CBSE | NCERT - Explanation of The modern periodic Table Chemistry - E Notes

The modern version of the periodic law is stated as :“
The physical and chemical properties of the elements are the periodic functions of their atomic masses”.

LONG FORM OF THE PERIODIC TABLE

There are many forms of the periodic tale. The long form of the periodic table is the most convenient and the most widely used and is presented here. The horizontal rows are called PERIODS. 

Elements having similar chemical and physical properties appear in vertical columns and are known as GROUPS or FAMILIES. Altogether there are seven periods and 18 groups.

PERIODS
1ST Period Contain only 2 elements namely 1H, 2He and is the shortest period.

IInd Period Contains 8 elements namely 3Li, 4Be, 5B, 6C, 7N, 8O, 9F and 10Ne is known as the short period.

IIIrd Period Contains 8 elements : 11Na, 12Mg, 13Al, 14Si, 15S, 16S, 17Cl, and 18Ar. 

The IIIrd period is also known as short period.

IV Period starts with potassium and contains 18 elements :

19K, 20Ca, 21Sc, 22Ti, 23V, 24Cr, 25Mn, 26Fe, 27Co, 28Ni, 29Cu, 30Zn, 31Ga, 32Ge, 33As, 34Se, 35Br and 36Kr and is known as long period.


V Period Starts with rubidium and contains 18 elements : 37Rb, 38Sr, 39Y, 40Zr, 41Nb, 42Mo, 43Tc, 44Ru, 45Rh, 46PD, 47Ag, 48CD 49In, 50Sn, 51Sb, 52Te, 53I, 54Xe.The fifth


VI Period Consists of 32 elements, starting from cesium (55Sc) and ending with radon. It is called as longest period. This period includes the 4f shell that includes Lanthanides (58Ce, …..71Lu).



VII Period Like the sixth period would have a theoretical maximum of 32 elements. This period however is incomplete and at present contains 19 elements starting from 87Fr (francium) to 92U. All these elements are naturally occurring but rest are radioactive with very short half lives. These also include a part of inner transition elements, 90Th, …. 103Lr.

GROUPS

The atom of the element in a single vertical column have the same or very similar electronic configurations in the highest occupied orbitals and are therefore said to belong to the same GROUP or FAMILY of elements. According to the new IUPAC recommendations the groups are numbered form 1 to 18. Bases on the electronic configuration, we can classify elements in to four types.

1. Noble Gases                                                        2. Representative Elements                

3. Transition Elements and                                     4. Inner Transition Elements

1. Noble Gases
The noble gases are found at the end of each period in group 18. With the exception of helium, these elements have ns2 np6 electronic configuration in the outermost shell. Helium has 1s2 configuration. All the energy levels that are occupied by the electrons are completely filled and this stable arrangement of electrons cannot be easily altered. These elements have very low chemical reactivity.

2. Representative Elements (s and p block elements)
The elements of Group 1 (alkali metals), Group 2(alkaline earth metals) and Group 13,– 17constitute the Representative Elements. For the representative elements the period in which the element is located equal the principal quantum number of the differentiating electron, i.e., if an element is in nth period then the electronic configuration will be either ns1-2 or np1-5 .

Group 1 Consists of 1H(1s1), 3Li(2s1), 11Na(3s1), 19K(4s1), 37Rb(5s1), 55Cs(6s1), 87Fr(7s1).

The common outermost electronic configuration is ns-1.  Elements belonging to this group are known as Alkali Metals.

Group 2 Contains 4Be(2s2), 12Mg(3s2), 20Ca(4s2), 38Sr(5s2), 56Ba(6s2), 88Ra(7s2). The elements belonging to this group are known as Alkaline Earth Metals. The common outermost electronic configuration of the elements of this group is ns2.

Group 3 Starts with 5B(2s2 2p1). The general electronic configuration of the elements of this group is ns2 np1 . This group is also known as the Boron Family.

Group 4 Starts with 6C(2s2 2p2). The general electronic configuration is ns2 np2. This group is

also referred to as the carbon family.

Group 5 Starts with 7N(2S2 2P3).the general outermost electronic configuration of the elements

of this family ns2 np3. This group is also referred to as the Nitrogen Family.

Group 6 Starts with 8O(2S2 2p4) and is known as Oxygen Family. The general outermost electronic configuration of the elements of this family ns2 np4. The elements of this family are also known as Chalcogens.

Group 7 Contains 9F(2S2 2p5), 17Cl(3s2 3p5), 35Br(4s2 4p5), 53I(5s2 5p5) and 85At(6s2 6p5) The elements of this group are commonly known as Halogens.

Typical Elements
Elements of the third period are known as Typical elements, examples, 11Na, 12Mg, 13Al, 14Si, 15P, 16S and 17Cl Properties of all elements present in a particular group e.g. of group 1 resemble with the properties of 11Na and not with 3Li.

Bridge Elements
Elements of second period are known as Bridge Elements. Properties of the bridge elements resemble with the properties of the diagonal elements of the third period. For example, Li resembles Mg; Be resemble Mg; Be resembles Al; B resembles Si etc.

Note:- Noble gases are also grouped with representative P – block elements as they come at the end of each period.
The chemical and physical properties of the representative elements is determined by the number of electrons in the outer most shell called the Valence Shell. The number of valence electrons for groups 1 and 2 is the same as the group number for group 13 – 17, this number is obtained by subtracting 10 from the group number.

3. Transition Elements (d- block elements)
These are the elements of Groups 3 to 12 in the center of the period table. The elements in which the last electron enters the d sub-shell of the penultimate energy level are called d block elements. The outer most configuration of these elements is (n-1)d1 - 10 ns1 - 2 They are metals. They form colored ions and exhibit variable valency. However, Zn, Cd and Hg which too have (n – 1) d1 – 10 ns2 configuration in their outermost shell do not form colored ions and are not regarded as transition elements.

4. The Inner Transition elements ( f – Block Elements)
The rows of elements at the bottom of the periodic table are called the Lanthanide and  actinide series. These elements in which the last electron enters the f sub-shell of the antipenultimate  (third to the outermost shell) shell are called f block elements Their outer  electronic configuration is (n-2)f1 – 14 (n-1) d0-1 ns2. The differentiating electron is an felectron.
They are all metals. Within each series the properties of the elements are quite similar .

PERIODIC TRENDS IN PROPERTIES
1. VALENCE:-
An important chemical property of the elements exhibiting periodic trends is their Valence. It is defined as combining capacity of an element. It can also be defined it terms of valence electrons (electrons in the outermost shells). The valency is equal to number of valence electrons (or equal to 8 minus the number of valence electrons.

1. the valence of representative elements is usually equal to the number of electrons in the outermost orbitals and /or equal to eight minus the number of outermost electrons.

2. Transition elements do not exhibit any general trend. The reason for this that those elements have variable valencies due to availability of vacant d- subshells in them.

3. Inner transition elements also do not exhibit any general trend in the valency.

2. ATOMIC AND IONIC RADII
It is impossible to define the size of atoms as we know that atoms have no shop boundaries due to the delocalized picture of electron cloud. An estimate of he atomic size can be made by knowing the distance between the atoms in the combined state. 

There are three operational concepts of atomic radius.

a. If the bonding is covalent, the radius is called covalent radius.

b. If the bonding is ionic, the radius is called ionic radius.

c. If the two atoms are not bounded by a chemical bond (as in noble gases), the radius is called Vander Walls radius.

a. Covalent Radius : 

It is half of the distance between the nuclei of two like atoms bounded together by a single bond. For Example, the bond distance in hydrogen molecule (H2) is 74 pm and half of this distance is taken as the atomic radius of hydrogen. This radius is known as the covalent radius.

b. Ionic Radius:

It is the effective distance from the nucleus of an ion up to which it has its influence on its electron cloud.

c. van der Wall’s radius : 

It is one half of the distance between the nuclei of two adjacent atoms belonging to two neighboring molecules of an element in the solid state. The covalent radius is always smaller than the van der Wall’s radius because in the formation of chemical bond, the tow atoms have to come closer to each other. This is why the inert gases (where covalent radius is generally not possible) tend to have a larger size.

1. The size of atoms increases as we go down a column of periodic table. This increase is attributed to the increase in the number of shells around the nucleus.

2. The size of the atoms decrease as we go across the period from left to right except group 18 (Noble Gases). This decreases in the size is attributed to the increases in the nuclear charge and hence the attraction.

3. A positive ion is always smaller in size than the corresponding neutral atom.

4. A negative ion is always bigger in size than the corresponding neutral atom.

5. The size of ions increases as we go down a group provided that we are comparing ions of same charge.

6. Atoms or ions with the same electronic configurations are called as iso-electronic. If we consider a series of iso-electronic species (atoms or ions), the size decreases with the increasing atomic number. 

To illustrate the concept consider the radius of the following iso-electronic species, all having 10 electrons   N3- > O2- > F- > Ne > Na+ > Mg2+ > Al3+

Note that the successive increase in the values of Z/e ratio decreases the values of ionic (atomic ) radii.

3. IONIZATION ENERGY
The chemical nature of an element depends on the ability of its atoms to accept of donate

electrons. A quantitative measure of these tendencies is the Ionization Energy or the Electron

or the Electron Affinity.

The ionization Energy (IE) is defined as the energy required to remove an electron from an isolated gaseous atom (M) in its ground state.
M(g) + IE --------> M+(g) + e-

The ionization energy is expressed in units of kJ mol-1 or in ev / electron.

The energy required to remove the second electron from the same element is known as the second ionization energy. 

The second ionization energy is higher than that required for the removal of the firs electron because it is more difficult to remove an electron from a positively charged species than the second and so on. If the term ionization energy is not qualified (i.e., if the first, second and third is not motioned), it is taken as the first ionization energy.

M+(g) + second IE--------------> M+++(g) + e-

M++(g) + third IE--------------> M+++(g) + e-

FACTORS INFLUENCING IONIZATION ENERGY :
The ionization energy depends upon the following factors :

Nuclear Charge
Ionization energy increases with increase in the nuclear charge. With increase in the nuclear charge the force with which the electron is bound with the atom increases and hence it becomes more difficult to remove the outermost electron. For example the ionization energy of He is 567 kcal/mol, whereas that of H is only 314 kcal/mole. This increase in the ionization energy can be explained on the basis of the increase in the nuclear charge.

Atomic Size
With the increase in the atomic size, the force with which the electron is bound to the atom decreases and hence with increase in the size, the ionization energy decreases. For example the ionization energy decreases as we don down the group.

Screening Effect
The force of attraction between the valence electrons and the nucleus is greatly shielded by the presence of core electrons. With the increase in the electrons in the inner sub shells the force of attraction between the outermost electron and the nucleus decreases and hence the ionization energy decrease.

Penetration of Electrons
The ionization energy also depends upon the penetrating power of the electrons. For example, the penetration of s-electrons. The f-electrons has the least penetration. The ionization energy increases with increasing penetration.

Stability of the Electronic Configuration
The half and fully filled orbitals are most stable as compared to their neighbours and hence the ionization energy of the fully filled or half filled orbitals is higher as compared to their neighbors. For instance, he ionization energy for N is higher that for C and O.

PERIODIC TRENDS
1. It is observed that the ionization energy of an element strongly depends on its electronic configuration and thus show periodic variations. The maxima are found at the noble gases which have completely filled electron shells. 

The high ionization energies of the noble gases can be connected with their extremely low chemical reactivity. Similarly, the high reactivity of alkali metals is reflected I their.


2. In a group : First ionization energy decreases as we go down a group in the table. It measures the ease of removing an electron from the outer shell. As we go down a group, this shell is farther away from the nucleus. As a result nuclear attraction decreases. Though the positive charge of nucleus increases, its effect is weakened due to the shielding supplied by the inner shells to the outermost shell.


3. In a period: 

As we go across a period from left to right, the atomic size decreases. As the number of shells in a particular period remain the same and the additional electrons are being continuously introduced in the same shell, the nuclear charge increases and hence the outer electrons are greatly attracted to the nucleus. Hence it becomes difficult to remove them and consequently ionization energy increases. The figure below shows the first ionization energy of elements of the second period as a function of the atomic number Z.

a. IE (B) < IE (Be). Be has its 2s orbital fully filled whereas B has one unpaired 2p electron and it is easier to remove a lone electron rather than that from a paired orbit. Hence extra-stability of fully filled sub- shell is the cause of this irregularity.

b. IE (O) < IE (N). nitrogen has an exactly half filled outermost electronic configuration and hence is extra stable. Thus extra-stability of half filled sub-shells is the cause of irregularity.

4. Whether all the outer shell electrons are removed, the next I.E. is much greater than the previous value of I.E. For the same element. Note that first I.E for the same is 72.64 eV.


4. ELECTRON AFFINITY
The electron affinity is the amount of energy releases when an isolated gaseous atom accepts an electron to form a monovalent gaseous ion.

X (g) + e- X- (g) + Energy

Electron affinities can be positive or negative. When energy is released in the process of attachment of an electron to an atom, the electron affinity is taken as positive and if energy is absorbed, electron affinity is taken to be negative like in inert gases. Thus the magnitude of electron affinity measures the tightness with which the atom can hold the additional electron. The larger value of E.A reflects the greater tendency of an atom to accept the electron.

Electron affinity values are influenced by

a. size of the atom b. Nuclear charge c. Electronic configuration

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