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Periodicity in Properties

According to the Modern Periodic Law, the properties of elements are periodic function of atomic numbers.
From the discussion of the periodic table, it is evident that those properties which depend upon the electron configuration of an atom will vary periodically with atomic number. On the other hand, those properties which depend upon the total number of electrons will show no such variations. Some of the more common properties which depend upon electronic configurations are:

Refer to the following video for periodicity in properties

Atomic Radius  or Atomic size

The term radius of an atom means the distance from the centre of the nucleus to the outermost shell of electrons. It is obviously impossible to isolate an individual atom and measure its radius; also the radius is altered by the nature of the other atom with which it is linked. Atomic radii are calculated from interatomic distances (i.e. the distance between the centres of two adjacent nuclei), which may be determined by X-ray and spectroscopic studies.

Note that the size of an atom is not constant; it depends on the chemical bonding of the atom, which in turn influences the electron cloud. The radius of a neutral atom is called its atomic or covalent radius. The sum of the radii (r1 + r2) of two bonded atoms equals the bond length.

Period

Group

IA

IIA

IIIA

IVA

VA

VIA

VIIA

Zero

1.

H

0.37

 

 

 

 

 

 

He

0.93

2.

Li

1.34

Be

0.90

B

0.82

C

0.77

N

0.73

O

0.74

F

0.72

Ne

1.31

3.

Na

1.54

Mg

1.30

Al

1.18

Si

1.11

P

1.06

S

1.02

Cl

0.99

Ar

1.74

Within a given period atomic radius decreases from left to right. This is due to the effect of increase in nuclear charge while the electrons are being added to the same shell. As the nuclear charge increases electrons are drawn closer to the nucleus. (Note also that the added electron repels its neighbouring electrons and tends to expand the electron cloud. But in a period when the electron is added to the same shell, the contracting effect of the increased nuclear charge exceeds the expanding effect of the increased electron repulsions).

Van der Waals’ radii:

It is one-half the distance between the nuclei of two adjacent atoms belonging to two neighbouring molecules in the solid state. Hence van der Waals’ radius is greater than covalent radius. We consider only the van der Waals’ radii of noble gases. Hence they are the largest in a period.

Ionic radii:

Ionic radius is the effective distance from the nucleus for an ion upto which its influence is felt on the electron cloud. The radius of an atom changes appreciably during loss or gain of electrons. When an atom loses an electron and forms a cation, the number of electrons decreases while the nuclear charge remains the same. The remaining electrons are drawn more towards the nucleus with the result that the cation is much smaller than the neutral atom.

For example, the atomic radius of Na is 1.54Å while the ionic radius of Na+ is 0.95Å.

Atom [M]

Li

Na

K

Rb

Cs

Atomic radius (Å)

1.34

1.54

1.96

2.11

2.25

Ionic radius (Å) [M]

0.60

0.95

1.33

1.48

1.69

On the other hand, when an anion is formed by the addition of an electron to a neutral atom there is increased electron-electron repulsion and the effective nuclear charge is reduced. So the anion is always larger than the neutral atom. e.g., the atomic radius of Cl is 0.99Å while the ionic radius of Cl is 1.81 Å.

Atom [X]

F

Cl

Br

I

Atomic radius (Å)

0.72

0.99

1.14

1.33

Ionic radius (Å) [M]

1.36

1.81

1.95

2.16

Across a period there is a decrease in ionic radius, while down a group there is an increase in ionic radius. This is with reference to the same type of ion (anion or cation).

In short we can say that..

  • Atomic radii decrease along the period.

  • Nuclear charge increases progressively by one unit on moving from left to  right across the period. As a result, the electron cloud is pulled closer to the nucleus by the increased effective nuclear charge, which causes decrease  n atomic size.

  • Atomic radii increase from top to bottom within a group. This increases is due to increase in the number of shells

Ionization Energy

  • Ionization energy is the amount of energy required to remove one electron electron from the outermost orbital of an isolated gaseous atom in its ground state

  • Ionization energy decreases with the increase in atomic size and increases with the increase in nuclear charge.

  • Ionization energy decreases with the increase in the number of inner electrons and increases with the increase in penetration power of electrons.

  • Atom which have a more stable configuration i.e. fully filled or half filled orbitals, has high value of enthalpy/energy.

  • Ionization energy increases with the increase in atomic number as we move from left to right across the period.

  •  Ionization energy decreases regularly with the increase in atomic number as we move down the group.

Electron Affinity/Electron Gain Enthalpy

  • Electron affinity is the enthalpy change taking place when an isolated gaseous atom accepts an electron to form a monovalent gaseous anion

  • Larger the value of electron gain enthalpy, greater is the tendency of an atom to accept electron.

  • Greater the magnitude of nuclear charge, larger will be the negative value of electron gain enthalpy.

  • Larger the size of the atom, smaller will be the negative value of electron gain enthalpy.

  • More stable the electronic configuration of the atom, more positive will be the value of its electron gain enthalpy.

  • Electron gain enthalpy become negative as we move from left to right  along the period.

  • Electron gain enthalpy becomes less negative down the group.

Electronegativity

  • It is the tendency of an atom in a molecule to attract the shared pair of electrons towards itself.

  • Greater the effective nuclear charge, greater is the electronegativity.

  • Smaller the atomic radius, greater is the electronegativity.

  • Electronegativity increases on moving from left to right in a period due to decrease in atomic radii and increase in effective nuclear charge.

  • Electronegativity decreases on moving down a group due to increase in atomic radii because of increased number of shells


Valency

  • It is the number of univalent atoms which can combine with an atom of the given element.

  • Valency is given by the number of electrons in outermost shell.

  • If the number of valence electrons ≤4: valency = number of valence electrons

  • If the number of valence electrons >4: valency = (8 - number of valence electrons)

  • Many elements exhibit variable valence (particularly transition elements and actinoids).

  • Valency increases in period from 1 to 4 and then decreases from 4 to zero on moving from left to right.

  • There is  no change in the valency of elements on moving down a group. All elements belonging to a particular group exhibit same valency.

Metallic Characterof an Element

  • Metals have tendency to lose the electrons.

  • Larger the size of atom, more easily it would release the electron and thus would show high metallic character.

  • Metallic character increases down the group due to increase in size of atom.

  • Metallic character decreases along the period due to decrease in atomic radii.

  • Non-metallic character is directly related to electronegativity and metallic character is inversely related to electronegativity.

  • Non-metallic character increases (and metallic character decreases) down the group.

  • Non-metallic character decreases (and metallic character increases) along the period.

  • This is why all the metals are present on left while non-metals are present on right side of the periodic table.

Solved Problems

Question 1:

Nitrogen has higher ionisation energy than oxygen – explain.

Solution:

In case of nitrogen, we see that it has a stable half filled shell 1s22s22p3. On the otherhand oxygen has neither a half filed nor a completely filled shell (1s22p22p4). Consequently N must have higher I.P. value than oxygen. However, although the first ionisation potential of N is greater than that of O, the second I.P. of N is much less than that of oxygen. This is because, after the removal of 1 electron from oxygen atom it attains stable 2p3 configuration (half filled shell).

_______________________________________________________

Question 2:

Why electron affinity of nitrogen is less than that of carbon?

Solution:

It is due to the comparatively stable half filled configuration (np3) of nitrogen and phosphorus and the tendency to acquire the stable np3 configuration by the gain of one electron in carbon and silicon (np)2.

_______________________________________________________

Question 3:

Generally second electron affinity is endothermic – explain

Solution:

When the anion A forms a bi negative ion A– –, it will not accept the second electron easily due to anion-electron repulsion. Hence work is to be done on the system to force the electron to enter inside the electronic structure. Consequently the second electron affinity must be endothermic (DH is positive).

Question 1: Lowest ionisation potential in a period is shown by

(a) Alkali metals

(b) Halogen

(c) Transition Elements

(d) Alkaline Earth Metals

Question 2: The second transition series contains the elements from

(a) K to Zn

(b) Se to Zn

(c) Y to Cd

(d) Sc to Hg

Question 3: Which has lowest ionisation potential

(a) Li

(b) Cs

(c) Cl

(d) I

Question 4: The outermost electronic configuration of the most electronegative element is

(a) ns2np3

(b) ns2np4

(c) ns2np5 

(d) ns2np6

Question 5: In the sixth period, the orbitals being filled are

(a) 5s, 5p, 5d 

(b) 6s, 6p, 6d, 6f

(c) 6s, 5f, 6d, 6p

(d) 4s, 4f, 5d, 6p

Q.1

Q.2

Q.3

Q.4

Q.5

a

c

b

c

c

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