CBSE Class 11 Chemistry Notes For Chapter 3 Periodic Trends In Properties Of The Elements
The properties of elements can be divided into two categories:
- Properties of individual atoms: The properties such as atomic and ionic radii, ionisation energy, electron affinity, electro¬negativity and valency are the properties of the individual atoms & are directly related to their electronic configurations.
- Properties of groups of atoms: The properties such as melting
point, boiling point, heat of fusion, heat of vaporisation, atomic volume, density etc. are the bulk properties i.e., the properties ofa collection of atoms and are related to their electronic configurations indirectly.
All these properties which are directly or indirectly related to the electronic configurations of the elements are called atomic properties. Since electronic configurations of the elements are periodic functions of their atomic numbers, these atomic properties are also periodic functions of atomic numbers of the elements. Thus, atomic properties are also called periodic properties.
Remember that, when we descendin a group, the chemical properties ofthe elements remain more or less the same due to same valence shell configuration, but there is a gradual change in physical properties due to gradual change in the size ofthe atoms owing to addition ofnewelectronic shells.
Atomic size or atomic radius
If an atom is assumed to be a sphere, the atomic size is given by the radius ofthe sphere, called the atomic radius.
Atomic size or atomic radius Definition:
The distance from the centre of the nucleus to the outermost shell containing the electrons is called the atomic radius
Difficulties in precise measurement of atomic radius:
- It is not possible to isolate a single atom for the measurement of its radius.
- The electron cloud surrounding the nucleus does not have a sharp boundary as the probability of finding an electron can neverbe zero even at alarge distance from the nucleus.
- The magnitude ofatomic radius changes from one bonded state to another.
Types of atomic radius:
As already mentioned, the size of an atom varies from one environment to another. Therefore, several kinds of atomic radii have been defined.
These are
- Covalent radius
- Metallic radius
- Van der
- Waals radius.
1. Covalent radius:
It is defined as one-half of the distance between the centres of two atoms of the same element bonded by a single covalent bond
Thus for homonuclear diatomicmolecules, covalent radius (r) = ½ internuclear distance.
Example: In H2 molecule, the internuclear distance is = 0.74 A
= 74pm. So, the covalent radius of hydrogen atom = ½ × 0.74
= 0.37 Å= 37pm.
1Å = 10-10m, 1 pm = 10-12m,
2. Metallic radius:
It is defined as one-half of the internuclear distance between two adjacent atoms in a metallic lattice.
Example:
The distance between two adjacent copper atoms in solid copper is 2.56 Å(determined by X-ray diffraction).
Hence, the metallic radius of copper is 1.28 Å. Similarly, the metallic radii of sodium and potassium have been determined as 1.86 Å and 2.31 Å respectively.
Note: Covalentradii ofNa and K are 1.54 Å and 2.03 Å
A metallic radius is always greater than a covalent radius.
3. Van der Waals radius:
It is defined as one-half of the distance between the nuclei of two non-bonded neighbouring atoms belonging to two adjacent molecules of an element in the solid state.
Example:
The distance between nuclei of two adjacent Cl -atoms of two adjacent chlorine molecules in solid states is 3.6 Å. So, the van der Waals radius of Cl-atom is
⇒ \(\frac{3.6}{2}\)= 1.8 A.
Since the van der Waals force of attraction is weak even in the solid state, the internuclear distances between the atoms of two adjacent molecules held by van der Waals forces are much larger than those between covalently bonded atoms (which involve mutual overlap of atomic orbitals).
Van der Waals radii are always greater than covalent radii. For example, the van der Waals radius of chlorine is 1.80 Å while its covalent radius is 0.99 Å
Sequenceofthree types of atomic radii: van der Waals radius >Metallic radius > Covalentradius
Variation of atomic radii or sizes in the periodic table
1. Variation of atomic radii or sizes across a period:
While moving from left to right across a period in the periodic table, atomic radii or atomic sizes progressively decrease.
Atomic radii across a period Explanation:
The principal quantum shell remains unchanged in the same period. So, the differentiating electrons enter the same shellbutdue to an increase in the number of protons, the positive charge of the nucleus also increases. So, attractive force of the nucleus for electrons in the outermost shell also increases. Hence, the atomic sizes, rather than the atomic radii of elements in the same period, gradually decrease with an increase in atomic number. In any period, atomic size of the element of group-1A is maximum and that of the halogen of group VILAisminimum.
Variation of atomic [covalent] radii of the elements of the third period (n = 3]:
2. Variation of atomic radii or size down a group:
On moving down along any group of the periodic table, the atomic sizes rather than the atomic radii of the elements increase remarkably.
Atomic radii or size down a group Explanation:
On moving down a group, a new electronic shell is added to each succeeding element, though the number of electrons in the outermost shell remains the same. This tends to increase the atomic size. At the same time, there is an increase in nuclear charge with an increase in atomic number. This tends to decrease the size.
However, the effect of the increased nuclear charge is partly reduced by the shielding effect of the inner electronic shells.In practice,it is found that the effect of the addition of a new electronic shell is so large that it outweighs the contractive effect of the increased nuclear charge. Hence, there occurs a gradual increase in atomic radii on moving down a group in the periodic table.
3. Variationofsizein agroupforheavierelements:
On moving down a group, the relative rate of increase of atomic radio is slow for heavier elements. So, the differences in size for heavier species such as Cs (6s¹) and Fr(7s¹) of group-1 or Ba(6s²) and Ra(7s²) of group-2 are very small. This is due to the presence of electrons in the inner d and f-orbitals having a poor screening effect. D and f-electrons do not screen the outer electrons effectively from the pull of the nucleus. There is only a small increase in size in spite of the addition of a new electronic shell.
Example: Na(1.54 Å), K(2.03Å), Rb(2.16 Å), Cs(2.35Å) etc
In the case of transition elements, such a decrease in the atomic size is relatively less. Here the differentiating electron instead of entering the outermost orbit, goes to the penultimate (n- f )d -subshell which is closer to the nucleus. However, due to the poor shielding effect of the additional d -electrons, there is a small increase in effective nuclear charge and hence, the decrease in the size with an increase in the atomic number is relatively small.
In the case of lanthanide elements, the differentiating electrons enter the (n – 2)f-orbital having a very poor shielding effect. So,In such cases, the increase in the effective nuclear charge is greater than that in the case of transition elements. Thus, the decrease in atomic sizes of the lanthanides is much more regular and distinct as compared to the transition elements. So, the difference in the extent of the decreased atomic sizes of these elements is due to their relative inability to screen the outermost electrons from attraction ofthe nucleus.
Covalent radii (pm] of representative element:
Variation of atomic radii with atomic numbers across the second period:
Variation of atomic radii with atomic numbers for group-1 metals:
Screening effect or shielding effect:
In multi-electron atoms, the electrons present in the inner shells shield the electrons in the valence shell from the pull of the nucleus. It means that the electrons of the inner shells act as a screen between the nucleus and the electrons in the valence shell. This is known as the screening effect shielding effect.
Screening effect or shielding Definition:
The ability of the inner electronic shells to shield the outer electrons from the attraction of the nucleus is called the screening effect or shielding effect.
Magnitude of the screening effect of electrons belonging to different subshells follows the sequence: s>p> d>f.
Important points regarding atomic (covalent) radius:
- The alkali metals occupying positions at the extreme left side ofthe periodic table have the largest size in period
- The halogens occupying positions at the right side of the periodic table have the smallest size in each period.
- The noble gases present at the extreme right of the periodic table have larger atomic radii than those of the preceding halogens because van der Waals radii (but not covalent radii) are taken into consideration for noble gases.
- In transition series {d -block elements), there is only a small decrease in size with successive increases in atomic number because the differentiating electrons enter into (n-1)d subshell, which partially shields the increased nuclear charge acting on the valence electrons. This is known as d contraction.
- For inner-transition series ( /-block elements), the decrease in atomic radii wide increase in atomic number is relatively greater and more regular as compared to the transition series elements. This is so because the differentiating electrons enter into (n- 2)/ -subshell having a very poor shielding effect.
In a group of representative elements, there is a continuous increase in atomic radii with an increase in atomic number.
On going down a group of transition elements, there is an increase in size from first member to second member as expected, but for higher members, the increase in size is quite small. This is due to lanthanide contraction. For example, Cu(1.28 Å), Ag(1.44 Å), Au (1.44 Å)etc