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.
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:
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.
This classification did not receive wide acceptance since only a few elements could be arranged into triads.
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.
The merits of Newlands’ Law of Octaves classification are:
The demerits of Newlands’ law of Octaves are:
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
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.
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:
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.
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.
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.
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.
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.
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).
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.
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 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.
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.