermal stability these hydrides decreases in the order of NH3>PH3>AsH3>SbH3>BiHi.e. down the group.


  • Since the general electronic configuration of group 15 elements is ns²np³ they show common oxidation states of +3, +5 and -3.
  • Nitrogen shows wide range of oxidation states from -3 to +3. From N to Bi +3 oxidation state increases and +5 oxidation state decreases.
  • Nitrogen in its solid state exist in two form


Abnormal Behaviour of Nitrogen

Nitrogen differs in other elements of the group 15 due to small size, high electronegativity, high ionization enthalpy and absence of d orbital.
  1. Nitrogen is gas while all the elements of the group are solids under normal temperature.
  2. Nitrogen exists in diatomic molecule (N2) while all the other elements of group exist in tetratomic molecules. (P4, As4, Sb4)
  3. Nitrogen can form pπ-pπ multiple bonds while other elements forms dπ-pπ overlapping.
  4. Nitrogen does not form Penta halides like other elements in group.
  5. Nitrogen shows variable range of oxidation states while other elements of group show limited oxidation states.
Nitrogen is highly electronegative element with absence of d orbital.

Preparation of dinitrogen (N2)

  • Commercial method: On commercial scale it is obtained by the fractional distillation of liquid air.
  • Laboratory method: It is prepared in the laboratory by reaction of ammonium chloride on sodium nitrite.
  • From ammonium dichromate: Dinitrogen is obtained by thermal decomposition of ammonium dichromate.
  • From sodium or barrium azide: Nitrogen is obtained in pure form by thermal decomposition of sodium or barrium azide.

Properties of dinitrogen (N2)

  • Dinitrogen is present in the atmosphere in gaseous state. It is odourless, colourless, tasteless, chemically inert and nontoxic gas.
  • It slightly soluble in water and it has low freezing point (195.3K).
  • At room temperature it is chemically inert. As temperature increases its reactivity also increases.
  • Dinitrogen combines with dioxygen to give nitric oxide at 2000K.
  • On reaction with carbide it gives calcium cyanamide.

Compounds of nitrogen

Ammonia (NH3)

Preparation of ammonia

  1. Laboratory method: Ammonia is obtained by heating a mixture of NH4Cl and slaked lime Ca(OH)2.
  2. Manufacturing method: Ammonia is manufactured by Haber’s process on large scale. In this process N2 and Hare in proportion 1:3 are allowed to react at high pressure 200×105 Pa and temperature 700K. The catalysts used in this process are Fe2O3, K2O and Al2O3.

Properties of ammonia

  • Ammonia is colourless gas having pungent smell and can be toxic.
  • Its freezing point is 198.4K and boiling point is 239.7K.
  • It is highly soluble in water to form NH4OH.
  • It burns in the excess of air to give dinitrogen and water.
  • It reacts with Cl2 to give NCl3.
  • It gives amide to react with active metal like Na and K.

Uses of ammonia

It is used to prepare ammonia nitrate, urea, ammonia phosphate, ammonia sulphate, fertilizer, nitric acid and it is also used as refrigerant.

Nitric acid (HNO3)

Laboratory method: By distillation of nitre with concentrated sulphuric acid.
Commercial scale: On large scale nitric acid is prepared by Ostwald’s process. It is prepared by three steps.
  1. NH3 is oxidized by atmospheric oxygen in presence of catalyst.
  2. Nitric oxide (NO) combines with nitrogen to give NO2.
  3. Nitrogen dioxide dissolves in water to gives nitric acid.

Properties of nitric acid

Pure nitric acid is colourless liquid having freezing point 231.4K and boiling point 355.6K. In gaseous state, it is planner molecule. The pure nitric acid is unstable.

Uses of nitric acid

  1. Nitric acid is used in the manufacture of organic dyes, varnishes, cellulose.
  2. It is used in the preparation of sulphuric acid in lead chambers process.
  3. And it is also used as oxidizer in rocket fuels.

Oxides of Nitrogen

Nitrogen can form number of oxides in its different oxidation states.
Name of oxideOxidation state of NitrogenPhysical appearance and chemical nature
Dinitrogen oxide (N2O)    +1Colourless gas, neutral
Nitrogen monoxide (NO)    +2Colourless gas, neutral, reactive
Dinitrogen trioxide (N2O3)    +3Brown gas, acidic, reactive
Nitrogen dioxide (NO)    +4Blue, solid, acidic
Dinitrogen tetraoxide (N2O2)    +4Colourless solid or liquid, acidic
Dinitrogen pentaoxide (N2O5)    +5Colourless ionic solid, unstable acidic


Allotropic forms of phosphorus

Phosphorus is present in three allotropic forms. Viz. White, red and black phosphorus.
Allotropic forms of phosphorus

  1. White phosphorus: It is translucent, white waxy solid insoluble in water. It is stored in water to protect from air. It has tetrahedral structure (P4).
  2. Red phosphorus: It is tasteless, odourless, non-poisonous reddish violet solid in powder form. It is less reactive than white phosphorus. In red phosphorus P4 tetrahedral units are linked to each other to form chain.
  3. Black phosphorus: It is thermodynamically most stable allotrope. It does not react with water. It has two different forms α-black phosphorus and β-black phosphorus.

Compounds of phosphorus

Phosphine (PH3)

Preparation: phosphine is prepared by action of water on calcium or aluminium phosphide.
Properties of phosphine: It is colourless gas having odour of rotten fish. It is highly poisonous and soluble in water. It does not react with acid and alkalies but reacts with oxygen to form explosive.
Uses of phosphine: It is used in preparation of smoke screens, organo phosphorus compound. It is used as powerful reducing agent.

Phosphorus halides

Phosphourus can form two types of halides namely phosphorus tricloride with general formula PX3 and phosphorus pentacloride with formula PX3 (X=F, Cl, Br, I)


Group 16 elements include oxygen (O), sulphur (S), selenium (Se), tellurium (Te) and polonium. Since first four elements of group forms ore i.e. oxides, sulphides etc. with number of metals they are called as chalcogens. We will also learn properties of p-block elements.


Oxygen is present by mass 46.6% and in atmosphere 21% by volume. Oxygen is most abundant element. Sulphur is less abundant. Selenium and tellurium occurs as sulphide while radioactive pollonium occurs by decay of thorium and uranium.


The valence shell electronic configuration of group 16 is ns²np⁴.
ElementsAtomic numberCondensed electronic configuration

Trends in physical properties

  • Except oxygen (O2) all the other elements of group are solid at common temperature.
  • Sulphur and Selenium is exist in crown puckered ring structure of eight atoms.
  • Due to addition of new shell the atomic as well as ionic radii increases down the group.
  • Metallic character increases down the group. Sulphur is non-metal, selenium and tellurium are semi – metals while polonium is metal.
  • Oxygen shows less possibility towards catenation. The tendency of catenation decreases down the group.
  • The density of group increases down the group due to increase in atomic mass
  • Melting and boiling point increases down the group (except Po).
  • Oxygen is second most electronegative after fluorine. The value of electronegativity decreases down the group.

Chemical properties

  • The reactivity of elements decreases down the group. The group 16 elements react with oxygen to form dioxides.
S + O2 → SO2
  • Sulphur reacts with alkalies to form sulphides and sulphites with non metal group 16 elements combine directly.
  • All the group 16 elements are weakly acidic and acts as weak diprotic acids.
  • All the hydrides of these groups are reducing agent (except H2O).
  • These elements of group can form large number of halides with formula AX6, AX4 and AX2 where A is elements of group 16.

Abnormal behavior of oxygen

  • oxygen differs from all the elements in the group ion several properties due to its
    • Small size
    • Absence of d orbital
    • High electromagnetivity
  • Oxygen is a gas while other elements are solid.
  • Hydride of oxygen is in liquid state i.e. H2O while others are in gaseous state.
  • Molecular oxygen is paramagnetic while others elements are diamagnetic.
  • Oxygen exist in diatomic molecule O2 while other elements are in polyatomic.

Dioxygen (O2)


1. Naturally by photosynthesis

Dioxygen is produced in atmosphere in large quantity by photosynthesis of green plants.
xH2O + xCOsunlight→ (CH2O)+ xO2 (g)

2. Laboratory method

In this method dioxygen is prepared by heating KClO3 using catalyst MnO3.
2KClO→ 2KCl + 3O2(g).

3. Commercial method

On large scale Ois prepared by two methods.
  1. From water: By electrolysis of acidified water, O gas is obtained at canoed.
  2. From air: O2 is obtained by fractional distillation of liquid air.

4. By thermal decomposition

  1. Dioxygen can obtained by thermal decomposition of oxygen rich salts and metallic oxides.
2KNO3 → 2KNO+ O2
2H2O→ 2H2O + O2

Properties of Dioxygen (O2)

  1. Dioxygen is odourless, colourless and tasteless gas which is heavier than air.
  2. It is paramagnetic, it is also present in water in dissolve form which help to sustaining life of plants and animals in water.
  3. It has three stable isotopes i.e. 16O, 17O, 18O.
  4. It does not show much chemical reactivity, it is quite stable at normal temperature.
  5. Dioxygen is not combustible but supports combustion. It forms oxides with different forms of metal.

Oxides are classified into three types:

  1. Acidic oxides: The oxide which combines with water to give an acid is called as acidic oxides.
  2. Basic oxides: The oxides which combines with water to give an base is called as basic oxides.
  3. Amphoteric oxides: The oxides which show acidic and basic characteristics are called as amphoteric oxides.

Uses of dioxygen (O2)

  1. It is most important gas for all living organism in respiration.
  2. Dioxygen is used firemen, sea divers, miners, aviators and also useful for artificial respiration in hospitals.
  3. It is used in rocket fuel, metallurgical processes and in preparation of large number of compounds.

Ozone (O3)

By the action of electrical discharges or ultraviolet radiation on dioxide (O2) ozone can be obtained.
It is present at the height of 20km from earth surface. This layer is called ozonosphere. Ozone (O3) prevents entry of harmful high energetic ultraviolet (UV) rays. It is formed in atmosphere naturally through photochemical reaction.

Preparation of ozone (O3)

By laboratory method slow dry stream of dioxide (O2) is passed through a silent electric discharge dioxide (O2) gets converted into ozone (O3).
    3O2 → 2O3            ΔH=+284KJ

Properties of ozone (O3)

  1. Ozone is thermodynamically unstable having pungent odour.
  2. Solid ozone is violet black, liquid ozone is blue black and gaseous ozone is blue.
  3. In higher concentration it is very dangerous explosive.

Uses of ozone (O3)

  1. It is used as disinfectant and germicide sterilize air and water bleaching agent for starch, oils, waxes, flour etc.
  2. In industry it is used in the manufacture of silk and camphor.
  3. It prevents the harmful ultraviolet (UV) rays of sun from reaching the earth surface.

Sulphur (S)

In free states sulphur occurs in the volcanic regions. It is also occurs in combined state as sulphides (feS2, ZnS, HgS, Cu2S) and sulphates (MgSO4.7H2O, CaSO4.2H2O). Sulphur exists in different allotropic forms as follows
  1. Rhombic sulphur
  2. Monoclinic sulphur
  3. Cyclo – S6
  4. Milk of sulphur
  5. Plastic sulphur
  6. Smolecular sulphur
  7. Colloidal sulphur

Compounds of sulphur (S)

Below are compounds of sulphur.

Sulphur dioxide (SO2)


  • Sulphur dioxide (SO2) is obtained when sulphur burnt in air.
              S(s) + O2(g) → SO2(g)
  • In laboratory method Sulphur dioxide (SO2) is obtained by the action of sulphuric acid on copper turnings.
     Cu + 2H2SO→ CuSO4 + SO2+ 2H2O
  • On commercial method it is on large scale as a by product of the roasting of pyrites.
     4FeS+ 11O → 2Fe2O+ 8SO2

Properties of sulphur dioxide (SO2)

  • Sulphur dioxide (SO2) is colourless with pungent and suffocating smell.
  • It is highly soluble in water to forms sulphurous acid, H2SO3.
SO+ H2O → H2SO3
  • Its boiling point is 263K. In solid state it forms discrete molecules.
  • It is reducing agent.

Uses of sulphur dioxide (SO2)

  1. It is used in the manufacture of H2SO4 and refining of petroleum and sugar industry.
  2. It is used as disinfectant, preservative germicide, bleaching agent and solvent

Sulphuric acid (H2SO4)

Manufacturing process

  • Contact process: This process involves four basic steps.
  • Preparation of sulphur dioxide (SO2) from sulphur
S(g) + O2(s) → SO2
  • Oxidation of sulphur dioxide (SO2) and sulphur trioxide (SO3) by using catalyst vanadium pentaoxide (VO5).
2SO+ O↔ 2SO2
  • By dissolution of sulphur trioxide (SO3) in sulphuric acid (H2SO4), oleum(H2S2O7) is obtained
SO3  +  H2SO4  → H2S2O7
  • Dilution of oleum (H2S2O7) to give sulphuric acid  (H2SO4).
H2S2O+ H2O → 2 H2SO4
  • Lead chamber process: in this process a mixture of SO2, NO and airis treated with steam to obtain H2SO4.
2SO+ O2 + 2H2O → 2 H2SO4

Properties of sulphuric acid (H2SO4)

  • Sulphuric acid (H2SO4) is colourless, dense, oily liquid with boiling point 611K and freezing point 283K.
  • It is strong acid, less volatile and highly viscous due hydrogen bonding. It is known as king of chemicals.
  • It acts as dehydrating agent with many organic compounds.
HCOOH( Formic acid ) H2SO4→ CO + H2O
C6H12O6 (Glucose) H2SO4→ 6C + 6H2O
  • It also acts as a oxidizing agent.
S + H2SO→ 3SO2 + H2O
P4O10 + 6H2O → 4H3PO4
  • When metals are treated with sulphuric acid (H2SO4) gives out hydrogen.
Fe + H2SO→ FeSO+ H2

Uses of sulphuric acid H2SO4

  1. It is used in preparation of HNO3, HCL, H3PO4, Na2CO3, sulphade, oleums, ether.
  2. It is most important chemical in many industries for manufacturing of dyes, fertilizers, detergents, explosives etc.
  3. It is in used lead storage batteries. It is also used as laboratory reagent dehydrating agent and oxidizing agent.


The group 17 elements include elements fluorine (F), chlorine (Cl), bromine (Br), iodine (I) and astatine (At). These elements are highly electronegative which are found in sea water as salts. Astatine is radioactive element.

Electronic configuration

The outermost shell electronic configuration is ns2np5. Each of them is only short of one electron to get stable noble gas electronic configuration.
Elements Atomic numberCondensed electronic configuration                                 
Fluorine (F)            9[He] 2s22p5
Chlorine (Cl)          17[Ne] 2s22p5
Bromine (Br)          35[Ar] 3d104s24p5
Iodine (I)          53[Kr] 4 d105s25p5
Astatine (At)          85[Xe] 4f145d106s26p5

Physical properties

  1. All these elements of group 17 are exist in diatomic form F2, Cl2, Br2 etc. F and Cl are in gaseous state while Br is liquid and I is in solid state at normal temperature.
  2. On moving the down the group 17 atomic and ionic radii increases while ionization enthalpy decreases.
  3. All the element of group possess very high value of electronegativity, fluorine is the most one. These values decrease down the group to iodine.
  4. All halogens are typical non metals having low melting and boiling points.
  5. To attain stable noble gas configuration halogen shows common oxidation state – 1. They also show positive oxidation state +1, +3, +5 and +7.

Chemical properties

  1. Halogens are highly reactive elements. The reactivity decreases down the group.
  2. Halogen has a strong tendency to accept an electron to form anions.
1/2X+ e– → X–    (X = F, Cl, Br, I)
They act as a strong oxidizing agent.
  • On reaction with metals and non metals halogen forms halides
Cu + F2 → CuF2
2AI + 3Br→ AI2Br6
  • Halogen reacts with hydrogen to give hydrogen halides.
H2 + F2 → 2HF
H+ Cl2 → 2HCl
  • All the halogen reacts with NaOH to form halides and hypohalides.
2NaOH + 2Cl2 → NaCl + NaOCl + H2O
  • Halogen combines amongst themselves from many compounds known as inter halogen compounds.

Abnormal behavior of Fluorine (F)

Fluorine’s abnormal behavior due to
  1. High electronegativity.
  2. Small size.
  3. Low F-F band dissociation energy.
  4. Non availability of d – orbital in its valance shell.
  5. It is highly reactive element which forms hydrogen bondings in its hydrides. It exhibits only in -1 oxidation state.

Chlorine (Cl)

The Chlorine is greenish yellow poisonous gas, occurs in extents of 0.14%in the earth crust. And, Chlorine is also presents in sea water as KCl, MgCl and CaCl2.

Preparation of Cl2

  • 1. By the oxidation of HCl on manganese dioxide (MnO2).
MnO+ 4HCl → MnCl2 + Cl2 + H2O
  • 2. On commercial scale it is obtained by
  • 2.1 Deacon process: HCI is oxidized by Atmospheric oxygen.
4HCl + O2 → 2Cl2 + H2O
  • 2.2 Electrolytic process: This process involves the electrolysis of an aqueous solution of brine (NaCl).

Properties of chlorine

  1. It is poisonous, greenish yellow gas with suffocating, pungent smell and heavier than air.
  2. It is soluble in water and the solution is called chlorine water.
  3. Chlorine is neither combustible nor supports combustion. It reacts with number of metals to form their chlorides.
2Na + Cl2 → 2NaCl
2K + Cl→2KCl
  • It dissolves in water to form mixture of HCl and HOCl (hypochloric acid).
Cl2 + H2O ↔ HCl + HOCl
  • Chlorine reacts with hydrogen with very high affinity to form HCl.
H+ Cl2 → 2HCl

Uses of chlorine

  1. It is use as bleaching agent, for sterilization, extraction of metals like gold and platinum.
  2. It is use in manufacture in bleaching powder, dyestuff, explosives, refrigerant, several poisonous gases, hydrochloric acid, plastics, insecticides and several industry solvents.
Hydrochloric Acid (HCl)
It is colourless gas with pungent odour. In liquid its boiling point is 189K and melting point is 159K. It is highly soluble in water. Hydrogen chloride (HCl) is prepared laboratory by heating a mixture of sodium chloride and concentrated sulphuric acid.
NaCl + H2SO4 → NaHSO4 + HCl
Uses of Hydrochloric Acid HCl
It is used in manufacture of chlorine, ammonium chloride, and glucose. It is also used in medicine, galvanizing, preparation of H2 and as a laboratory reagent.

Inter halogen compound

The compound obtained from different halogen are called inter halogen compound. The general composition of inter halogen compounds are XX’, XX’3, XX’5 and XX’7 where X is electropositive group halogen elements.
IClICl3 (unstable)  
These compounds can be prepared by the direct combination or by action of halogen on lower interhalogen compounds. Inter-halogen compounds are covalent, volatile, unstable and diamagnetic in nature.


Group 18 elements include helium, neon, argon, krypton, xenon and radon. These elements are also called as noble or inert gases.
All the elements of group 18 exist in the form of gases in atmosphere. About 1% by volume of air is occupied by noble gases. They are chemically unreactive.
Noble gas have very stable electronic configuration. The valence shell electronic configuration of these elements is ns2np6.
ElementsAtomic numberCondensed electronic configuration
Helium (He)21s2
Neon (Ne)10[He]2s22p6
Argon (Ar)18[Ne]3s23p6
Krypton (Kr)36[Ar]3d104s24p6
Xenon (Xe)54[Kr]4d105s25p6
Radon (Rn)86[Xe]4f145d106s26p6

Physical properties

  1. All the elements are colorless, odorless, and tasteless.
  2. All the elements are always occurs in free the state in nature in monoatomic.
  3. Atomic radii and tendency to liquify increases on moving down the group.
  4. They require a very large amount of energy to remove an electron because of their stable electronic configuration hence value of ionization enthalpy is very high.
  5. Noble gases have low melting and boiling point. They are slightly soluble in water.

Chemical unreactivity

Due to –
  • 1) completely filled valence shells
  • 2) High and positive value of electron gain enthalpy and
  • 3) High ionization enthalpy
noble gases do not take part in chemical reactions i.e. they are chemically unreactive elements. 
There are so many uses of noble gases in our daily life.
Some important concepts in p-block elements are described below:

Electrovalency by accepting electrons and forming anions.

Electrovalency, also known as the ionic valency. It is a term used to describe the valence of an atom or ion when it accepts or donates electrons to form the ionic bonds. When an atom accepts electrons to form negatively charged ions (anions), it does so because it has a tendency to fill its outermost electron shell to achieve a stable electron configuration, typically resembling that of a noble gas.

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