As on today 112 elements are known. Which can form millions of compounds. Inorder to remember the properties of chemical elements and their compounds, different Scientists studied properties of all the elements and came to know that some of the elements exhibit similar properties. They were succeeded to know the root cause of similarity in properties. According to Mendeleef, "Periodicity of physical and chemical properties of elements are due to their atomic weights" (fundamental property). From indepth, systematic study of adjacent elements and their compounds, he was able to predict new elements EKa Boron, EKa Aluminium, EKa Silicon, later these elements discovered as Scandium, Gallium and Germanium respectively. In certain pairs of elements, the increasing order of atomic weights was not obeyed, called as "anomalous pairs". They are Ar & K, Co & Ni, Te & I, Th & Pa.
According to Mosleys X-ray spectral studies, he concluded that atomic number instead of atomic weights is the fundamental property.
Based on this, modern periodic law is stated as "The physical and chemical properties of the elements are periodic functions of their electronic configuration." Periodic table used today is long form of periodic table. In which elements are arranged into 7 horizontal rows (Periods) and 18 vertical columns (Groups), 2nd period elements are known as bridge elements. 3rd period elements are known as typical elements. According to IUPAC system, group numbers are given from 1 to 18 from left to right. Where as according to American System they are given I A to VII A, 0 and I B to VIII B. I A to VII A are called representative elements. '0' group are called inert elements. I B to VIII B are d-block elements, excluding II B are transition elements. Lanthanides, Actinides are called as inner transition elements. The elements of a group will have similar properties due to similar outer electronic configuration.
Classification of Elements on the Basis of Electronic Configurations into Blocks
Based on the entry of last electron into s, p, d, f sub-levels, the elements in the periodic table are classified into 4 blocks. They are
's' Block Elements:
The elements in which the last electron enters in 's' sub shell of the valency shell. Since the maximum capacity of s-subshell is 2, 's' block contains two groups IA and IIA. They are placed at the left side of the long form of the periodic table. Their general electronic configuration is ns1 (for I A) and ns2 (for II A). Except H, rest of the elements are metals. They are soft, having low M.Ps & B.Ps they have low ionization potentials. They form ionic compounds. They impart characteristic colours to the flame. They are strong reducing agents.
'p' Block Elements:
The elements in which the last electron enters in 'p' sub shell of valency shell. Since the maximum capacity of p-sub shell is 6, 'p' block contains 6 groups III A to VII A and 0. They are placed at the right side of the long form of the periodic table. Their general electronic configuration is ns2 np1 to ns2 np6. p block consists metals, non metals and semi-metals. They form mostly covalent compounds. '0' group elements are inert and mono atomic. 'p' block elements can also form ionic compounds. As the 16th group elements produce minerals, called "Chalcogens". As the 17th group elements are Sea Salt producers, called "Halogens". Halogens are good oxidising agents.
'd' Block Elements:
The elements in which the last electron enters 'd' sub shell of penultimate shell. Since 'd' sub level accommodates 10 electrons, 'd' block consists 10 groups I B to VII B and VIII B. General electronic configuration of valency shells is (n-1) d1-10 ns1-2. They are placed between s and p block elements in four rows (3d, 4d, 5d & 6d series). All are hard metals with high M.Ps and B.Ps. They form coloured compounds, complex compounds, alloys. They act as catalysts. They are generally paramagnetic. They exhibit variable oxidation states due to participation of n s and
(n-1) d electrons.
'f' Block Elements:
The elements in which the last electron, enters 'f' sub shell of anti penultimate shell. As 'f' orbitals can accommodate 14 electrons, there should be 14 groups. Due to similarity in Chemical behaviour, they are placed at the bottom of the main table in 2 series (4 f - Lanthanides,
5f- actinides). The outer electronic configuration is (n-2)f 1-14 (n-1)d 0-1 ns2. They are heavy metals with high M.Ps and B.Ps. They form coloured, complex compounds. They exhibit variable oxidation states. Actinides are radio active.
Classification of Elements Based on Properties and Reactivity
All the elements are divided into 4 types on the basis of their properties, complete and incomplete electron shells. They are
Noble Gas Elements:
The elements in which 's' & 'p' sub shells of valency shell are completely filled. 'He' has 1s2, other elements have ns2 np6 configuration. They are arranged in '0' group, present at the extreme right side of the periodic table. They are called 'Aerogens' as present in air. They are gases. They are mono atomic and inert. Under controlled conditions Xe can form compounds with F & O.
Representative elements:
The elements in which outer most shell is incompletely filled. The outer electronic configuration is ns1-2 np0-5. Groups IA to VII A are representative elements. Metals, Non metals, Metalloids are present. These elements are chemically reactive and get stable octet configuration either by losing, gaining or by sharing electrons.
Transition elements:
The elements in which outermost and penultimate shells are incompletely filled. The outer electronic configuration is (n-1)d1-10 ns1-2. IB, IIIB to VIIIB groups belongs to this type. All are metals with high M.Ps and B.Ps. They form coloured compounds, complex compounds, alloys. They acts as catalysts. They are generally paramagnetic. They exhibit variable oxidation states. They form interstitial compounds. Most common oxidation state is +2.
Inner transition elements:
The elements in which last 3 shells are incompletely filled. The outer electronic configuration (n-2)f1-14 (n-1)d0-1 ns2. Lanthanides (rare earths), Actinides of IIIB group belongs to this type. Most common oxidation state is +3. The elements after Uranium (trans uranium elements) are man made and radio active. Due to similarity in chemical behaviour, they are placed at the bottom of the main table in 2 series (4f- Lanthnides, 5f - Actinides.) They are heavy metals with high M.Ps and B.Ps. They form coloured, complex compounds.
Periodic Properties & their trends
The reoccurence of similar properties of elements after certain regular intervals 2, 8, 8, 18, 18, 32 when they are arranged in increasing order of their atomic numbers are called "Periodicity" and the properties are called "Periodic Properties".
1. Atomic Radius:
The distance between the centre of the nucleus to the outermost shell of electrons.
They are 3 types of atomic radius. They are
a) Metallic radius: Half of the distance between centres of nuclei of two adjacent metal atoms in the metal crystal.
e.g.: Metallic radius of Na is 1.86 A°.
b) Covalent radius: Half of the distance between centres of nuclei of two atoms of an element bonded by a covalent bond. Twice of the covalent radius is equal to bond length.
e.g.: Covalent radius of Cl in Cl2 is 0.99 A°.
c) Vander Waals radius: Half of the distance between the nuclei of two similar adjacent atoms belong to two neighbouring molecules.
e.g.: Vander Waals radius of Cl2 is 1.8 A°.
Atomic radii in a group increases from top to bottom as newer shells are added. In a period atomic radius decreases from left to right upto 17th group. Atomic radius abruptly increases in noble gases. For other elements, we measure the covalent radius, whereas for noble gases we measure Vander Waals Radius.
e.g.: At radius of F is 0.72 A°, Ne is 1.60 A°.
Atomic radii in transition elements remains unchanged. In a period as the differentiating electron enters the inner (n-1)d orbitals and the screening effect increases. Atomic radii in inner transition elements decreases gradually as the differentiating electron enters anti penultimate (n-2)f orbitals and screening effect becomes poor. The steady decrease in atomic size of Lanthanides is known as "Lanthanide Contraction". Due to this contraction, the properties of 4d & 5d transition elements are more closer, becomes very difficult to separate Lanthanides from their mixture as they have similar crystal structure and properties. Ionic radius is the distance between the centre of the nucleus of ion to which it has influence on its electron cloud in valency shell. Positive ion is smaller than the parent atom due to increased effective nuclear charge. The size of negative ion is always greater than parent atom due to decreased effective nuclear charge. Ions or atoms or molecules having same number of electrons but with different nuclear charge are called ''ISOELECTRONICS''. In this series, size of the ion decreases as the number of positive charges increase and the negative charges decrease. Size of the ion increases as the number of positive charges decrease and the negative charges increase.
e.g.: N-3 > O-2 > F - > Ne > Na+ > Mg+2 > Al+3
2. Ionization Enthalpy:
The minimum amount of energy required to remove one electron from valency shell of neutral, isolated gaseous atom is called first Ionization Enthalpy (I.E1).
M (g) + I.E1 → M+ (g) + e-
The amount of energy required to remove one electron from outermost shell of gaseous unipositive ion is called Second Ionization Enthalpy(I.E2).
M+(g) + I.E2 → M+2 (g) + e-
In unipositive ion, number of protons are more than that of electrons. As a result, effective nuclear charge is increased, it becomes more difficult to remove one electron from unipositive, as it requires high amount of energy, So IE2 greater than IE1.
Factors influencing I.E.:
i) Atomic radius: As atomic radius increases, the distance between nucleus and valence electron increases, attractive force decreases. Hence it requires less amount of energy, i.e., I.E. decreases in a group. In a period, I.E. increases from left to right due to decrease in atomic radius.
iv) Penetration of electrons: Size of orbital increases from s to f. Due to small size, spherical shape, electrons in 's' orbital are more penetrating towards nucleus than that of p orbital. The more the penetrating power of electron towards the nucleus, the more would be the ionization potential. I.E. of electrons in different orbitals in a given shell is s > p > d > f.
I.E. 
penetrating power of electron
v) Electronic Configuration: Atoms whose degenerate orbitals (orbitals with same energy) are either half filled or completely filled are more stable, requires more energy to remove electron. So their I.E. values are higher. I.E. of Be > B and N > O. Element with highest I.E. is "He" and lowest I.E. is Cs.
3. Electron gain enthalpy (or Electron affinity)
The energy released when an electron is added to isolated gaseous neutral atom to make it anion.
X(g) + e- → X - (g) + E1
It can be calculated from Born-Haber cycle. It is a measure of oxidation power of an element. It is zero for zero group elements and Be, Mg.
It increases from left to right in a period, decreases from top to bottom in a group. Second electron affinity (E2) is endothermic.
X - (g) + e-→ X-2 (g) - E2
Electron affinity (or gain enthalpy) of Cl is greater than F and S > O due to repulsion between the added electron and valance electrons of valency shell (of F or S). Cl is having highest, Cs is having lowest electron affinity in the periodic table.
4. Electronegativity:
The tendency of an atom in a compound to attract the shared pair of electrons more towards itself in a covalent bond. It has no units. It increases from left to right in a period, decreases from top to bottom in a group. Most E.N. element is F and least E.N. element is Cs. E.N. can be calculated by Pauling Scale, based on bond energies.
XA - XB = 0.208 
(if ∆ is in K.Cal / Mole)
= 0.1017 (if ∆ is in K.J. / Mole)
∆ = bond polarity = EAB - 
EA-A + EB-B
Mullikan E.N. Valens are 2.8 times greater than Pauling E.N. values. E.N. of zero group elements is zero. Pauling assigned arbitrary valens of Electronegativity 2.1 for H and 4.0 for F. According to Mullikan's scale
Electronegativity is a measure of oxidising power. The nature of the bond if electronegativity difference between two atoms is
= 1.7 (50% ionic + 50% covalent)
> 1.7 (ionic)
< 1.7 (covalent)
= 0 (100% covalent)
5. Valency:
Number of H atoms or double the number of 'O' atoms that can combine with one atom of an element.
In a group all the elements have same valency.
6. Oxidation Number:
The possible +ve or -ve charge exhibited by an atom of an element in the ionic compound. In general oxidation states for IA is +1, IIA is +2, Noble gases is zero. Due to inert pair effect (The reluctance of pair of ns2 electrons present in the valency shell to take part in the bonding), the stable oxidation states of Tl is +1, Pb is +2, Bi is +3. F always exhibit -1 oxidation state due to its high E.N. Oxidation states may be +ve, -ve, 0 or fractional. It should not exceed its group number. Maximum oxidation state +8 is shown by Ru, Os and Xe.
7. Electro-positive character:
The tendency of an atom to lose electrons. Metals exhibit this character. It can be measured in terms of I.P. and electrode potential. This character increases down a group and decreases across a period.
8. Metallic character:
If an element has tendency to lose 1 or 2 or 3 electrons from the valency shell, that element is called metal. This character is inversely proportional to E.N. It increases down a group and decreases across a period.
9. Non metallic character:
If an element has tendency to gain electrons is known as non metal. This character decreases down a group and increases across a period.
10. Acidic & Basic nature of oxides
Oxides of metals are basic, non metals are acidic, metalloids are amphoteric in nature. In a group, basic nature of oxides increases, acidic nature decreases. In a period basic nature decreases, acidic nature increases from left to right.
11. Diagonal relationship
The phenomenon of exhibiting similar properties by second period elements with the third period elements of next group is known as diagonal relationship. It occurs only upto IVA group. The reason for diagonal relationship is similar Ionic size, E.N, polarising power (ionic charge/ ionic radius2)
Li is diagonally related to Mg
