d = r_Na +r_F-
i.e r_Na + r_F- = 231 pm
We know that
\(\frac { { r }_{ Na }+ }{ { r }_{ F- } } =\frac { (Zeff)F- }{ (Zeff){ Na }^{ + } } \)
(Z_eff)_F- = Z -Z
= 9 - 4.15
=4.85
(Z_eff)Na⁺ = 11 - 4.15
=6.85
\(\frac { r{ Na }^{ + } }{ r{ F }^{ - } } =\frac { 4.85 }{ 6.85 } \)
\(\Rightarrow\) r_Na+ = 0.71 x r_F-
substituting (2) in (1)
(1) \(\Rightarrow\) 0.71 r_F- + r_F- = 231 pm
1.71 r_F- = 231 pm
\(r{ F }^{ - }=\frac { 231 }{ 1.71 } =135.1\) pm
Substituting the value of r_F- in equation (1)
r_Na+ 135.1 =231
r_Na+ = 95.9 pm

(i) Bond length (d_H-H) in hydrogen molecule = 0.74A
 \(\therefore\) Covalent radius of hydrogen = \(\frac { d_{ H-H } }{ 2 } -\frac { 0.74 }{ 2 } =.37A\)
(ii) Bond length (d_CI-Cl) in chlorine molecule = 1.98A
\(\therefore\) Covalent radius of chlorine = \(\frac { d_{ cl-CL } }{ 2 } =\frac { 1.98 }{ 2 } =0.99A\)
(iii) Bond length (C - C) in diamond = 1.54A
\(\therefore\)  Covalent radius of carbon = \(\frac { 1.54 }{ 2 } =0.77A\)
(iv) Bond length of (d_si-c) in carborundum = 1.93A
\({ d }_{ si-c }={ r }_{ si }+{ r }_{ c }\)
\(1.93\quad =\quad { r }_{ si }+0.77\)Refer (iii)
\({ r }_{ si }=1.93-0.77=1.16A\)
(v) Bond length (d_C-CI) in CCl₄ = 1.76A
\({ d }_{ c-cl }={ d }_{ c }+{ d }_{ cl }\)
1.76 \(=r_{ c }+0.99\) Refer (ii)
\({ r }_{ c }=1.76-0.99=0.77A\)

Factors influencing ionization enthalpy:
(i) Size of the atom:

• If the size of an atom is larger
• Outermost electron shell from the nucleus is also larger.
• The outermost electrons experience lesser nuclear force of attraction.
• It would be more and easier to remove an electron from the outermost shell.
• Thus, ionization energy decreases with increasing atomic sizes.

(ii) Magnitude of nuclear charge:

• As the nuclear charge increases, the force of attraction between the nucleus and valence electrons also increases.
• Consequently, more energy is required to remove a valence electron.

(iii) Screening or shielding effect of the inner electrons:
Ionization enthalpy decreases when the shielding effect of inner electrons increases. This is because when the inner electron shells increases, the attraction between the nucleus and the outermost electron decreases.
(iv) Penetrating power of subshells s, p, d & f:
The penetration power of the electrons in various orbitals decreases in a given shell in the order: s > p > d > f.
(v) Electronic configuration:
If an atom has half-filled or completely filled sub-levels, its ionization enthalpy is higher. This is because such atoms have extra stability and hence it is difficult to remove electrons from these stable configuration.

(i) ₄Be - 1s² 2s² ; ₅B - 1s² 2s² 2p¹

• The I.E. of Be is more than that of B though the nuclear charge of boron atom (Z = 5) is greater than that of beryllium atom (Z = 4). This can be explained as follows.
• Boron atom (Z = 5; 1s² 2s² 2P_x¹ 2p_y⁰ 2p_z⁰) is having one unpaired electron in the 2p-subshell. Be-atom (Z = 4; 1s² 2s²) is having paired electrons in the 2s-subshell.
• As the fully filled 2s-subshell in Be-atom is more stable than B-atom due to symmetry, more energy would be needed to remove an electron from Be-atom. Hence, Be has high I. E. Hence I.E of Be > B

(ii) ₇N - 1s² ,2s² ,2p_x¹, 2p_y¹, 2p_z¹
₈O -1s² ,2s² ,2p_x², 2p_y¹, 2p_z¹
∴ O has lower I.E. than N.

Z = 17
Z* = Z -S
= 17- [(0.35 \(\times\) No. of other electrons in nth shell) + (0.85\(\times\) No. of electrons in (n -1)^th shell) + (1.00 \(\times\) total number of electrons in the inner shells)]
= 17 - [(0.35 \(\times\) 6) + (0.85 \(\times\) 8) + (1 \(\times\) 2)]
= 17 - 10.9 = 6.1
Z* =  6.1

The electronic configuration of K atom is
K₁₉ = (1s²)(2s²p⁶)(3s²3p⁶)4s¹
Effective nuclear charge (Z*) = Z - S
Z* = 19 - [(0.85 \(\times\) No. of electrons in (n -1)t^h shell) + (1.00 total number of electrons in the inner shells)]
= 19-[0.85 \(\times\) (8) + (1.00 \(\times\) 10)]
Z* =  2.20

Electron gain Enthalpy	Electro negativity
It is the tendency of an isolated gaseous atom to attract an electron.	It is the tendency of an atom in a molecule to attract the shared pair of electrons
It is measured in electron volts/atom or kcal/mole or kj/mole.	It is a number and has no units
It is the property of an isolated atom.	It is property of a bonded atom.
An atom has an absolute  value of electron gain enthalpy	An atom has a relative value of electronegativity depending upon its bonding state. For example, sphybridized carbon is more  electronegative than sp2-hybridized carbon which, in turn, is more electronegative than sp,- hybridized carbon.
It does not change regularly in a period or group.	It changes regularly in a period or a group

Factors influencing electron gain enthalpy:
(i) Size of the atom : Electron affinity is inversely proportional to the size of the atom. As the size of atom increases, the effective nuclear charge decreases or the nuclear attraction for adding electron decreases.
(ii) Nuclear charge: Electron affinity is directly proportional to effective nuclear charge. Greater the nuclear charge more will be the tendency to accept electrons so E.A is more.
(iii) Electronic configuration: An atom with stable electronic configuration has no tendency to gain an electron. Such atoms have zero or almost zero electron gain enthalpy.
(iv) Shielding effect: Electron affinity is inversely proportional to shielding effect. Electronic energy state, lying between nucleus and outermost state hinder the nuclear attraction for incoming electron. Therefore, greater the number of inner lying states, less will be the electron affinity.

Chemical Properties	Metals	Non-Metals
Oxidising and Reducing Action	Generally good reducing agents	Generally good oxidising agents
Nature of oxides	Metallic oxides are basic in nature	Non-metallic oxides are acidic in nature
Nature of Hydrides	Form unstable hydride	Form stable hydrides
Action with acid	Dissolve in mineral acid to form salt.	Do not react with mineral acids
Arrangement of valence electrons	They have one, two or three valence electrons.	They have 4,5,6, 7 valence electrons

Characteristics of f-block elements:
(i) The elements in which the last electron enters in f-orbital are called f-block elements.
(ii) This block consists of 4f block elements and 5f block elements.
Lanthanoid series (4f) ₅₈Ce^- ₇₁Lu
Actinoid series (5f) ₉₀Th^- ₁₀₃Lr
(iii) Their general electronic configuration is (n - 2) f^1-14 (n - 1)d^0-1 ns².
Lanthanoid - 4f^1-14 5d^0-1 6s²
Actinoid - 5f^1-14 6d^0-1 7s²

Characteristics of p-block elements:
(i) General electronic configuration is ns²np^1-6 (Group 13 to 18).
(ii) They are generally non-metals but some metals and metalloids are present.

• Non-metallic character increases from left to right along period.
• Metallic character increases on moving down a group.

(iii) They have high electron affinity and ionization enthalpy.

• Both the properties increase along the period.
• Both the properties decrease on moving down the group.

(iv) Oxidising character.

• Increases on moving from left to right along a period.
• Reducing character increases from top to bottom in a group.

(v) s-block and p-block elements are collectively called representative elements.

Elements	Belonging to Groups	Electronic Configuration	Reason for their Name
s-block	1 and 2	ns1 and ns2	Valence electron enters the s-orbital
p-block	13 to 18	ns2np1 to ns2np6	Valence electron enters the p-orbital
d-block	3 to 12	(n-1)d1-10ns(0-2)	Valence electron enters the d-orbital
f-block	lanthanides and Actinoids	(n-2)f1-14(n-1)d0-1ns2	Valence electron enters the f-orbital

Anomalies of Mendeleev's periodic table:
(i) Some elements with similar properties were placed in different groups whereas some elements having dissimilar properties were placed in same group. Eg: Tellurium (127.6) was placed in VI group but iodine (127.0) was placed in VII group.
(ii) Some elements with higher atomic weights were placed before lower atomic masses in order to maintain the similar chemical nature of elements. This concept was called inverted pair of elements concept. Eg: \(_{ 27 }^{ 59 }{ Co }._{ 28 }^{ 58.7 }{ Ni }\) .
(iii) Isotopes did not find any place in Medeleev's periodic table.
(iv) Position of hydrogen could not be made clear.
(v) He did not leave any space for lanthanides and actinides which were discovered later on.
(vi) It could not explain the variable valencies of elements.
(vii) Elements with different nature were placed in one group as 'A' and 'B'; Eg: Alkali metals & coinage metals were placed together.
(viii) Diagonal and horizontal relationships were not explained.

(i) According to the recommendation of IUPAC, the groups are numbered from 1 to 181A.
(ii) There are 18 vertical columns which constitute 18 groups or families.
(iii) There are 7 horizontal rows of the periodic table known as periods.
(iv) The first period contains two elements. One present in first group and the other in 18th group.
(v) Second and third periods contain 8 elements in each.
(vi) Fourth and fifth periods are completely filled as they contain 18 elements in each.
(vii) The sixth period contains 32 elements. The seventh is incomplete. Fourteen elements of both sixth and seventh periods are placed in separate panels at the bottom of the table.
(viii) This periodic table is important and useful because we can predict the properties of any element using periodic trend.