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Periodic Classification of Elements

periodic classification of elements 1

Before we should study about elements let’s first know about the matter. Since childhood you know that matter is anything which has mass, occupies space (have volume), may or may not be visible to our naked eyes. Every day we see objects and touch it is matter. Matters are composed of atoms. Atoms interacted together by subatomic particles. We always want to know about several matters. In every individual’s mind, there remains a question by birth that “what is the last particle when we break any matter?”. Different types of object are composed of different particles. These ultimate particles are elements. In this article, we will study different elements by the simple means of classification. We say it the periodic classification of elements.

Humans are always involved in searching of new elements. It was difficult for them to identify and remember properties of a number of elements. So they made a chart of elements based on their similarity in property. It made easy for them to study about elements. The chart of elements approved by IUPAC (International Union of Pure and Applied Chemistry) contains 117 elements. IUPAC adds the elements to the chart when they are discovered. Now this chart has a special name i.e. “THE MODERN PERIODIC TABLE”. In the first decade of 18th century there were less than 30 elements known. Now we have 117 elements in the periodic table.

Scientists made easy to study about each elements and their specific properties.Also,

  • Matters are classified into three types i.e. elements, compounds and mixtures.
  • Element is one of a class of substances that cannot be separated into simpler substances by chemical means.
  • Till date there are 117 elements known.
  • Each elements has their specific properties.
  • Elements having some similar properties are classified into several groups. This made easy in study of elements.
  • This classification of elements in specific rows and columns in which each next element are periodic in properties is said to be periodic table.

History of the periodic classification of elements:

As the number of elements discovery increased, it became difficult to learn about each elements. In the Year of 1829, there was only 31 elements known. A German chemist Johann Dobereiner tried to arrange elements in the order, keeping elements with similar properties together. The groups made by him was named as Dobereiner’s Triads.

Dobereiner’s Triads:

  • Johann Wolfgang Dobereiner, a chemist, arranged the elements having similar properties into groups.
  • Each group of elements were three in number called
  • When the three elements in a triad were written in the order of increasing atomic masses some similar property was obtained.
  • In the triad, the atomic mass of the middle element was roughly the average of the atomic masses of the other two elements.
  • The middle elements of the triad has intermediate of the properties of other two elements of the triad.
Elements Atomic
Elements Elements
Li 6.9 Ca Cl
Na 23.0 Sr Br
K 39.0 Ba I

doberneir's triads periodic classification of elements

Drawbacks of Dobereiner’s Triads:

  • Dobereiner triads were limited for a few elements. Only 3 triads followed the triad rule.
  1. Li, Na, K
  2. Ca, Sr, Ba
  3. Cl, Br,I
  • As the new elements discovered, they were not accommodated in triads.

Dobereiner’s triad classification failed as all the elements of the periodic table cant be put in triad form. It was a great coincidence that made these triads. After Dobereiner’s failure a French Geologist A.E.B de Chancurtois in 1862 classified the elements in another way.

Chancurtois cylindrical periodic classification of elements:

  • The credit of the first person to make use of atomic weights to produce a classification of periodicity goes to Alexandre-Émile Béguyer de Chancourtois.
  • He arranged elements as a continuous spiral around a metal cylinder and divided it into 16 parts.
  • The atomic weight of oxygen (16) was taken as standard against which all the other elements were compared.

Chancurtois periodic classification of elements

This periodic classification of elements did not took much attention toward it. Widely it was not accepted. In 1865 an English chemist John Alexander Newlands proposed a new table of classification of elements which he called law of octaves.

Newlands’ Law of Octaves:

  • Newland arranged the known elements in order of increasing atomic masses.
  • He found that every eighth element had properties similar to that of the first element.
  • He compared it with music octaves and named it as law of octaves.
  • After lithium the 8th element was sodium and they both exhibited the similar property. Similarly beryllium and magnesium exibits similar properties.

















H Li Be B C N O
F Na Mg Al Si P S
Cl K Ca Cr Ti Mn Fe
Co & Ni Cu Zn Y In As Se
Br Rb Sr Ce & La Zr ·           ·          

Drawbacks of Newlands’ law of Octaves:

  • Newlands’ Law of Octaves was applicable only up to calcium (say for lighter elements only).
  • Newland assumed that only 56 elements existed in nature no more elements will be discovered.
  • Properties of newly discovered elements did not fit to the law of octaves.
  • Newland adjusted two elements in the same slot, also put unlike elements in same slot.

After the failure of Newlands classification, elements classification grabbed attention of many scientists. In 1869, two scientists Russian chemist Dmitri Mendeleev and German chemist Lothar Mayer without knowing each-others work gave a common periodic law. The law stated that “Properties of elements are a periodic function of their atomic weights”. Both Mendeleev and Lothar Meyer made the arrangements of elements in the form of rows and columns. They kept the elements having similar properties in same column. Mendeleev took an extra step. He studied about the chemical properties of elements (the compound formed by them) and changed the order of atomic masses in some cases. He also left space for undiscovered elements. Points related to Mendeleev’s periodic table are:

Mendeleev’s periodic table:

  • Dmitri Ivanovich Mendeleev, a Russian chemist takes the credit for base of development of periodic table.
  • When elements are arranged in order of increasing atomic masses, there is periodic repetition of elements with similar physical and chemical properties.
  • The properties of elements are periodic in properties of their atomic masses.
  • Mendeleev arranged the elements on the basis of its chemical properties.
  • Mendeleev’s periodic Table contains vertical columns called ‘groups’ and horizontal rows called ‘periods’.

mendeleev's periodic classification of elements

Achievements of Mendeleev’s Periodic Table:

  • Some space were left empty for undiscovered elements.
  • Gallium (Ga), Germanium (Ge) and Scandium (Sc) discovered later found its space in periodic table.
  • Position based prediction of properties of elements was there.
  • There was separate space for noble gases as they can be accommodated when discovered without disturbing the main elements.

Drawbacks of Mendeleev’s Periodic Table:

  • There was no explanation of position of isotopes.
  • The position of Hydrogen was not defined.
  • Some atomic masses did not followed the order.

Henry Mosley (English physicist) in 1913 proposed a chart of classification of elements named Modern Periodic Table. This periodic table was mainly based on the studies and observation through X-ray diffraction. The observation made it clear that there is something more fundamental quantity than atomic mass. It was number of protons present in the nucleus of an atom, than called atomic number. It is unique for each element. According to him Atomic number is regarded as the perfect identity of element which is like fingerprint of every individual humans.

Thereafter, Henry Mosley coined Modern Periodic Law which states “the properties of elements are a periodic function of their atomic numbers”. Atomic number also provided an idea of undiscovered elements. Between two known elements if there is gap of atomic number then we can predict how many elements are yet to be discovered. Noble gases also took a separate space in the periodic table.

The main points of Mosley’s Periodic table are:

The Modern Periodic Table:

  • Henry Moseley showed that the atomic number of an element is a more fundamental property than its atomic mass.
  • Properties of elements are a periodic function of their atomic number.
  • Modern Periodic table is the arrangement of elements in order of increasing atomic number Z.
  • Position based prediction of properties of elements gave more precised result when arranged according to atomic number.
  • This periodic table cleared all misconceptions of Mendeleev’s periodic table.

modern periodic table periodic classification of elements

Atomic Number: The number of protons present in nucleus of an atom is its atomic number.

Elements in the Modern periodic table:

  • There are 18 vertical columns known as ‘groups’ and 7 horizontal rows known as ‘periods’.
  • The elements present in any one group have the same number of valence electrons.

Example: See the second group elements, each has 2 electrons in outermost shell.

Be – 2, 2

Mg – 2, 8, 2

Ca – 2, 8, 8, 2

  • The number of shells increases as we go down the group.
  • Elements which have the same number of shells are accommodated in same period.

Example: Na (2,8,1), Mg (2,8,2), Al(2,8,3), Si (2,8,4), P (2,8,5), S (2,8,6), Cl (2,8,7) and Ar (2,8,8). In each of the following there are 3 shells so belongs to third period.

  • Maximum number of electrons that can be accommodated in a shell is given by 2n2. Where ‘n’ is the number of shell.


For K shell 2(1)2 = 2

For L Shell 2(2)2 = 8

For M shell 2(3)2 = 18, but outermost shell can accommodate only 8 electrons.

  • Number of elements in a period is equal to the number of electrons in that numbered shell. As 1st period contain 2 elements same as no. of electrons in K shell i.e. 2. Similarly
  • 2nd period contains 8 elements same as no. of electrons in L shell i.e. 8.
  • Chemical reactivity is determined by the position of an element in the Periodic Table.
  • The valence electrons tells about type and number of bonds formed by an element.

Trends of Periodic Table:

Periodic trends is all about its feature to remain at its position in the periodic table. Periodic table is the proof of Mosley’s statement “the properties of elements are periodic in function of their atomic number”. There are some common trends which is followed throughout the periodic table. Let’s discuss them:

  • Valency:

Everything want to achieve its stable state. We determine the reactivity of elements on the basis of Law of octet. According to this law “An element either gain or loose electron to achieve noble gas configuration. As we know that noble gas configuration means having 8 electrons in their outermost shell. Those elements which has 1, 2, 3 electrons in their outermost shell they achieve noble gas configuration by loosing electrons. The elements which have 5, 6, 7 elements in their outermost shell gain 1, 2, 3 electrons respectively to achieve noble gas configuration. The main points of valency are:

  1. The number of electrons present in outermost shell of an atom determines its valency.
  2. When we move left to right in a period its valency increases from 1 to 4 then decreases to 0. Thus follow specific periodic trends.
  3. Valency remains same in a group.
3rd Period
Na Mg Al Si P S Cl Ar
Valency 1 2 3 4 3 2 1 0
  • Atomic radius:
    It is the radius of an atom. Atomic size is referred as the distance between centre of atom and its outermost shell.
    Atomic radius is the specific followed periodic trend.
  • The atomic radius decreases in moving from left to right along a period. As positive charge (ENC- Effective Nuclear Charge) increases and it pulls electron more towards nucleus, thus size contracts.
3rd period
Atomic radii
(in pm)
Na 186
Mg 160
Al 143
Si 118
P 110
S 104
Cl 99
  • The atomic radius increases down the group. As new shells are added as we move down the group.

atomic redii in group 1 periodic classification of elements

  • Metallic and Non-metallic Properties:
  1. Metallic character is the tendency of an atom to lose electrons.
  2. The metallic character decreases from left to right along a period as the positive charge on nucleus increases and it holds electrons more.
  3. The metallic character increases down the group as the outermost shell are farther from nucleus which make loose of electrons easy.
  4. Electropositive nature of metal also helps it in loosing electrons easily while forming bonds.
  5. Metals are situated at left side of the modern periodic table.
  6. Metallic oxides are basic in nature.


  • Non-metallic character:
  1. Non-metallic character of an atom is the tendency to gain electrons.
  2. On moving left to right across a period Non-metallic character increases as the ENC (Effective nuclear Charge) increases which helps to gain electron easily.
  3. While going down the group Non-metallic character decreases as outer shell is farther from the nucleus gaining of electron becomes difficult.
  4. At the middle of the periodic table i.e. between metals and non-metals there are metalloids which exhibits the property of metals and non-metals both.
  5. Non-metallic oxides are acidic in nature.

Valency, Atomic radius, Metallic and non-metallic character is the basic trends which is followed by the modern periodic table. Other than these there are several important trends which are followed in periodic table. Some of them are ionization enthalpy, electronegativity, Electron gain enthalpy, Electron affinity and many more. The Modern Periodic Table has given us all the possible explanations of the behavior of all the elements present in it. But each and every day our scientists finds something new so to accommodate all these they are still searching for a periodic classification of elements which can provide us each and every solutions without any exceptions.

Properties of Periodic table periods:

Period Orbitals Filled Number of elements Atomic Number
Noble gases
1 1s 2 1-2 He
2 2s 2p 8 3-10 Ne
3 3s 3p 8 11-18 Ar
4 4s 3d 4p 18 19-36 Kr
5 5s 4d 5p 18 37-54 Xe
6 6s 5d 4f 5d 7p 32 55-86 Rn
7 7s 6d 5f 6d 7p 32 87-118 1.        

Periodic table elements:

Elements of periodic table are:

  1. 7 periods
  2. 18 Groups
  3. 4 blocks – (s-block, p-block, d-block, f-block)

Categories of elements present in modern periodic table:

  1. Representative elements (s & p Block)
  2. Transition elements (d-block)
  3. Inner transition elements (f-block)
  4. Noble gases

Properties of Groups in periodic table:

Each group in the periodic table is assigned with its own property which is different from other groups.

Name table of group 1 elements:

Group number Name of the group
Group 1 or IA Alkali metals
Group 2 or IIA Alkaline earth metals
Group 13 or IIA Boron family
Group 14 or IIIA Carbon family
Group 15 or IVA Nitrogen family
Group 16 or VA Oxygen family
Group 17 or VIA Halogen Family
Group 18 Zero Group


Group 1 or IA:

Physical properties of Group 1 elements:

  • Group 1 elements of modern periodic table is said to be alkali metals.
  • Atomic size increases from Lithium to Francium.
  • Alkali metals have the largest atomic radii within the group.
  • Alkali metals are grey solids with shiny surfaces like silver.
  • These elements are soft solids and can be easily cut.
  • Alkali metals have low density as compared to iron like heavy metals.
  • Also, they are good conductors of heat and electricity.
  • These elements can be compressed into sheets or drawn into wires.
  • The density increases as we move down the group from Li to Fr (exception Na is heavier than K).
  • Alkali metals have low melting and boiling points when compared with heavy metals like iron.

atomic radius trends

Chemical properties of Group 1 elements:

  • All the element within the group exibits similar chemical property.
  • Each element has 1 valance electron in group 1.
  • They are highly reactive and reactivity increases as we move down the group.
  • Acts as a reducing agent.
  • These elements have low values of ionization enthalpies as it is easier to remove their valence electrons to attain stability.
  • They are electropositive.
  • Effective nuclear charge (Effective Nuclear Charge) decreases down the group.
  • Carbonate, nitrate, chloride, sulphate, bromide and iodide salts of alkali metals are white solids and are soluble in water.
  • Alkali metals, when exposed in open, can react with oxygen and water vapour in the air.
Elements Electronic Arrangements
Lithium (Li) 2.1
Sodium (Na) 2.8.1
Potassium (K)
Rubidium (Rb)
Caesium (Cs)
Francium (Fr)

Group 2 or IIA:

Physical properties:

  • Group 2 elements of modern periodic table is said to be alkaline earth metals.
  • Atomic size increases from beryllium to radium.
  • Alkaline earth metals are soft solids with surfaces like silver.
  • These metals have a little higher density as compared to alkali metals.
  • Also, they are good conductors of heat and electricity.
  • These elements are compressed into sheets or drawn into wires.
  • These elements have low metallic character than group 1.
  • Group 2 elements have low melting and boiling points as size is larger thus bound loosely.

Chemical properties:

  • All the element within the 2nd group exibits similar chemical property.
  • Each element has 2 valance electron in group 2.
  • Oxidation state of group 2 elements are +2.
  • They are reactive and reactivity increases as we move down the group.
  • Also, acts as a reducing agent.
  • They are electropositive.
  • Effective nuclear charge (Effective Nuclear Charge) decreases down the group.
  • Carbonate, nitrate, chloride, sulphate, bromide and iodide salts of alkaline earth metals are white solids and are soluble in water.
  • They have low ionization energies.
  • Alkaline earth metals and Lithium generally form hydrated salts.
  • The extent of hydration decreases down the group as it decreases with size.
Elements Electronic Arrangements
Beryllium (Be) 2.2
Magnesium (Mg) 2.8.2
Calcium (Ca)
Strontium (Sr)
Barium (Ba)
Radium (Ra)

Group 13 or IIIA elements:

  • P-block starts with group 13 or IIIA elements.
  • Elements in this group are Boron (B), Aluminium(Al), Gallium(Ga), Indium (In), Thallium(Th).
  • Atomic radius increases as we move down the group (Tl has the largest atomic radius). Also, Atomic radius of Gallium (Ga) is less than Aluminium (Al) due to poor shielding (say screening) effect of d-orbital electrons.
  • Reactivity decreases down the group so electrode potential increases when we move down the group.
  • Ionization Energy also decreases going down the group.
  • The only metalloid in the group is Boron while the rest of the elements are metals.
  • Gallium is the only element in this group which is liquid at temperatures over 30o
  • These elements are accommodated with 3 electrons in their valance shell.
  • Electronic configuration is ns2np1.
  • Density increases down the group with Thallium (Tl) having the highest density.
  • As we move down the group, boiling point decreases.
  • Electronegativity decreases down the group up to Aluminium but then increases. This is due to the abnormality in atomic size.
  • Bottom elements are more stable in their +1 oxidation state. Therefore, Thallium is more stable than Aluminium in +1 state.
  • Lanthanide contraction is seen as indium has lesser atomic size than thallium.
  • At high temperatures each element of this family combines with Oxygen to form X2O3.
  • Boron doesn’t show any chemical activity at room temperature.
  • Oxides of Boron is Acidic in nature.
  • Oxides of Aluminium and gallium is Amphoteric in nature As it reacts with acid as well as base.
  • Oxides of thallium is basic in nature.
  • Acidic Nature of Hydroxides reduces down the group.
  • These elements are Lewis acids as they are eager to accept the lone pair of electrons.

Industrial use of Group 13 or IIIA elements:

  • One of the useful product of boron is Borax (Na2B4O7.10H2O). Borax is used in:
  1. Bleaching agent as Home cleaning products.
  2. Gel electrophoresis as a pH buffer.
  3. Synthesize of organic compounds.
  4. When metaborates are exposed to flames it shows several colour characteristics so used in lab test for transition metals.
  5. In soldering as flux.
  • Another product of boron is Orthoboric acid (H3BO3) which is used in:
  1. Fibreglass and borosilicate glass manufacturing.
  2. Making of talc or dry lubricants for carom boards
  3. LCD display screen glass manufacturing.
  • The element of high Industrial use is Aluminium. Its uses are:
  1. Almost every household materials make the use of Aluminium as cooking vessels, refrigerators, kitchen wares etc.
  2. Aluminium foil is used in packaging food materials and medicines.
  3. Electric power line wires is also made up of aluminium.
  4. Manufacture of aircraft bodies.
  • Gallium Arsenide is used in manufacture of semiconductors, enhancers, solar cells.
  • Indium is used in display gadgets and warm reflectors.
  • Products of Lithium and boron’s uses:
  1. LiAlH4 (Lithium Aluminium Hydride) is a widely used reducing agents.
  2. NaBH4 (Sodium borohydride) is also a better reducing agent.

Group 14 or IVA Elements:

Element Symbol Atomic No. Electronic Configuration
Carbon c 6 [He] 2s2 2p2
Silicon Si 14 [Ne] 3s2 3p2
Germanium Ge 32 [Ar] 3d10 4s2 4p2
Tin Sn 50 [Kr] 4d10 5s2 5p2
Lead Pb 82 [Xe] 5d10 6s2 6p2


  • Group 14 is the 2nd group in p-block.
  • Elements in this group are Carbon (C), Silicon (Si), Germanium (Ge), Tin (Sn), Lead (Pb) and flerovium (Fl).
  • Carbon is the most general element of the earth and is constituted in almost everything.
  • The atomic radii increase down the group but it increases much to silicon but increases less from silicon (Si) to lead (Pb) poor screening effect of full filled d and f orbitals.
  • Metallic properties increase down the group.
  • In this group carbon is a non-metal, silicon and germanium are metalloids, tin and lead are poor metals.
  • Their valance shell contains 4 electrons.
  • Electronic configuration for this group is ns2np2.
  • The Group 14 elements has oxidation states of +4 but for the heavier elements it is +2 due to the inert pair effect.
  • Ionization enthalpies decrease down the group as the distance from the nucleus increases. We see that there is a substantial decrease of ionization enthalpy from carbon to silicon. But after that, the difference is less considerable.
  • Also, there is slight increment in ionization enthalpy from tin to lead due to the poor screening effect of the d and f orbitals.
  • When we talk of electronegativity it is almost same from silicon to lead, but is more electronegative than corresponding group 13 elements.
  • Melting and Boiling points are higher than the corresponding elements of group 13.
  • As we move down the group Boiling point decreases from silicon (Si) to lead (Pb).
  • Bottom elements (germanium to lead) are more stable in +2 oxidation state rather than in +4. Lead acts as oxidizing agent in +4 state but is stable in +2 state.
  • Density decreases down the group.
Element  Carbon Silicon  Germanium  Tin Lead 
Density (g/cc) (3.5-diamond
2.34 5.32 7.26 11.3
  • Except carbon the elements of this group has a better tendency to form the complexes. As they Electron pairs from donors easily.
  • When these elements heated with oxygen, two types of oxides are formed, monoxides and dioxides. Dioxides are more acidic than monoxides.
  • When we move down the group Acidic nature of dioxides decreases. Also, tin and lead dioxide is amphoteric.
  • As in earlier section we have learnt that carbon has a unique property called catenation by which it can form long chains with other carbon atoms.
  • Catenation property decreases down the group. Also, the bottom element does not show catenation due to its large size.
  • Except Lead (Pb), all the other elements of 14 or IV A group show allotropy.
  • Different allotropes of carbon as graphite, diamond and buckminsterfullerenes exist due to catenation and formation of pπ– pπ multiple bonds by carbon.
  • Carbon combines with hydrogen to form methane (CH4). The hydride formed is covalent in nature.
  • Carbon reacts with all halides to form a tetrahalide – MX4 . These halides is thermally very stable.
  • Carbon forms both organic acids (acids having -COOH group) and inorganic acids(H2CO3).

Uses of Group 14 or IVA:

  • Uses of carbon:
  1. An allotrope of carbon i.e. graphite is used in pencils, rockets and missile components.
  2. Diamond another allotrope of carbon is used in jewellery and in cutting equipment.
  • Uses of silicon:
  1. A silicon compound “Silicone” is used to make lubricants and for vacuum pumps.
  2. Silicon oxide is used in making glass.
  3. Also used in the manufacture of cement.
  • Germanium and silicon are important semiconductors.
  • Tin and lead are used in alloys and to make pigments and paints. Lead is used to shield various types of radiation.

Group 15 or VA Elements:

Element Symbol Atomic No. Electronic Configuration
Nitrogen N 7 [He] 2s2 2p3
Phosphorous P 15 [Ne] 3s2 3p3
Arsenic As 33 [Ar] 3d10 4s2 4p3
Antimony Sb 51 [Kr] 4d10 5s2 5p3
Bismuth Bi 83 [Xe] 5d10 6s2 6p3


  • Group-15 is the 3rd group of P-block
  • Elements present in this group are Nitrogen (N), Phosphorous (P), Arsenic (As), Antimony (Sb) and Bismuth (Bi).
  • Nitrogen being the primary member is the main constituent of the earth’s air, and records for 78% of it by volume.
  • Metallic character increases down the group.
  • Nitrogen and Phosphorous are non-metals, arsenic and antimony are metalloids while bismuth is a metal.
  • Nitrogen is a diatomic gas, while the other elements are solids in nature.
  • These elements contain five electrons in the valance shell.
  • Electronic configuration of this group is (ns2np3). There is filled s-orbital and three electrons in the p-orbital.
  • Elements are more stable in this group as the p-orbital is half-filled.
  • When we move down the group covalent and ionic radii increase.
  • As atomic size increases down the group thus ionisation enthalpy decreases down the group.
  • Ionisation enthalpy of this group is higher than the corresponding elements of group 15 due to extra stability of half-filled p orbitals.
  • Electronegativity decreases down the group.
  • Boiling point first increases up to arsenic and then decreases till bismuth.
  • As we move down the group stability of -3 and +5 oxidation state decreases.
  • All elements of this group shows allotropy except Nitrogen.
  • Arsenic exists in three allotropic structures – black, grey, and yellow.
  • Antimony also has three main allotropic structures, yellow, metallic and explosive.
  • The special thing is that Nitrogen can form multiple pπ– pπ bonds, won the other hand phosphorous and arsenic can form dπ– pπ bonds.
  • Oxides of group 15 elements formed are of two types having oxidation states +3 as well as +5.
  • Oxides of elements having higher oxidation state are more acidic.
  • When we move down the group acidic character of these oxides decreases.
  • All elements react with metals and the compounds formed has -3 oxidations states.
  • Dinitrogen is a diatomic substance with a triple bond between the two molecules.

Uses of Nitrogen group elements:

  • Uses of nitrogen are:
  1. It is used as a coolant,
  2. In making fertilizers and is also an important raw material of chemical industries.
  3. Used in making dyes.
  4. In manufacturing plastics, nylon and explosives.
  5. It is also used in cryopreservation, pharmaceuticals and X-ray detectors.
  6. Making ammonia by Haber’s process.
  • Uses of phosphorous are:
  1. It is used in fertilizers along with nitrogen,
  2. Red phosphorous is used in making matches to light fires,
  3. softening hard water
  4. it is also used in making smoke bombs.
  5. Phosphorus is also important in the production of steel.
  6. Phosphates are also used in the production of special glasses and fine chinaware.
  • Uses of Arsenic are is used
  1. in semiconductors in the form of gallium arsenide.
  2. Arsenic-lead alloys are used to manufacture bullets.
  3. in wood preservation
  4. It is also used in insectisides.
  • Antimony is used as a dopant in semiconductors also is alloyed with lead to increase lead’s durability.
  • Bismuth is used as a catalyst to make acrylic fibres and in cosmetics.

Group 16 or VI A elements:

Element Symbol Atomic No. Electronic Configuration
Oxygen O 8 [He] 2s2 2p4
Sulphur S 16 [Ne] 3s2 3p4
Selenium Se 34 [Ar] 3d10 4s2 4p4
Tellurium Te 52 [Kr] 4d10 5s2 5p4
Polonium Po 84 [Xe] 5d10 6s2 6p4


  • Group-16 or VI A is the 4th group in P-block
  • The elements present in this group are Oxygen (O), Sulphur (S), Selenium (Se), Tellurium (Te) and Polonium (Po).
  • Metallic character decreases down the group.
  • Oxygen (O) and Sulphur (O) are non-metals, selenium (Se) and tellurium (Te) are metalloids, whereas polonium (Po) is a metal.
  • These elements have 6 electrons in their valance shell i.e. a filled s-orbital and four electrons in the p-orbital. (ns2np4).
  • Atomic radii and Ionic radii increases down the group.
  • Melting and boiling points increase down the group.
  • Ionization enthalpy as in other p-block groups decreases down the group. As there is increase in size.
  • When we compare the increment in size with group 15 than is less because they have more stable Half-filled p-orbitals.
  • When we move down the group its electronegativity decreases.
  • Every elements in this group exhibit allotropy.
  • Oxygen has two allotropes, O2 (Dioxygen) and O3 (Ozone).
  • Bottom elements less stable in oxidation states of -2 and +6.
  • Oxygen shows only -2 oxidation state, but there is one exceptional case of OF2(+2 oxidation state is seen in oxygen).
  • Due to inert pair effect stability of +4 oxidation state increases as we go down the group.
  • Sulfur commonly having -2 oxidation state can also exist at a +4 and +6 state.
  • Se, Te, and Po can exist in +2, +4, and +6 oxidation states.
  • Acidic character of hydrides of elements and their reducing ability increases as we move down the group. Also not considering H2
  • On the other hand reducing property of oxides decreases as we move down the group and their oxides show acidic nature.

Uses of Group 16 or VIA elements:

  • Uses of Oxygen are:
  1. used as antifreeze
  2. to make polyester
  3. Oxygen is used as liquid fuel in rockets.
  4. Oxygen Helps Clean Waste Water.
  5. in oxy-acetylene flames used for welding
  6. in smelting iron into steel
  7. as a breathing gas by living organisms on earth.
  • Uses of Sulphur:
  1. In making gun powder, matches and in fireworks.
  2. Sulphur in the form of sulphuric acid is used in the manufacture of fertilizers.
  3. In manufacturing paints, dye and storage batteries.
  4. Also in petroleum refining.
  5. Sulphur is used in ointments for curing skin diseases.
  6. Cleansing of metals
  7. As a laboratory reagent.
  8. Sulphur is also used in the vulcanization of rubber.
  • Selenium is used in manganese electrolysis and in alloys with bismuth also used in alternative medicine as an aid to treat Hashimoto’s thyroiditis
  • Tellurium is used in solar panels. Also its addition to lead decreases the corrosive action of sulfuric acid on lead and improves its strength and hardness.

Group 17 or VII A elements:

Element Symbol Atomic No. Electronic Configuration
Fluorine F 9 [He] 2s2 2p5
Chlorine Cl 17 [Ne] 3s2 3p5
Bromine Br 35 [Ar] 3d10 4s2 4p5
Iodine I 53 [Kr] 4d10 5s2 5p5
Astatine At 85 [Xe] 5d10 6s2 6p5
  • Group-17 is the 5th group of p-block and is known as halogen family.
  • Elements present in group 17 are Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I) and Astatine (At).
  • These elements have 7 electrons in their valance shell.
  • Electronic configuration for this group is (ns2np5).
  • The speciality of this group is that its elements have the smallest atomic radii in their respective periods.
  • Atomic size and Ionic radii increases as we move down the group.
  • Melting and boiling points also increases down the group.
  • Ionisation enthalpy due to atomic size factor decreases down the group.
  • Electronegativity also decreases down the group.
  • Oxidising power as well decreases as we move down the group.
  • Affinity towards hydrogen and ionic character of metal halides in this group also decreases as we move down.
  • When we talk of Bond dissociation enthalpy ( the energy released when a bond is broken) decreases as we go down the group, however, due to small size and large electron-electron repulsion in fluorine, its bond dissociation enthalpy is close to that of chlorine.
  • Oxidation states seen by the elements of this group are -1, +1, +3, +5 and +7.
  • Fluorine exists in -1 oxidation state only.
  • The acidic strength of hydrides increases down the group (I>Br>Cl).
  • Also, stability of oxides increases in the same order.
  • Halogens generally form diatomic molecules (of the form X2 , where X denotes a halogen atom) in their elemental states.
  • These elements act as oxidizing agents as they exhibit the property to oxidize metals.
  • Halogens form salts when they react with metals

Use of Group 17 elements:

  • Uses of Fluorine:
  1. Its compound are used in toothpaste.
  2. Also in drinking water to prevent dental cavity.
  3. It is used to etch the glass in light bulbs and in other products.
  4. Used to make nuclear material in reactors.
  5. Also, it is used to insulate electrical equipment in the large transformers that bring energy from the power plant to our house.
  • Uses of Chlorine:
  1. Generally it is used as an antiseptic.
  2. Also is used to make drinking water safe.
  3. Used to treat swimming pools.
  4. In a wider extent it is used in the production of paper products, plastics, dyes, textiles, medicines, antiseptics, insecticides, solvents and paints.
  • Uses of Bromine:
  1. One of the major uses of bromine is in water purification.
  2. Bromine is used in photography.
  3. Brominated compounds are used for water treatment in swimming pools and hot tubs.
  4. It is used to control algae and bacterial growth in industrial processes.
  • Iodine reduces thyroid hormones. Iodine is used to make photographic film and, when iodine mixed with alcohol acts as an antiseptic for external wounds.

Group 18 or VIII A elements:

Element Symbol Atomic No. Electronic Configuration
Helium He 2 1s2
Neon Ne 10 [He] 2s2 2p6
Argon Ar 18 [Ne] 3s2 3p6
Krypton Kr 36 [Ar] 3d10 4s2 4p6
Xenon Xe 54 [Kr] 4d10 5s2 5p6
Radon Rn 86 [Xe] 5d10 6s2 6p6
  • Group 18 or VIII A is the last group of p-block elements.
  • Elements present in this group are: Helium (He), Neon (Ne), Argon (Ar), Krypton (Kr), Xenon (Xe) and Radon (Rn).
  • Except helium each element has p-
  • Only gases are there in this group.
  • Their valance shell is completely filled with 8 electrons.
  • Electronic configuration of elements present in this group is ns2np6.
  • Atomic radii increase down the group.
  • Melting and boiling points also increases as we move down the group.
  • Fairly nonreactive.
  • Complete outer electron or valence shell (oxidation number = 0)
  • High ionization energies due to higher stability and decreases as we move down the group.
  • These elements possess very low electronegativities.
  • All are monatomic gases at room temperature.
  • No color, odor, or flavor under ordinary conditions
  • Nonflammable
  • At low pressure, they will conduct electricity and fluoresce.

Uses of Noble Gas:

  • Uses of helium:
  1. Helium is used as a cooling medium for the Large Hadron Collider (LHC).
  2. It is also used as the superconducting magnets in MRI scanners and NMR spectrometers.
  3. It is also used to keep satellite instruments cool.
  4. Helium is used to fill decorative balloons, weather balloons and airships because of its low density.
  5. Deep-sea divers uses mixture of 80% helium and 20% oxygen as an artificial atmosphere for working under pressurised conditions.
  • Uses of Neon:
  1. Widely used in ‘neon signs’ for advertising.
  2. Neon is also used to make high-voltage indicators and switching gear, lightning arresters, diving equipment and lasers.
  • Uses of Argon:
  1. Argon is used in electrical light bulbs, fluorescent tubes, photo tubes, and glow tubes.
  2. Argon is used arc welding such as gas metal arc welding and gas tungsten arc welding,
  3. Used in the processing of titanium and other reactive elements.
  4. An argon atmosphere is also used for growing crystals of silicon and germanium.

D-Block elements (transition metals):

Those elements that have partially filled d orbitals are transition metals. They have high tensile strength, electrical and thermal conductivity, malleability, ductility, as well as metallic lustre.

Some of the common properties of D-Block elements are:

  • Transition metals have large charge/radius ratio.
  • These are hard and have high densities.
  • Also they have high melting and boiling points.
  • They form compounds which are often paramagnetic.
  • We calculate spin-only magnetic moment by the formula: μ = √n(n+2). Where n is the number of unpaired electrons. Its unit is Bohr magneton (BM).
  • Always show variable oxidation states. This leads to a large number of compounds different types of reactions involved.
  • D-block elements form coloured ions and compounds.
  • They form compounds with profound catalytic activity.
  • Also, they form stable complexes.
  • We consider only M2+/M and M3+/M2+ reduction potentials for the states M2+ and M3+

To practice questions based on this chapter please click here.  Q/A on Periodic Classification of Elements

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