Periodic Classification of Elements

Aufbau Principle

In a ground state of the atoms, the orbitals are filled in order of their increasing energies." i.e. an electron will initially occupy an orbital of lower energy level and when the lower energy level orbitals are occupied, then only they shall start occupying the higher energy level orbitals.

Period Length and Sublevels in the Periodic Table

Period	Number of Elements in Period	Sublevels in Order of Filling
1	2	1s
2	8	2s 2p
3	8	3s 3p
4	18	4s 3d 4p
5	18	5s 4d 5p
6	32	6s 4f 5d 6p
7	32	7s 5f 6d 7p

Atomic Radius of an Element

The linear distance between the centre of the nucleus and valence shell of an atom of the element is called atomic radius of that element. It is denoted by letter 'r' and is expressed in nm (10⁻⁹ m) or Å (10⁻¹⁰ m).

According to modern concepts, electrons are taken as diffused clouds, what we call orbitals. Therefore, atomic radius has no precise physical significance since the electron density does not suddenly fall to zero at a particular distance from the nucleus or better to say that the atom has no definite boundary. At best, however, we can explain atomic radius of an element as the distance between the centre of nucleus and the point in three dimentional space where electron density of its atom falls to zero (i.e. probability of finding an electron is maximum).
It is therefore at best be explained on the basis of the physical state and the nature of bonds of an atom of the element. So, we express the atomic radius as-
1. Covalent Radius
2. Van der Waal's Radius
3. Metallic Radius
4. Ionic Radius

Covalent Radius

It is one half of the distance between centres of nuclei of the two covalently bonded atoms of the same element in a molecule. The distance between the centres of nuclei of two H-atoms in H₂ molecule is 0.74Å and hence 0.37Å is the covalent radius of hydrogen. The covalent radius decreases with the increase in the value of bond order. Covalent radius of single, double and triple bonded carbon-carbon atoms are 0.154nm, 0.142nm and 0.120nm respectively.

Van der Waal's Radius

It is one half of the distance between centres of nuclei of the two adjacent atoms belonging to two neighbouring molecules of an element in the solid state. The two adjacent atoms belonging to two molecules of iodine in solid state are shown below-

Therefore, the covalent radius of iodine is 1.33Å. Van der Waal's radius is much larger than covalent radius of an element.

Metallic Radius

It is one half of the distance between centres of two adjacent kernels (atoms stripped of valence electrons) in the metallic close packed lattice.
The distance between centers of two adjacent kernels (M⁺ⁿ) in metallic closed pack lattice of Al is shown below-
Al -----> Al⁺³ + 3e
K²L⁸M³         K²L⁸

Hence, the metallic radius of Al is 0.43Å. The average bond order of metallic bond is generally less than one. Therefore, the metallic radius is longer than a covalent radius but shorter than Van der Waal's radius because bonding forces in metallic crystal lattice are much stronger than Vander Waal forces.

Ionic Radius

We know the atom is neutral because the number of electrons (negative charges) is equal to the number of protons (positive charges). When the number of electrons and protons are not equal, ions are formed from the atom.
(e = p) -----> (e ≠ p)
 Atom             Ion
Hence, there are two types of ions-cation and anions. The number of protons is always greater than the number of electrons in cations whereas the number of electrons is always greater than the number of protons in anions. If the ion has 'n' electrons more or less than its atom, then the ion has '− n' and + n charges.
    M − ne = M⁺ⁿ
 Atom      Cation
    M + ne = M⁻ⁿ
 Atom      Anion
So, we have cationic radius as well as anionic radius. The ionic radius is the distance between the centre of the nucleus of the ion and the point where electron density falls to zero. The cationic radius is always shorter than that of its atom, whereas anionic radius is always larger than that of its atom.

Classification of Elements

s-Block Elements

1. The elements of the periodic table in which the last electron enters in s-orbital, are called s-block elements.
2. s-orbital can accommodate a maximum of two electrons.
3. Their general formulae are ns^1-2, where n = (1 to 7)
4. IA group elements are known as alkali metals because they react with water to form alkali while II A group elements are known as alkaline earth metals because their oxides react with water to form alkali and these are found in the earth.
5. Total number of s-block elements are 14
6. All the elements are soft metals
7. They have low melting and boiling points
8. They are highly reactive
9. Most of them impart colours to the flame
10. They generally form ionic compounds
11. They are good conductors of heat and electricity
12. ₅₇Fr and ₈₈Ra are radioactive elements while H and He are gaseous elements.
13. Cs and Fr are liquid elements belonging to s-block.

p-Block Elements

1. The elements of the periodic table in which the last electron enter in the p-orbital, are called p-block elements
2. p-orbital can accommodate a maximum of six electrons. Therefore, p-block elements are divided into six groups which are III A to VII A and zero group
3. The general formula of p-block elements is ns²np^1-6, (where n = 2 to 6)
4. The zero group elements having general formula ns² and ns²np⁶
5. The total number of p-block elements in the periodic table is 30 (excluding He)
6. There are nine gaseous elements (Ne, Ar, Kr, Xe, Rn, F₂, Cl₂, O₂ and N₂) belonging to p-block. Gallium(Ga) and bromine (Br) are liquids
7. The compounds of p- block elements are mostly covalent in nature
8. They show variable oxidation states
9. In moving from left to right in a period, the non-metallic character of the elements increases
10. The reactivity of elements in a group generally decreases downwards
11. At the end of each period is a noble gas element with a closed valence shell ns² np⁶ configuration
12. Metallic character increases down the group

d-Block Elements

1. The elements of the periodic table in which the last electron enter in the d-orbital, called d-block elements
2. The d-block elements are placed in the groups named I B to VII B and VIII (does not have sub group)
3. In d-block elements the electron gets filled up in the d-orbital of the penultimate shell
4. d-block elements lie between s & p block elements. (Transition Elements)
5. The general formula of these elements is ns^1-2(n-1)d^1-10. Where n = 4 to 7
6. All of these elements are metals(Transition Metals)
7. Out of all the d-block elements, mercury is the only liquid element.
8. They are all metals with high melting and boiling points
9. The compounds of the elements are generally paramagnetic in nature
10. They mostly form coloured ions, exhibit variable valence (oxidation states)
11. They are used as catalysts

f-Block Elements

1. The element of the periodic table in which the last electron enter in the f-orbital, called f-block elements
2. The f-block elements are from atomic number 58 to 71 and from 90 to 103
3. The lanthanides occur in nature in low abundance and therefore, these are called rare earth elements
4. There are 28 f-block elements in the periodic table.
5. The elements from atomic number 58 to 71 are called lanthanides because they come after lanthanum(57). The elements from 90 to 103 are called actinides because they come after actinium (89)
6. All the actinide elements are radioactive
7. All the elements after atomic number 92 (i.e. ₉₂U) are transuranic elements
8. The general formula of these elements is (n−2)f^0-14(n−1)d^0-1)ns²
9. They are all metals. Within each series, the properties of the elements are quite similar
10. Most of the elements pf the actinoid series are radio-active in nature

Ionization Potential

The minimum amount of energy required to remove an electron from an isolated gaseous atom to produce a cataion is called ionization potential.
Since it is energy so expressed in KJ/mol or in ev/atom.(1ev = 97.5KJ)
                             ∆H_(S)
              M_(solid) ------> M_(gas)
                             I.P_I
               M_(gas) -----> M_(gas) ⁺ + e
                             I.P_II
               M_(gas) ⁺ -----> M_(gas) ⁺² + e
                             I.P_III
               M_(gas) ⁺² -----> M_(gas) ⁺³ + e
Experimental value of I.P is 2.18 x 10^-18 x Z*/n²J.
Or, I.P is directly proportional to Z*/n²
where, Z* is effective nuclear charge and n is orbit number
Greater the Z*, greater the I.P and larger the atomic radius lower is the I.P.
It is the property of metals. Smaller the value I.P., easier the ease of formation of cation.
The value of IP_I is smaller than that of IP_II and the value of IP_II is less than that of IP_III.
i.e. IP_I < IP_II < IP_III

Factors affecting the ionization potential

1. Nuclear Charge
2. Atomic Size
3. Penetrating effect of the electron
4. Screening effect of the inner shell electron And
5. Electronic configuration

Nuclear Charge

I.P increases with increase in nuclear charge because, with increase in nuclear charge the electron in outer shell are more tightly held by the nucleus and thus more energy required to remove an electron from the atom.

Atomic Size:

I.P decreases with increasing the atomic size because, the distance of outer electron from the nucleus increases with size and thus the force of attraction decreases and electron in the outer shell held with nucleus less firmly and less amount of energy is required to remove an electron from the atom.

Penetrating Effect

I.P increases as penetrating of the electron increases. Within the same shell the penetrating effect is in order--- s > p > d > f. If the penetrating effect of the electron is more, the electrons are closer to the nucleus and held firmly. Consequently more energy is required to remove an electron from the atom.

Screening Effect

I.P decreases when screening effect (repulsive force felt by the valence shell electrons from the electrons present in the inner shell) increases.

Electronic Configuration

Atoms having the half filled or full filled configurations are more stable and extra energy required to remove an electron.

Variation of Ionization Potential in Periodic Table

Electron Gain Enthalpy (Electron Affinity)

The minimum amount of energy released when an electron is added to an isolated gaseous atom so as to convert it into a negative ion.
During the addition of an electron, energy can either be released or absorbed.
Since it is energy so expressed in KJ/mol or in ev/atom.(1ev = 97.5KJ)
X_(s) → X_(g)
X_(g) + e → X_(g)⁻
Electron gain enthalpy is generally negative for all elements except few elements. For example, the electron gain enthalpy for halogens is highly negative because they can acquire the nearest noble gas configuration by accepting an extra electron.
In contrast, noble gases have large positive electron gain enthalpies because the extra electron has to be placed in the next higher energy level thereby producing highly unstable electronic configuration.
After the addition of one electron, the atom becomes negatively charged and the second electron is to be added to a negatively charged ion. But the addition of second electron is opposed by electrostatic repulsion and hence the energy has to be supplied for the addition of second electron. Thus the second electron gain enthalpy of an element is positive.
For example, when an electron is added to oxygen atom to form O^– ion, energy is released. But when another electron is added to O^– ion to form O^–2 ion, energy is absorbed to overcome the strong electrostatic repulsion between the negatively charged O^– ion and the second electron being added.
Electron Gain Enthalpy and Electron Affinity are related to each other by the given formula.
Electron gain enthalpy = electron affinity – 5/2 RT
where, R= universal gas constant and T= temperature in Kelvin scale

Factors affecting the Electron Gain Enthalpy:

1. Atomic size
2. Nuclear charge
3. Electronic Configuration

Atomic size

As the size of an atom increases, the distance between its nucleus and the incoming electron also increases and electron gain enthalpy becomes less negative.

Nuclear charge

With the increase in nuclear charge, force of attraction between the nucleus and the incoming electron increases and thus electron gain enthalpy becomes more negative.

Electronic Configuration

The atoms with symmetrical configuration (having fully filled or half filled orbitals in the same sub-shell) do not have any urge to take up extra electrons because their configuration will become unstable.
In that case the energy will be needed and electron gain enthalpy will be positive. For example, noble gas elements have positive electron gain enthalpies.

Variation of Electron Gain Enthalpy in Periodic Table

In period from left to right electron gain enthalpy increases while down the group it is decreases.

Exception in Electron Gain Enthalpy

In the case of Chlorine and Fluorine, Chlorine has a higher negative electron gain enthalpy value due to very small size of fluorine
In between Sulphur and Oxygen, Sulphur has a higher negative value than oxygen due to very small size of oxygen.

Electronegativity, Factors Affecting its value and Importance