Periodic Classification of Elements

Before the beginning of the eighteenth century, only a few elements were known, so it was quite easy to study and remember the properties of those elements and their compounds individually. However, by the middle of the nineteenth century, more the than sixty elements had been discovered.

The number of compounds formed by them was also enormous. With the increasing number of elements, it was becoming more and more difficult to study their properties individually. Therefore, the need for their classification was felt. This led to the classifications of various elements into groups which helped in the systematic study of elements.

Development of Classification

Scientists after many attempts were successful in arranging various elements into groups. They realised that even though every element is different from others, yet there are a few similarities among some elements. Accordingly, similar elements were arranged into groups which led to classification.

The first classification of elements was into 2 groups-metals and non-metals. This classification served only limited purpose mainly because some elements like germanium and antimony showed the properties of both metals and non-metals. They could not be placed in any of the two classes.

Four major attempts were made for classification of elements:

  1. Dobereiner’s Triads
  2. Newlands’ Law of Octaves
  3. Mendeleev’s Periodic Law & Periodic Tables
  4. Modern Periodic Table

Dobereiner’s Triads

In 1829, J.W. Dobereiner, a German chemist made groups of three elements each and called them triads. All three elements of a triad were similar in their physical and chemical properties. He proposed a law known as Dobereiner’s law of triads. According to this law, when elements are arranged in order of increasing atomic mass, the atomic mass of the middle element was nearly equal to the arithmetic mean of the other two and its properties were intermediate between those of the other two.

law of triads

This classification did not receive wide acceptance since only a few elements could be arranged into triads.

Newlands’ Law of Octaves

In 1864, an English chemist John Alexander Newlands arranged the elements in the increasing order of their atomic masses (then called atomic weight). He observed that every eighth element had properties similar to the first element. Newlands called it the Law of Octaves.

Starting from lithium (Li), the eighth element is sodium (Na) and its properties are similar to those of the lithium. Similarly, beryllium (Be), magnesium (Mg) and calcium (Ca) show similar properties. Fluorine (F) and chlorine (Cl) are also similar chemically.

law of octaves

The merits of Newlands’ Law of Octaves classification are:

  • (i) Atomic mass was made the basis of classification.
  • (ii) Periodicity of properties (the repetition of properties after a certain interval) was recognised for the first time.

The demerits of Newlands’ law of Octaves are:

  • (i) It was not applicable to elements of atomic masses higher than 40 u. Hence, all the 60 elements known at that time, could not be classified according to this criterion.
  • (ii) With the discovery of noble gases, it was found that it was the ninth element which had the properties similar to the first one and not the eighth element. This resulted in the rejection of the very idea of octaves.

Mendeleev’s Periodic Law and Periodic Table

Mendeleev studied the properties of all the 63 elements known at that time and their compounds. On arranging the elements in the increasing order of atomic masses, he observed that the elements with similar properties occur periodically. In 1869, he stated this observation in the form of the following statement which is known as the Mendeleev’s Periodic Law.

The chemical and physical properties of elements are a periodic function of their atomic masses.

A periodic function is the one which repeats itself after a certain interval. Mendeleev arranged the elements in the form of a table which is known as the Mendeleev’s Periodic Table.

Mendeleev’s Periodic Table

Mendeleev arranged the elements in the increasing order of their atomic masses in horizontal rows till he came across an element whose properties were similar to those of the first element. Then he placed this element below the first element and thus started the second row of elements.

The success of Mendeleev’s classification was due to the fact that he laid more emphasis on the properties of elements rather than on atomic masses. Occasionally, he could not find an element that would fit in a particular position. He left such positions vacant for the elements that were yet to be discovered.

Main Features of Mendeleev’s Periodic Table

  1. The elements are arranged in rows and columns in the periodic table.
  2. The horizontal rows are called periods. There are six periods in the periodic table. These are numbered from 1 to 6 (Arabic numerals). Each one of the 4th, 5th and 6th periods have two series of elements.
  3. Properties of elements in a given period show regular gradation (i.e. increase or decrease) from left to right.
  4. The vertical columns present in it are called groups. There are eight groups numbered from I to VIII (Roman numerals).
  5. Groups I to VII are further divided into A and B subgroups. However, group VIII contains three elements in each of the three periods.
  6. All the elements present in a particular group are chemically similar in nature. They also show a regular gradation in their physical and chemical properties from top to bottom.

Merits of Mendeleev’s Periodic Classification

1. Classification of all elements

Mendeleev’s classification included all the 63 elements known at that time on the basis of their atomic mass and facilitated systematic study of elements.

2. Correction of atomic masses

Atomic masses of some elements like Be (beryllium), Au (gold), In (indium) were corrected based on their positions in the table.

3. Prediction of new elements

Mendeleev arranged the elements in the periodic table in increasing order of atomic mass but whenever he could not
find out an element with expected properties, he left a blank space. He left this space blank for an element yet to be discovered.

4. Valency of elements

Mendeleev’s classificaiton helped in understanding the valency of elements. The valency of elements is given by the group number. For example, all the elements in group 1 i.e. lithium, hydrogen, sodium, potassium, rubidium, caesium have valency 1.

Defects of Mendeleev’s Periodic Table

1. Position of Hydrogen

The position of hydrogen which is placed in group IA along with alkali metals is ambiguous as it resembles alkali metals as well as halogens (group VII A).

2. Position of Isotopes

All the isotopes of an element have different atomic masses therefore, each one of them should have been assigned a separate position. On the other hand, they are all chemically similar; hence they should all be placed at the same position. In fact, Mendeleev’s periodic table did not provide any space for different isotopes.

3. Anomalous Pairs of Elements

At some places, an element with greater atomic mass had been placed before an element with lower atomic mass due to their properties. For example, cobalt with higher atomic mass (58.9) was placed before nickel with lower atomic mass (58.7).

4. Grouping of chemically dissimilar elements

Elements such as copper and silver have no resemblance with alkali metals (lithium, sodium etc.), but have been grouped together in the first group.

5. Separation of chemically similar elements

Elements which are chemically similar such as gold and platinum have been placed in separate groups.

Modern Periodic Law

In 1913, Henry Moseley, an English physicist discovered that the atomic number and not the atomic mass is the most fundamental property of an element.

Atomic number (Z) of an element is the number of protons in the nucleus of its atom.

Modern Periodic Law

The Modern Periodic Law states that the chemical and physical properties of elements are periodic functions of their atomic numbers i.e. if elements are arranged in the order of their increasing atomic number, the elements with
similar properties are repeated after certain regular intervals.

Taking atomic number as the basis for classification, removed major defects from it such as anomalous pairs and position of isotopes.

Cause of Periodicity

In the electronic configuration of alkali metals i.e., the first group elements with atomic numbers 3, 11, 19, 37, 55 and 87 (i.e., lithium, sodium, potassium, rubidium, caesium and francium), all the elements have one electron in the outer most shell and so they have similar properties which are as follows:

  1. They are good reducing agents.
  2. They form monovalent cations.
  3. They are soft metals.
  4. They are very reactive and, therefore, found in nature in combined state.
  5. They impart colour to the flame.
  6. They form hydrides with hydrogen.
  7. They form basic oxides with oxygen.
  8. They react with water to form metal hydroxides and liberate hydrogen.

All the elements having similar electronic configuration have similar properties. Thus, the re-occurrence of similar electronic configuration is the cause of periodicity in properties of elements.

Modern Periodic Table

The periodic table based on the modern periodic law is called the Modern Periodic Table. Presently, the accepted modern periodic table is the Long Form of Periodic Table.


There are 18 vertical columns in the periodic table. Each vertical column is called a group. The groups have been numbered from 1 to 18 (in Arabic numerals).

All elements present in a group have similar electronic configurations and have same number of valence electrons.

All elements of group 1 have only one valence electron. Li has electrons in two shells, Na in three, K in four and Rb has electrons in five shells. Similarly all the elements of group 17 have seven valence electrons however the number of shells is increasing from two in fluorine to five in iodine.


There are seven horizontal rows in the periodic table. Each row is called a period. The elements in a period have consecutive atomic numbers. The periods have been numbered from 1 to 7 (in Arabic numerals).

In each period a new shell starts filling up. The period number is also the number of the shell which starts filling up as we move from left to right across that particular period. For example, in elements of 3rd period (N = 3), the third shell (M shell) starts filling up as we move from left to right. The first element of this period, sodium (Na: 2,8,1) has only one electron in its valence shell (third shell) while the last element of this period, argon (Ar: 2,8,8) has eight electrons in its valence shell.

Types of Elements

Main Group Elements

The elements present in groups 1 and 2 on left side and groups 13 to 17 on the right side of the periodic table are called representative or main group elements. Their outermost shells are incomplete,which means their outermost shell has less than eight electrons.

Noble Gases

Group 18 on the extreme right side of the periodic table contains noble gases. Their outermost shells contain 8 electrons except He which contains only 2 electrons.

  • They have 8 electrons in their outermost shell (except He which has 2 electrons).
  • Their combining capacity or valency is zero.
  • They do not react and so are almost inert.
  • All the members are gases.

Transition Elements

The middle block of periodic table (groups 3 to 12) contains transition elements. Their two outermost shells are incomplete. Since these elements represent a transition (change) from the most electropositive element to the most electronegative element, they are named as transition elements.

  • All these elements are metals and have high melting and boiling points.
  • They are good conductors of heat and electricity.
  • Some of these elements get attracted towards magnet.
  • Most of these elements are used as catalyst.
  • They exhibit variable valencies.

Inner Transition Elements

These elements, also called rare-earth elements, are shown separately below the main periodic table. These are two series of 14 elements each. The first series called lanthanoids consists of elements 58 to 71 (Ce to Lu). They all are placed along with the element 57, lanthanum (La) in the same position (group 3, period 6) because of very close resemblance between them.

The second series of 14 rare-earth elements is called actinoids. It consists of elements 90 to 103 (Th to Lr) and they are all placed along with the element 89, actinium (Ac) in the same position (group 3, period 7).

In all rare-earths (lanthanoids and actinoids), three outermost shells are incomplete. They are therefore called inner transition elements.


Metals are present in the left hand portion of the periodic table. The strong metallic elements; alkali metals (Li, Na, K, Rb, Cs, Fr) and alkaline earth metals (Be, Mg, Ca, Sr, Ba, Ra) occupy groups 1 and 2 respectively.


Non-metals occupy the right hand portion of the periodic table. Strong non-metallic elements i.e., halogens (F, Cl, Br, I, At) and chalkogens (O, S, Se, Te, Po) occupy groups 17 and 16 respectively.


Metalloids are the elements that show mixed properties of both metals and non-metals. They are present along the diagonal line starting from group 13 (Boron) and going down to group 16 (Polonium).

Merits of the Modern Periodic Table

1. Position of isotopes

All isotopes of an element have the same atomic number and therefore, occupy the same position in the modern periodic table.

2. Anomalous pairs

The anomaly regarding all these pairs disappears when atomic number is taken as the basis for classification. For example, cobalt (at. no. 27) would naturally come before nickel (at. no. 28) even though its atomic mass is little more than that of nickel.

3. Electronic configuration

This classification is according to the electronic configuration of elements, i.e., the elements having a certain pattern of electronic configuration are placed in the same group of the periodic table. It relates the properties of elements to their electronic configurations.

4. Separation of metals and non-metals

The position of metals, non-metals and metalloids are clearly established in the modern periodic table.

5. Position of transition metals

It makes the position of the transition elements quite clear.

6. Properties of elements

It reflects the differences, the trends and the variations in the properties of the elements in the periodic table.

Periodic Trends in Properties

In a given group, the number of filled shells increases. The number of valence electrons is the same in all the elements of a given group. However, these valence electrons but they are present in higher shells which are farther away from the nucleus. This decreases the force of attraction between the outermost shell and the nucleus as we move downwards in a group.

In a given period, the nuclear charge and the number of valence electrons in a particular shell increase from left to right. This increases the force of attraction between the valence electron and nucleus as we move across a period from left to right.

Atomic Size

Atomic size is the distance between the centre of nucleus and the outermost shell of an isolated atom. It is also known as atomic radius. It is measured in picometre, pm (1 pm = 10–12 m).

Variation of atomic size in periodic table

The size of atoms decreases from left to right in a period but increases from top to bottom in a group.

In a period the atomic number and therefore the positive charge on the nucleus increases gradually. As a result, the electrons are attracted more strongly and they come closer to the nucleus. This decreases the atomic size in a period from left to right.

In a group as one goes down, a new shell is added to the atom which is farther away from the nucleus. Hence electrons move away from the nucleus. This increases the atomic size in a group from top to bottom.

Metallic and Non-metallic Character

The tendency of an element to lose electrons to form cations is called electropositive or metallic character of an element. Alkali metals are most electropositive. The tendency of an element to accept electrons to form anions is called electronegative or non-metallic character of an element.

Variation of Metallic Character in a Group

Metallic character increases from top to bottom in a group as tendency to lose electrons increases. This increases the electropositive character and metallic nature.

Variation of Metallic Character in a Period

Metallic character decreases in a period from left to right. It is because the ionization energy increases in a period. This decreases the electropositive character and metallic nature.