The ionic radius of a cation is smaller than the parent atom because. the loss of one or more electrons increases the effective nuclear charge. As a result force of attraction of the nucleus for the electron increases as a result, the electron comes closer to the nucleus.
In contrast, the ionic radii of anion is always greater than the parent atom because addition of one or more electrons decreases the effective nuclear charge. As a result, force of attraction of the nucleus for the electrons decreases and ionic radii increase.

d_{H - Cl} = r_H + r_cl - .0.09 (r_cl - x_H)
1.28 = r_H+ 0.09 - 0.09 (3 - 2.1)
1.28 = r_H + 0.09 - 0.09 (0.9)
1.28 = r_H+ 0.09 - 0.081
1.28 = r_H + 0.909
\(\therefore\)  r_H = 1.28 - 0.909 = 0.317 A

The covalent radius of individual atom can also be calculated using the internuclear distance (d_A-B) between two different atoms A and B. The simplest method proposed by Schomaker and Stevenson is as follows.
d_A-B = r_A+r_B  - 0.09 (x_A - x_B)
where x_A and x_B are the electronegativities of A and B respectively in Pauling units.
Here x_A > x_B and radius is in A

(A) Atomic number 35:
1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁵
A receives the last electron in 4p orbital and hence belongs to 'p' block and 4th period.
Group number = 10 + 7 = 17
(B) Atomic number 18:
1s² 2s² 2p⁶ 3s² 3p⁶
Block - 'p' ; 3^rd period; (10 + 8) = 18th group
(C) Atomic number 28:
1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s²
Block - 'd' ; 4th period; (8 + 2) = 10th group.
(D) Atomic number 38:
1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁶5s²
Block - 's' ; 5th period; 2nd group

(i) The physical and chemical properties of the elements are correlated to the arrangement of electrons in their outermost shell (valence shell). Different elements having similar outer shell electronic configuration possess similar properties.
(ii) All the elements are arranged in the modem periodic table which contains 18v ertical columns and 7 horizontal rows. The vertical columns are called groups and the horizontal rows are called periods. Groups are numbered 1 to 18 in accordance with the IUPAC recommendation which replaces the old numbering scheme IA to VIlA, IB to VIIB and VIII.
(iii) Each period starts with the element having general outer electronic configuration ns¹ and ends with np⁶. Here 'n' corresponds to the period number (principal quantum number).

(i) Cu is largest due to less effective nuclear charge. It has 29 electrons, 29 protons. Cu⁺ has 28 electrons and 29 protons, Cu²⁺ has 27 electrons and 29 protons.
(ii) He (helium) has highest ionisation energy.
(iii) Mg is more metallic due to lower ionisation energy.

The energy required to remove the most loosely held electron from an isolated gaseous atom is called ionization energy.
(i.e) Atom_(g) + Energy ⟶ Positive Ion_(g)+Electron
Li_(g) + 520Kj mol^-1 ⟶ Li_(g)⁺ + e^-

Reason: The ionization energy decreases with the increasing size of atom. This is due to the fact, in larger atoms electrons are loosely bound to the nucleus, i.e., lesser energy is required for removal of electrons from larger atoms than the smaller one. Hence ionization energy is lower for larger atoms.

As we move down a group, the number of inner shell electrons increases. which in turn increases the repulsive force exerted by them on the valence electrons, thus decreases the attractive force acting on the valence electrons by the nucleus. Therefore the ionisation energy decreases.

(i) CI- 1s² 2s² 2p⁶ 3s² 3p⁵; Cl^- - 1s² 2s² 2p⁶ 3s² 3p⁶
(ii) Generally anions are larger than their parent atom.
(iii) CI^- had 18 electrons and CI has 17 electrons. The later experiences a greater effective inward pull by the nucleus.
(iv) This decreases the size of CI atom when compared to Cl^- ion.

The largest species is Mg
The smallest species is Al³⁺
Reason:
(i) Atomic radii decreases across a period.
(ii) Cations are smaller than their parent atoms.
(iii) Among the isoelectronic ions, the greater the charge on the cation smaller will be the radius.

(i) Noble gases have completely filled electronic configuration and are stable.
(ii) It is difficult to remove the electrons from the valence shell. So the ionisation energy is maximum.

Na - 1s² 2s² 2p⁶ 3s¹
Mg - 1s² 2s² 2p⁶ 3s²
Second I-E of Na is higher because it acquires nearest noble gas configuration after losing one electron and 2nd electron is to be removed from the noble gas configuration.

The term periodicity of properties indicates that the elements with similar properties reappear at certain regular intervals of atomic number in the periodic table.

Na⁺ - 10 electrons 1s² 2s² 2p⁶
Cl^- - 18 electrons 1s² 2s² 2p⁶ 3s² 3p⁶
In case of Na⁺ the nuclear charge is more. So the electrons are strongly attracted by the nucleus, nuclear charge so the size decreases.
In CI^- the number of orbital electrons becomes greater, nuclear charge decreases and hence the electrons are not strongly attracted by the nucleus so ionic radius increases.

A cation is formed by loss of one or more electrons. The nuclear charge remains the same but the number of electrons becomes less than the parent resulting in the increase in the effective nuclear charge per electron. This causes decrease in size.

Covalent radius is one half of the internuclear distance between two identical atoms linked together by a single covalent bond.

Metallic radius is defined as one half of the distance between two adjacent metal atoms in the closely packed metallic crystal lattice.

\({ O }_{ (g) }+{ e }^{ - }\longrightarrow { O }^{ - }_{ (g) };{ \Delta }_{ eg }H=-141\quad kj\quad { mol }^{ -1 }\)
\({ O }_{ (g) }+{ e }^{ - }\longrightarrow { O }^{ 2- }_{ (g) };{ \Delta }_{ eg }H=+780\quad kj\quad { mol }^{ -1 }\)
When an electron is added to oxygen atom to form ion, energy is released. Hence, first electron gain enthalpy of oxygen is negative.
But when another electron is added to O^- ion to form O^2- ion, it feels stronger electrostatic repulsion. Hence, addition of second electron takes place with absorption of energy. That's why the second electron gain enthalpy of oxygen is positive.

Halogens act as good oxidising agents.
Due to high electron gain enthalpy they act as good oxidising agents as they can gain electrons easily

The increasing order of non-metallic character as follow
Non metallic character increases across the period and decreases down the group.

On moving down the group, electronegativity decreases because atomic size increases. Iodine is the largest, therefore it is least electronegative.

(i) F-1s² 2s² 2p_x² 2p_y² 2p_z¹
Cl- -1s² 2s² 2p⁶ 3s² 3p_x² 3p_y² 3p_z¹
(ii) Because of small size of fluorine, the 2p subshell becomes compact.
(iii) There occurs electron-electron repulsion
(iv) Due to its small size, large crowding of electrons occurs which reduces the nuclear charge. So the electron affinity of fluorine is less than chlorine.

(i) Noble gases have completely filled electronic configurations.
(ii) It has no tendency to attract the electrons towards itself
(iii) so the electron affinity of noble gases are almost zero.

(i) Fluorine, F (Z = 9)
(ii) Potassium, K (Z = 19)
(iii) Strontium, Sr (Z = 38)
(iv) Ruthenium, Ru (Z = 44)
(v) Praseodymium, Pr (Z = 59)
(vi) Xenon, Xe (Z = 54)

Zn, Cd and Hg have completely filled d orbital and so. They do not show the properties of transition elements. ∴ They are not regarded as transition elements. On the basis of properties, all transition elements are d-block elements. On the basis of electronic configuration, all d-block elements are not transition elements.

(i) In a period, all elements have equal number of shells.
(ii) In a group, all elements have equal number of electrons in the outermost shell.

The elements in which extra electron enters (n-2) f-orbitals are called f-block elements. These elements form a transition series within the transition elements. So these elements are called inner transition elements.

118, Uuo.

Transition metals form a bridge between the chemically active metals of s-block and the less active elements of groups of 13th and 14 and thus are called as Transition elements.

Noble gases having completely filled electronic configuration so neither accept nor lose their electron readily and hence they are chemically inert in nature.

(i) ₁₃Al =1s², 2s², 2p⁶, 3s², 3p¹
[Third period and p-block]
(ii) ₂₄Cr = 1s², 2s², 2p⁶, 3s², 3p⁶, 4s¹, 3d⁵
[Fourth period and d-block]
(iii) ₂₉Cu = 1s², 2s², 2p⁶, 3s², 3p⁶, 4s¹, 3d¹⁰
[Fourth period and d-block]
(iv) ₁₁Na = 1s², 2s², 2p⁶, 3s¹
[Third period and s-block]

The element belongs to d- block, Zinc (Z=30).

Metallic character increases down a group and decreases along a period as we move from left to right.Hence the order of increasing metallic character is, P < Si < Be < Mg < Na.

The element with Z = 117, would belong to the halogen family (group 17) and the electronic configuration 89 would be. [Rn] 4f¹⁴ 5d¹⁰ 7s² 7p⁵ . the element with Z = 120, will be placed in group 2 (alkaline earth metals), and will have the electronic configuration [Uuo]8s² .

The first period n= 1, it is the first energy level and has only 1s orbital which can accommodate a maximum of 2 electrons hence the number of elements in the first period is only 2.

Elements that show properties that are characteristic of both metals and non-metals are called as Semi Metals or Metalloids. Example- germanium, silicon, arsenic, antimony and tellurium.

s-block and p-block (Groups 13-18) elements are collectively called representative elements or Main Group Elements.

Gallium and Bromine.

In the present set up of the long form of the periodic table, we have eighteen groups, seven periods (i.e. principal quantum number, n = 7) and four blocks (s, p, d and f -block elements).
Therefore, the maximum number of elements which can be accommodated in the present set up of the long form of the periodic table in accordance with Aufbau principle is
1s2, 2s², 2p⁶, 3s², 3p⁶, 4s², 3d¹⁰, 4p⁶, 5s², 4d¹⁰,
5p⁶, 6s², 4f¹⁴, 5d¹⁰, 6p⁶, 7s², 5f¹⁴, 6d¹⁰, 7p⁶ = 118.

Electronic configuration of the element is 3d¹ 4S².
The element belongs to fourth period and group 3.
The element is scandium (Z = 21).

Third period contains eight elements in it. The third period begins with the electrons entering the third energy shell, i.e n = 3. Out of nine orbitals of this energy level [one 3s, three 3p (3px, 3py, 3pz) and five 3d (3dxy, 3dyz, 3dxz, 3dx² - y², dz²)], the five 3d orbitals have higher energy than 4s orbital. As such only four orbitals (one 3s and three 3p) are filled before the fourth energy level begins to be filled. Hence, there are only eight elements in the third period.

The elements beyond fermium are called as transfermium elements.

In periodic table the elements placed beyond uranium (Z = 92) are all synthetic elements and are called as transuranium elements.

(i) Position of hydrogen is not defined till now.
(ii) Lanthanides and actinides still find place in the bottom of the table.

The Mendeleev periodic law states that, "The physical and chemical properties of the elements are periodic functions of their atomic weights.

(i) The number of electrons increases by the same number as the increase in the atomic number.
(ii) As the number of electrons increases, the electronic structure of the atom changes.
(iii) Electrons in the outermost shell of an atom (valence electrons) determine the chemical properties of the elements.

(i) The cause of periodicity in properties is the repetition of similar outer electronic configuration after certain regular intervals.
(ii) Eg: Consider elements of group 1s
₃Li - 1s² 2s¹
₁₁Na -1s² 2s²2p⁶ 3s¹
-Due to similar outermost electronic configuration they exhibit similar properties.

Newlands law of Octaves:
When elements are arranged in increasing order of their atomic weights every eighth element resembles its properties with the first one just like the eight note of a musical table.

Limitations: Failed for heavier elements beyond calcium.

(i) Lothar Meyer plotted the physical properties such as atomic volume, melting point and boiling point against atomic weight and obtained a periodically repeated pattern.
(ii) Lothar Meyer observed a change in length of that repeating pattern.
(iii) In 1868, Lothar Meyer had developed a table of the elements that closely resembles the modern periodic table.

Helium is placed in group 18 since it has a completely filled valence shell and so resembles characteristics of noble gases.

Lanthanoids and actinoids are placed in separate rows at the bottom of the periodic table to maintain its structure and to preserve the principle of classification by keeping elements with similar properties in a single column.

Selenium (Se)
Xenon (Xe)

(i) Sodium (Z = 11)
(ii) krypton (Z = 36)
(iii) Bromine (Z = 35)
(iv) Scandium (Z = 21)