Dobereiner’s Triads: Dobereiner observed that when some elements are arranged in triad, the atomic weight of the middle element is roughly equal to the average of atomic masses of two elements. Since it could be true only for a few elements hence Dobereiner’s Law of Triads was dismissed as coincidence.
Newlands’ Law of Octaves: As per this law, when elements are arranged in ascending order of atomic weights, every eighth element has property similar to the first element.
Mendeleev’s Periodic Table: Mendeleev arranged elements in increasing order of their atomic weights, and based on his observations in similarities in chemical properties he made a periodic table. Mendeleev’s periodic law states that chemical property of elements is a periodic function of their atomic masses. Another scientist Lothar Mayor worked on similar concept and made similar periodic table. But since Mendeleev’s work was published earlier hence Mendeleev is credited with the first periodic table. Mendeleev made some adjustments in order of atomic weights so that elements with similar properties could be fit in a particular group. Moreover, Mendeleev boldly predicted discovery of more elements in future.
Modern Periodic Law
During Mendeleev’s time little was known about internal structure of atom. But many new discoveries about subtomic particle could be possible from the beginning of the 20th century. Henry Moseley (1913) observed regularities in the characteristic X-ray spectra of the elements. When a graph was plotted for X-ray spectra Vs atomic number, a straight line was observed. It was not the case when atomic mass was used instead of atomic number. Thus, Moseley showed that atomic number is a more fundamental property than atomic mass. Accordingly, Mendeleev’s Periodic Law was modified to formulate the Modern Periodic Law. The Modern Periodic Law is as follows:
“The physical and chemical properties of the elements are periodic functions of their atomic numbers.“
Nomenclature of Elements with Atomic Number > 100
|Notation for IUPAC Nomenclature of Elements|
Electronic Configuration of Elements in Periodic Table
Electronic Configuration in Periods
- The period number is associated with the value of n for the outermost valence shell. So, in successive periods, a shell with higher energy level is added. Number of elements in each period is twice the number of orbitals available in the new energy level.
- First period has s orbital, i.e. 1 orbital. So, number of elements in 1st period is 2 × 1 = 2
- Second period has two new orbitals, viz. 2s and 2p. There is one orbital in 2s and three orbitals in 2p. So, number of elements in 2nd period is 2 × 4 = 8
- Subsequently, orbitals 3s and 3p are filled in third period and so the third period has 8 elements.
- The fourth period begins with filling of 4s orbital. After that electrons start filling 3d orbital (3d1 to 3d10). It marks the beginning of 3-d transition elements which begins with Scandium (Z = 21) and ends at Zinc (Z = 30). The 4th period ends at Krypton with filling up of 4p orbital.
- Similarly, the fifth period contains the 4-d transition series.
- The sixth period contains the 4-f transition series; also called the lanthanoid series.
- The seventh period contains the 5-f transition series; or actinoid series.
Electronic Configuration in Groups: Elements in a particular group have same electronic configuration in the outermost orbital.
Electronic Configuration and Types of Elements
Elements in a group constitute a group of elements with similar chemical property. Similarity in chemical property is because of same distribution of electrons in their outermost orbitals. Depending on the type of orbital being filled, elements can be classified into these groups: s-block elements, p-block elements, d-block elements and f-block elements. There are two exceptions to this rule, He and H. Helium belongs to s-group but is kept with p-group elements because helium has fully filled outermost orbital. Hydrogen can behave both as electropositive and electronegative element. Hence, hydrogen has got special place in the periodic table.
- s-block Elements: These are the elements with ns1, ns2 electronic configuration in outermost orbital. These are reactive metals with low ionization enthalpies. They readily lose electrons to form 1+ or 2+ ions. Metals forming 1+ ions are called alkali metals, while those forming 2+ ions are called alkaline earth metals. Metallic character and reactivity increase when we go down the group. They are not found in pure form in nature. These elements generally make ionic compounds, but lithium and beryllium are exceptions.
- p-block Elements: These elements belong to Group 13 to 18. Together with the s-block elements, these are called the Representative Elements or Main Group Elements. Outermost electronic configuration varies from ns2 np2 to ns2 np6 in each period. Each period ends with a noble gas which has fully filled outermost shell. Because of fully filled outermost shell, noble gases show very low chemical reactivity. Halogens (Group 17) and chalcogens (Group 16) precede the noble gases and are important non-metals. Halogens and chalcogens show high negative electron gain enthalpy and readily gain one to two electrons to attain the noble gas configuration. Non-metallic character increases when we move from left to right across a period. Metallic character increases when we move down a group.
- d-block Elements: These elements are in Group 3 to 12 in the centre of the Periodic Table. Filling of d orbitals starts in these elements so they are called d-block elements. The general outer electronic configuration is (n-1)d1-10 ns0-2. All the elements of this group are metals. They generally form colorful ions, exhibit variable valence (oxidation states), paramagnetism and are often used as catalysts. But Zn, Cd and Hg do not show most of the properties of transition elements. These elements form a bridge between chemically active s-block elements and less active elements of Group 13 and 14. So, they are called Transition Elements.
- f-block Elements: The outer electronic configuration of these elements is (n-2)f1-14 (n-1)d0-1 ns2. Since the last electron added to each element is filled in f-orbital, hence they are called f-block elements. They are also called Inner Transition Elements. They are all metals. They are kept at the bottom of the Periodic Table in two groups, viz. actinoids and lanthanoids. Due to the large number of possible oxidation states, the chemistry of early actinoids is more complicated than that of corresponding lanthanoids. Actinoids are radioactive elements. Elements after uranium are called Transuranium Elements.
Metals, Non-metals and Metalloids
Metals: 78% of all known elements are metals and they appear on the left side of the Periodic Table. Metals are generally solid at room temperature, have high melting and boiling points, are good conductors of heat and electricity, and are ductile and malleable.
Non-metals: Non-metals are located at the top right hand side of the Periodic Table. Non-metals are generally solids or gases at room temperature, have low melting and boiling points, are poor conductors of heat and electricity and solids are brittle. The change from metallic to non-metallic is not abrupt; as shown by a thick zig-zag line in the Periodic Table.
Metalloid: Bordering the zig-zag line are elements which show properties of both metals and non-metals. These are called semi-metals or metalloids. Silicon, germanium, arsenic, antimony and tellurium are metalloids.