Need and Genesis of Classification
* A periodic table helps classify all the known elements in rows and columns such that the periodicity in their properties becomes apparent. The periodic law defines this periodicity.
* The classification of elements in a periodic table helps predict the properties of unknown elements.
* Johann Dobereiner was the first scientist to observe trends in the properties of elements in 1817; he stated the law of triads based on his observations.
* The first periodic table that applied to all the known elements was formulated by A.E.B de Chancourtois in 1862.
* Chancourtois’s cylindrical periodic table was based on atomic weights of elements and displayed the periodic recurrence in their properties.
* John Alexander Newlands gave the law of Octaves which states that when the elements are arranged in the increasing order of their atomic weights, the properties of every eighth element are similar to that of the first element
* Lothar Meyer and Dmitri Mendeleev independently proposed the Periodic Law, as we know it today, in 1869.
* According to Mendeleev’s periodic law, the physical and chemical properties of the elements are a periodic function of their atomic weights.
* Due to the boldness of Mendeleev’s quantitative predictions and their eventual success, he is considered the father of Modern Periodic
Modern Periodic Table
* The English physicist, Henry Moseley observed regularities in the characteristic X-ray spectra of the elements. He proved that the atomic number is a more fundamental property of an element than its atomic mass.
* Based on his findings, Moseley revised Mendeleev’s Periodic Law and named it as Modern Periodic Law.
* The Modern Periodic law states that the physical and chemical properties of the elements are periodic functions of their atomic numbers.
* The most widely-used periodic table based on the modern periodic law is called the ‘long form’ of the periodic table.
* The long form of periodic table arranges elements based on their atomic numbers and outer electronic configurations in 7 periods and 18 groups.
* Elements having the same principal quantum number are arranged in periods, and elements with similar outer electronic configurations are arranged in groups of the periodic table.
Electronic Configurations of Elements
* The distribution of electrons into the shells and orbitals of an atom is called its electronic configuration.
* The electronic configuration of an atom defines its physical and chemical properties.
* The elements in the long form of periodic table are organized based on the principal quantum numbers of the last-filled orbital.
* The periods in the periodic table indicate the value of principal quantum number ‘n’ for the
valence shell.
* The number of elements in a period is two times the number of atomic orbitals available in the energy level that is being filled. Therefore, the seven periods contain 2, 8, 8,18,18, 32 and 32 elements, respectively.
* The elements in the same vertical column or group have similar valence shell electronic configurations, the same number of valence electronic configurations, the same number of valence electrons, and therefore , similar properties.
Classification of Elements In s, p, d, f Blocks
* The elements of a periodic table are categorized as s-block, p-block, d-block and f-block elements based on the name of the orbital which receives the last electron. The exceptions are Hydrogen and Helium.
* Group 1-2 Constitute the s-block, Groups 3-12 constitute the d- block and Groups 13-18 constitute the p-block.
* Elements in the last two rows at the bottom of the periodic table are called f-block elements or inner transition elements.
* The f-block elements of the series in which the electrons are in 4f orbitals are called lanthanoids.
* The elements of the series in which the electrons are in 5f orbitals are called actinoids.
* Apart from classifying the elements into s, p, d and f blocks, the periodic table also shows the classification of elements into metals, non-metals and metalloids.
* Metals occupy the left-side of the periodic table, followed by few metalloids in the p-block and finally by the non-metals.
* The metallic property of elements increases as we move down a group and decreases as we move left to right across a period in the periodic table.
Periodic Trends: Chemical Reactivity
* The trends in the chemical reactivity of elements can be explained based on the trends in fundamental properties like atomic and ionic radii, ionization enthalpy and electron gain enthalpy.
* The alkali metals and halogens show a higher tendency to combine with oxygen as compared to the elements in the center. The maximum chemical reactivity among the alkali metals is exhibited by the loss of an electron and in halogens is shown by the gain of an electron.
* The alkali metals and halogens of a period can easily combine with oxygen and form acidic and basic oxides.
* Oxides formed by elements in the centre are amphoteric or neutral.
* Down a group, the chemical reactivity generally increases for the alkali and the alkaline earth elements and decreases for the transition elements.
Periodic Trends: Physical Properties II
* The elements in a periodic table demonstrate periodicity in terms of physical properties such as atomic and ionic radii, ionization enthalpy, electron gain enthalpy and electronegativity.
* The radius of an atom is the distance between the nucleus and the outermost orbital of the electrons surrounding the nucleus. It can only be measured approximately as covalent, metallic or Van der Waals radius.
* Atomic radii generally decrease across a period and increase down a group in a periodic table; ionic radii exhibit a similar trend.
* Ionic radius of a cation is always smaller than its parent atom and the ionic radius of an anion is larger than that of the parent atom.
* The more positive the charge on an ion, the smaller the radius. The more the negativity of the charge, the larger the radius.
* Ionization enthalpy is the energy required to remove the most loosely bound electron from an isolated gaseous atom (X) in its ground state. It increases across a period and decreases down a group in a periodic table.
* Electron gain enthalpy is the energy associated with the gain of an electron by an isolated gaseous atom (X) in its ground state. Electron gain enthalpy becomes more negative from left to right in the period and becomes less negative as we move down the group.
* Electronegativity is a qualitative measure of the ability of an atom in a chemical compound to attract shared pair of electrons towards itself.
* Electronegativity increases across a period and decreases down the group.
Periodic Trends: Chemical Properties
* Oxidation states of an element can be deduced from its electronic configuration.
* For representative elements, oxidation number is usually equal to the number of electrons in the outermost shell or eight minus the number of outermost electrons.
* Transition and inner-transition elements exhibit variable oxidation states.
* Oxidation state of an element in a compound is as the charge acquired by its atom based on the electronegativity of other atoms in the molecule.
* Oxidation number of the elements first increases and then decreases across a period.
* The oxidation number in a group remains constant and depends on its generic electronic configuration.
* The number of oxidation stated depicted by the elements varies with groups.
* The second period elements show differences in some properties from the other elements in their respective groups and similarities with the elements present diagonally across them.
* The anomalous behavior of second period elements is due to smaller size, larger charge to radius ratio and higher electronegativity of these elements as compared to the other members in the group.
* The availability of lesser number of orbitals and the peculiarity of their bonding capabilities also contribute to the anomalous behavior of period 2 elements.