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WHAT IS PERIODIC PROPERTY?

Periodic properties, in simple terms, are the predictable patterns or trends exhibited by elements as you move across a row (period) or down a column (group) in the periodic table. These trends arise due to the arrangement of electrons in atoms and how they interact with each other.

Elements display various periodic properties, including:
  1. Atomic and Ionic Sizes
  2. Ionization Energy
  3. Electron Affinity
  4. Electronegativity
  5. Metallic and Non-metallic Character

AFBAU'S PRINCIPLE

It states that electrons are filled into atomic orbitals in the increasing order of orbital energy level. According to the Aufbau principle, the available atomic orbitals with the lowest energy levels are occupied before those with higher energy levels. 
Afbau principle of filling electrons

HUND'S RULE

It states that orbitals of the same energy will be filled partially ( i.e. single electron filling before pairing). Because electron prefers to remain far as possible to acquire minimum energy, repulsion, and more stability
Hunds rule of electron filling

ORBIT VS ORBITALS

 

ORBIT

ORBITALS

Definition

A circular path is followed by an electron in the Bohr model of the atom.

A three-dimensional region within an atom where the probability of finding an electron is high.

SIZE

Orbits are well-defined and have fixed radii at specific energy levels.

Orbitals do not have fixed sizes or precise boundaries; their shape and size depend on the quantum numbers that define them.

CAPACITY

Orbits can accommodate only a limited number of electrons (2n2, where n is the principal quantum number).

Each orbital can accommodate a maximum of two electrons with opposite spins, following the Pauli exclusion principle. Multiple orbitals can exist within a given energy level.


EFFECTIVE NUCLEAR CHARGE

The actual charge felt by the valence electrons is called effective nuclear charge and the repulsive force felt by the valence shell electrons from the electrons present in the inner cells is called the shielding effect or screening effect. Therefore, the effective nuclear charge (Zeff) is given by the relation

Zeff= Total nuclear charge (Z) — screening constant (S) 

where screening constant (S) takes into account the screening effect of the electrons present in the inner shells. Obviously, the greater the number of electrons in inner shells, the larger will be the screening effect. As the screening effect increases, the effective nuclear charge decreases.

Penetration Effect

The ability of an electron to get close to the nucleus is called the penetration of the electron.
Therefore, for the same shell value (n) the penetrating power of an electron in orbitals (subshells) is given as under:
s>p>d>f
And for different values of the shell (n) and subshell (l), the penetrating power of an electron follows this trend:
Is > 2s > 2p > 3s > 3p > 4s > 3d > 4p > 5s > 4d > 5p > 6s > 4f
and the energy of an electron for each shell and subshell follows the order
Is < 3s < 3p < 4s < 3d < 4p

Periodic Trends Due to Penetration and Shielding

Periodic Trends Due to Penetration and Shielding:
  1. Effective Nuclear Charge (Zeff): The effective nuclear charge increases from left to right and increases from top to bottom in the periodic table.
  2. Atomic Radius: The atomic radius decreases from left to right in a period, and increases from top to bottom in a group.
  3. Ionization Energies: The ionization energies increase from left to right, and decrease from top to bottom.
  4. Electronegativity: The electronegativity of the elements is highest near flourine. In general, it increases from left to right in a period and decreases from top to bottom in a group.

GENERAL ELECTRONIC CONFIGURATIONS:

S BLOCK:  ns1-2
P BLOCK: ns1-2 np1-6 
D BLOCK:  (n-1)d0-10 ns1-2

ATOMIC RADIUS

Atomic radius may be defined as the distance from the center of the nucleus to the
point upto which the density of the electron cloud (i.e. probability of finding the electron)
is maximum.

Types of Atomic Radius

There are three operational concepts that have been widely used. These are:
  1. Covalent radii
  2. Van der Waals' radii
  3. Ionic radii.

Covalent Radii

It is defined as one-half of the distance between the nuclei of two covalently of the same element in a molecule. Thus, for a homonuclear diatomic molecule,
r covalent= 1/2 * [Internuclear distance between two bonded atoms]
covalent radii (Types of atomic Radius)

Vander Waal's Radii

It is defined as one-half of the distance between the nuclei of two non-bonded isolated atoms or two adjacent atoms belonging to two neighboring molecules of an element in the solid state.

Ionic and Crystal Radii

It is defined as the effective distance from the nucleus of the ion up to which it has an influence in the ionic bond.

Periodic Trends in Ionic Radii

  1. Variation of ionic radii in a period. For an isoelectronic series (ions having the same number of electrons). The ionic radius decreases as the nuclear charge increase.
  2. Variation of ionic radii in a group. As we go down a group, the ionic radii increase with the increase in atomic number like atomic radius.

IONIZATION ENERGY

The ionization energy of an element is defined as the amount of energy required to remove the most loosely bound electron from an isolated gaseous atom.
M (g) + Ionization Energy —> M+ (g) + e- (g)
Ionization energy is also called ionization potential because it is measured as the amount of potential. It is expressed in terms of either kcal/mol or kJ/moI or electron volt/atom.

FACTORS AFFECTING IONIZATION ENERGY

The magnitude of ionization energy for an atom depends upon the following factors :
  • Size of the atom. The ionization energy depends upon the distance between the electron and the nucleus, i.e. the size of an atom. As the size of the atom increases, the outermost electrons are less strongly attracted by the nucleus because the force of attraction is inversely proportional to the square of the distance between the charged particles. As a result, it becomes easier to remove the electron, and, therefore, ionization energy would tend to decrease with the increase in the size of the atom.
  • Charge on the nucleus. The attractive force between the nucleus and the electrons increases with an increase in nuclear charge. This is because the force of attraction is directly proportional to the product of charges on the nucleus and that on the electrons. Therefore, with the increase in nuclear charge, it becomes more difficult to remove an electron, and ionization energy increases.
  • Screening effect of the inner electrons. In multi—electron atoms, the outermost electrons are shielded or screened from the nucleus by the inner electrons. This is the screening effect. As a result of this, the outermost electron will not feel the full charge of the nucleus. The actual charge felt by an electron is an effective nuclear charge. An effective nuclear charge will decrease thus, ionization energy will decrease.
  • Penetration effect of electrons. It is well known that in the case of multi-electron atoms if the penetration of the electron is more (s>p>d>f), it will be closer to the nucleus and will be held firmly. Consequently, ionization energy will be high. This means that ionization energy increases with the increase in the penetration power of the electrons.
  • Electronic arrangement or configuration. The ionization energy also depends upon the electronic configuration of the atom. It has been observed that certain electronic configurations are more stable than others. For example, half-filled and completely filled shells have extra stability associated with them. Consequently, it is difficult to remove electrons from these stable configurations and ionization energy is high. for example, noble gases have the most stable electronic configurations (ns2 np6) in each period and have the highest ionization energies.

VARIATION ALONG PERIOD

  1. On moving across a period from left to right, the nuclear charge increases.
  2. The atomic size decreases along the period but the energy level remains the same.

VARIATION DOWN A GROUP

  1. In going from top to bottom in a group, the nuclear charge increases.
  2. There is a gradual increase in atomic size due to an additional main shell (n).
  3. There is an increase in the shielding effect on the outermost electron due to an increase in the number of inner electrons.

SUCCESSIVE IONIZATION ENERGY

The energies required to remove subsequent electrons from the atom in the gaseous state are known as successive ionization energies. The term first, second, third, and more ionization energy refers to the removal of first, second, third, or more electrons respectively.

ELECTRON AFFINITY

The amount of energy released when an electron is added to an isolated gaseous
atom is called electron affinity.
When an electron is removed from an atom energy is required for the process of removal
of electron but energy is released when an electron is added to a neutral atom.
The process may be expressed as:
Cl (g) + e¯ (g) —> Cl-(g) + Energy (electron affinity)

FACTORS AFFECTIVE ELECTRON AFFINITY

There are many factors that govern the electron affinity but the following are some
important factors on which it mostly depends :
  1. Nuclear charge. The electron affinity increases as the nuclear charge increases. This is due to the greater attraction for the incoming electron if the nuclear charge is high.
  2. Size of the atom. With the increase in the size of the atom, the distance between the nucleus and the incoming electron increases, and this results in lesser attraction. Consequently, the electron affinity value will decrease.
  3. Electronic configuration. The element having stable electronic configurations of half and completely-filled valence subshells show a very small tendency to accept additional electrons and thus, electron affinities are low or almost zero in certain.

Why Electron affinity of fluorine is unexpectedly less than that of chlorine?

The low value of the F atom is due to the very small size of the F atom. As a consequence of its small size, they are strong inter—electronic repulsions in the relatively compact 2p subshell of fluorine and thus the incoming electron does not feel much attraction. As a result, its electron affinity value is small. On the other hand, the electron affinity of the Cl atom is larger than the F atom where the electron is added to a relatively large 3p—orbital which can easily accommodate the additional electron.

ELECTRONEGATIVITY

The tendency of an atom to attract the shared electron pair in a molecule towards itself is called electronegativity.
Different atoms have different tendencies to attract the bonded electron pair toward themselves. The greater the ability of an atom to attract electrons in a bond, the larger the value of its electronegativity.

FACTORS EFFECTING ELECTRONEGATIVITY