(i) s-block elements: ns^1-2.where n = 2 -7.
(ii) p-block elements: ns² np^1-6 where n = 2- 6.
(iii) d-block elements: (n -1) d^1-10 ns^0-2where n = 4 - 7.
(iv) f-block elements: (n - 2) f^0-14 (n -1) d^0-1ns² where n = 6 - 7.

Ionization enthalpy, among other things, depends upon the electronic configuration (number of electrons) and nuclear charge (number of protons). Since isotopes of an element have the same electronic configuration and same nuclear charge, they have same ionization enthalpy.

Electron gain enthalpy refers to tendency of an isolated gaseous atom to accept an additional electron to form a negative ion. Whereas electronegativity refers to tendency of the atom of an element to attract shared pair of electrons towards it in a covalent bond.

A cation is smaller than the parent atom because it has fewer electrons while its nuclear charge remains the same. The size of anion will be larger than that of parent atom because the addition of one or more electrons would result in increased repulsion among the electrons and a decrease in effective nuclear charge.

Within a group atomic radius increases down the group.
Reason: This is due to continuous increases in the number of electronic shells or orbit numbers in the structure of atoms of the elements down a group.
Atomic radius: From left to right across a period atomic radii generally decreases due to increase in effective nuclear charge from left to right across a period.

The elements in a group have same valence shell electronic configuration and hence have similar physical and chemical properties.

The sixth period corresponds to sixth shell. The orbitals present in this shell are 6s, 4f, 5p and 6d. The maximum number of electrons which can be present in these sub-shell is 2 + 14 + 6 + 10 = 32. Since the number of elements in a period corresponds to the number of electrons in the shells, the sixth period should have a maximum of 32 elements.

The valency of an element may be defined as the combining capacities of elements. The electrons present in the outermost shell are called valence electrons and these electrons determine the valency of the atom.

As we move down from top to bottom in a group, electronegativity decreases due to increased atomic radius. Fluorine has the highest value of electronegativity among all the elements.

As we move from left to right in a period, electronegativity increases. This is due to the following reasons:
(a) Nuclear charge increases in a period
(b) Atomic size decrease in a period
Halogens have the highest value of electronegativity in their respective periods.

(i) Effective nuclear charge: As the nuclear charge increases, electronegativity also increases along the periods.
(ii) Atomic radius: The atoms in smaller size will have larger electronegativity.

When an electron is added to oxygen, the added electron goes to the 'L' shell (n = 2). As the 'L' shell possess smaller region of space, the added electron feels significant repulsion from the other electrons present in this level.
E.A of O=-141 KJ mole^-1
E.A of S = -200 KJ mole^-1

When an electron is added to F, the added electron goes to the L shell (n = 2). As the 'L' shell possess smaller region of space, the added electron feels significant repulsion from the other electrons present in this level.
E.A of F = -328 KJ mole^-1 .
E.A of CI = -349 KJ mole^-1.

The electron gain enthalpy of an element is the amount of energy released when an electron is added to the neutral gaseous atom
A + electron  \(\rightarrow\)  A^- + energy (E.A)
Unit of electron affinity is KJ mole^-1.

₇N 1s² 2s² 2p³ (or) Nitrogen has exactly half filled valence electrons, which requires high I.E₁.
₈O 1s² 2s² 2p⁴ (or) Oxygen has incomplete valence shell electrons, which requires comparatively less I.E₁ N > I.E₁O

Be (Z = 4) 1s² 2s². It has completely filled valence electrons, which requires high IE₁.
B (Z = 5) 1s² 2s².2p¹. It has incompletely filled valence electrons, which requires comparatively less IE₁.Hence I.E₁ Be > I.E₁B.

The energy required to remove the most loosely held electron from an isolated gaseous atom is called ionization energy.
M_(g)+ energy ⟶ M⁺_(g) + electron
The unit of ionization energy is KJ mole^-1

When an atom gain one or more electrons it forms anion. During the formation of anion, the number of orbital electrons become greater than the nuclear charge. Hence, the electrons are not strongly attracted by the lesser number of nuclear charges. Hence anionic radius is higher than the corresponding neutral atom.

(i) When an neutral atom lose one or more electrons it forms cation.
Na ⟶ Na⁺ + e^-
(ii) The radius of this cation (r_Na⁺)is decreased than its parent atom (r_Na)
(iii) When an atom is charged to cation, the number of nuclear charge becomes greater than the number of orbital electrons. Hence the remaining electrons are more strongly attracted by the nucleus. Hence the cationic radius is smaller than its corresponding neutral atom.

The ionic radius of an ion is the distance between the centre of the ion and the outermost point of its electron cloud.

The term periodicity of properties indicates that the elements with similar properties reappear at certain regular intervals of atomic number in the periodic table.
Example:
(i) Atomic radius 
(ii) Ionisation energy
(iii) Electron affinity
iv) Electronegativity.

(i) Some elements in the periodic table show properties that are characteristic of both metals and non-metals. These elements are called semi-metals or metalloids.
(ii) Example: Silicon, germanium, arsenic, antimony and tellurium.

(i) The elements, whose last electron enters into the f-orbital of the ante-penultimate shell (n-2) are having similar properties are called f-block elements. In these elements (n-2) f orbitals are being filled progressively.
(ii) The two rows of elements placed at the bottom of the periodic table.They are lanthanides and actinides.

d-block elements form a bridge between the chemically active metals of s-block elements and the less active elements of groups of 13^th and 14^th and thus take their familiar name transition elements.

(i) Zinc,cadmium and mercury are metals with low melting points.This is because they have an especially stable electronic configuration.
(ii) Mercury is so poor at forming metallic bonds that it is liquid at room temperature.
(iii) Zinc and cadmium are soft metals that oxidize to the +2 oxidation states.

(i) Zn, Cd and Hg are having completely filled d-orbitals (d¹⁰ electronic configuration).
(ii) They do not have partially filled d-orbitals like other transition elements. So they do not show much of the transition elements properties.

(i) The elements whose last electron enters into the d-orbital of the penultimate shell (n-l) are having similar properties and called as d-block elements.
(ii) The groups of 3 to 12 in the centre of the periodic table belongs to d-block.

(i) Group 16 (chalcogens) and Group 17 (halogens) are interested to add two or one electrons respectively to attain stable noble gas configuration.
(ii) Because of this interest these elements have highly negative electron gain enthalpies.

Noble gases having closed valence shell configuration as ns² np⁶. The valence shell orbitals of noble gases are completely filled by electrons and it is very difficult to alter this stable arrangement by the addition or removal of electrons. Because of these reasons noble gases do not show much of chemical reactivity.

(i) The elements whose last electron enters into the p-orbital of the outermost shell are having similar properties and thus form a group. The 'np' orbital of these elements is being progressively filled. Hence, these elements are named as p-block elements.
(ii) The groups of 13^th to 18^th in the periodic table belongs to p-block.

(i) Alkaline earth metals readily lose their outermost electrons to form +2 ion.
(ii) As we go down the group, their metallic character and reactivity are increased.

1. Alkali metals readily lose their outermost electron to form +1 ion.
2. Alkali metals are soft metals with low melting and boiling points.

Elements of 1^st group are called alkali metals.Their electronic configuration is ns¹ Elements of 2^nd group are called alkaline earth metals. Their electronic configuration is ns².

A vertical column of the periodic table is called a group. A group consists of a series of elements having similar configuration of the outermost shell. There are 18 groups in the periodic table.It may be noted that the elements belonging to the same group are said to constitute a family. For example, elements of group 17 are called halogen family.

Z = 101 IUPAC name: Mendelevium
Z = 102 IUPAC name: Nobelium
Z = 109 IUPAC name: Meitnerium
Z = 110 IUPAC name: Darmstadtium

The long form of periodic table need clarification about the following:
(i) Position of hydrogen is not defined till now.
(ii) Lanthanides and actinides still find place in the bottom of the table

Henry Moseley studied the X-ray spectra of several elements and determined their atomic numbers (Z). He noticed that the frequencies of X-ray emitted from the elements concerned could be correlated by the equation \(\sqrt { \upsilon } =a(Z-b)\)
Where, ሀ = Frequency of the X-ray emitted by the element.
a and b = Constants and have same values for all the elements.
Z = Atomic number of the element.

This law states that "The physical and chemical properties of elements are a periodic function of their atomic weights."

The Law of octaves states that, "when elements are arranged in the order of increasing atomic weights, the properties of the eighth element are a repetition of the properties of the first element".

In this system, elements that differed from each other in atomic weight by 16 or multiples of 16 fell very nearly on the same vertical line. Elements lying directly under each other showed a definite similarity. This was the first periodic law.

Johann Dobereiner noted that elements with similar properties occur in groups of three which he called triads. It was seen that invariably, the atomic weight of the middle number of the triad was nearly equal to the arithmetic mean of the weights of the other two numbers of the triad.
For e.g., 

Cl(g) + e^- ⟶Cl^-(g)    AH = 348 kJ mol^-1
For one mole (35.5g) 348 kJ is released.
\(\therefore \text { For } 17.5 \mathrm{~g} \text { chlorine, } \frac{348 \mathrm{~kJ}}{35.5 \mathrm{~\not g}} \times 17.75 \mathrm{~\not g} \text { energy leased. }\)
∴ The amount of energy released =\(\frac { 348 }{ 2 } \) =174 kJ

Noble gases: Ioniation energy ranging from 2372 KJmol^-1 to 1037 kJ mol^-1.
For element X, the IE₁ value is in the range of noble gas, moreover for this element both IE₁ and IE₂ are higher and hence X is the noble gas.
For Y, the first ionisation energy is low and second ionisation energy is very high and hence Y is most reactive metal. For Z, both IE₁ and IE₂ are higher and hence it is least reactive.

Electronic Configuration: (n - 1)d² ns²
for n = 5, the electronic configuration is,
1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 4d² 5s²
Atomic number: 40
4^th group 5^th period (d block element) = Zirconium

Atomic number: 120
IUPAC temporary symbol: Unbinilium
IUPAC temporary symbol: Ubn
Possible electronic configuration :[Og] 8s²

The main basic difference between Mendeleev's periodic table and modem periodic table is that first one is constructed on the basis of atomic weight and the later is constructed on the basis of atomic number.

(i) \(\sigma \) = 1 x 0.35  + 18 x 0.85 + 10 x 1 = 25.65
(ii) \(\sigma \) = 9 x 0.35  +18 x 1.0 = 21.15

The last member of the sixth period is the element Radon with atomic number 86. Therefore, the element with atomic number 114 belongs to the seventh period. In the seventh period, 7s, 5f, 6d and 6p orbitals are filled. 7s, 5f and 6d orbitals filling corresponds to 26 elements i.e., and member will be with atomic number 112. It means the element in question is the second 6p element. i.e. it belongs to group 10 + 4 = 14

Mendaleev's periodic law based on the properties of elements are periodic functions of their atomic weights while the modem periodic law based on the properties of elements are periodic function of atomic numbers.

Na₂O reacts with water and forms a base i.e, sodium hydroxide.
Na₂O + H₂O \(\rightarrow\)2NaOH
Cl₂ O₇ reacts with water and forms perchloric acid
Cl₂O₇ + H₂O \(\rightarrow\) 2HClO₄

Electron gain enthalpy becomes more negative down a period. CI > S >P. Electron gain enthalpy becomes less negative down a group. However, adding an electron to 2p orbital leads to greater repulsion due to high electron density than adding an electron to a larger 3p orbital. Hence, chlorine is more negative electron gain enthalpy than fluorine. Therefore, phosphorous has least electron gain enthalpy and chlorine has the most electron gain enthalpy.

The nuclear charge in Mg⁺² is more than Na+ and therefore electrons are drawn more closely, and hence size decreases

Metals have low ionisation energy and highly electropositive. As the ionisation energy of the metals decreases down the group, the electro positive character of elements increases. The basic character of hydroxides of alkaline earth metals increases down the group

The basic nature of the oxide decreases in a period while acidic nature increases. In a group, basic nature of the oxide increases, while acidic nature decreases

Generally metallic oxides are basic in nature white non metallic oxides are acidic in nature
Na₂O \(\Rightarrow\) basic
Cl₂O₇ \(\Rightarrow\) acidic

Metallic nature decreases in a period while non metallic character increases. Metallic nature increases in a group while non metallic character decreases