(i) In the 1^st group, lithium differs in many aspects from its own family elements. Similarly, in the 2^nd group,  beryllium differs in many aspects from its own family.
(ii) For example, lithium forms compounds with more covalent character. But other alkali metals of this group form only ionic compounds.
(iii) Similarly, beryllium forms compounds with more covalent character while the other elements of this family form only ionic compounds.
(iv) Lithium and beryllium resemble more with the elements lying at their right hand side in the 3^rd period than with the other members of their own family.
(v) These kinds of anomalies are also observed from 13^th to 17^th groups.
(vi) This sort of similarity is commonly referred to as diagonal relationship in the periodic properties.
(vii) The anomalous behaviors are attributed to the following factors:
(a) Smaller atomic size
(b) Higher ionization enthalpy
(c) High electronegativity

(a) The energy required to remove the most loosely held electron from an isolated gaseous atom is called as ionization energy.
(b) (i) Variation in a period:
On moving across a period from left to right, the ionization enthalpy value increases. This is due to the following reasons:
1. Increase of nuclear charge in a period
2. Decrease of atomic size in a period
Because of these reasons, the valence electrons are held more tightly by the nucleus, thus ionization enthalpy increases.
Hence, ionization energy is a periodic property.
(ii) Variation in a group:
As we move from top to bottom along a group, the ionization enthalpy decreases. This is due to the following reasons:
1. A gradual increase in atomic size.
2. Increase of screening effect on the outermost electrons due to the increase of number of inner electrons.
Hence, ionization enthalpy is a periodic property.

Factors influencing ionization enthalpy:
(i) Size of the atom:
If the size of an atom is larger, the outermost electron shell from the nucleus is also larger and hence the outermost electrons experience lesser force of attraction. Hence it would be more easy to remove an electron from the outermost shell. Thus, ionization energy decreases with increasing atomic sizes.
Ionization enthalpy \(\infty{1\over Atomic \ size}\)
(ii) Magnitude of nuclear charge:
As the nuclear charge increases, the force of attraction between the nucleus and valence electrons also increases. So, more energy is required to remove a valence electron. Hence I.E increases with increase in nuclear charge.
Ionization enthalpy \(\alpha \ nuclear \ charge\)
(iii) Screening or shielding effect of the inner electrons:
The electrons of inner shells form a cloud of negative charge and this shields the outer electron from the nucleus. This screen reduces the coulombic attraction between the positive nucleus and the negative outer electrons. If screening effect increases, ionization energy decreases.
Ionization enthalpy  \(\infty{1\over Screening\ effects}\)
(iv) Penetrating power of subshells s, p, d, and f:
The s-orbital penetrate more closely to the nucleus as compared to p-orbitals. Thus, electrons in s-orbitals are more tightly held by the nucleus than electrons in p-orbitals. Due to this, more energy is required to remove a electron from an s-orbital as compared to a p-orbital. For the same value of 'n', the penetration power decreases in a given shell in the order.
s>p>d>f.
(v) Electronic configuration:
If the atoms of elements have either completely filled or exactly half filled electronic configuration, then the ionization energy increases.

Atomic radius is the distance between the centre of its nucleus and the outermost shell containing the electron. Atomic radius is a periodic property.
(i) Variation in periods:
The atomic radius decreases while going from left to right in a period, As we move from left to right in a period, the nuclear charge increases by one unit in each succeeding element. But the number of the shell remains same. Hence, the electrons are attracted strongly by the nucleus. Hence the atomic radius decreases along the period.
In 2^nd period r_Li > r_Be > r_B > r_C > r_N> r_O > r_F
(ii) Variation in a group:
The atomic radius of elements increases with increase in atomic number as we move from top to bottom in a group. The attraction of the nucleus for the electrons decreases as shell number increases. Hence atomic radius increases along the group.
In 1^st group r_Li < r_Na < r_K < r_Rb < r_Cs
Hence, atomic radii is a periodic property.

(a) Atomic radius is the distance between the centre of its nucleus and the outermost shell containing the electron.
(b) Difficulties in determining atomic radius
(i) The size of an atom is very small (~1.2\(\mathring A\) i.e 1.2 x 10 ^-10)
(ii) The atom is not a rigid sphere; it is more like a spherical cotton ball rather than like a cricket ball.
(iii) It is not possible to isolate an atom and measure its radius.
(iv) The size of an atom depends upon the type of atoms in its neighborhood and also the nature of bonding between them.

Based of chemical behavior:
(i) Main group elements:
All s-block and p-block elements excluding 18^th group elements are called representative elements.
(ii) Noble gases:
The 18^th group elements are exclusively called noble gases. They have completely filled electronic configuration as ns² np⁶, These elements are highly stable.
(iii) Transition elements:
The elements of d-block are called transition elements. These include elements of groups from yd to 12^th lying between s-block and p-block elements.
(iv) Inner transition elements.
The elements of f-block are called inner-transition elements. These consist of lanthanides and actinides, with 14 elements in each. Based of physical properties:
(i) Metals:
Metals comprise more than 78% of all known elements. They are usually solids at room temperature (except Hg, Ga and Cs). They have high melting and boiling points. They are good conductors of heat and electricity.
(ii) Non-metals:
Non-metals are usually solids or liquids or gases at room temperature with low melting and low boiling points (except boron and carbon). They are poor conductors of heat and electricity. Most of the non-metallic solids are brittle and are neither malleable nor ductile.
(iii) Metalloids or Semi-metals:
Some elements in the periodic tables show properties that are characteristic of both metals and non-metals, They are called metalloids. Example: Silicon, germanium, arsenic, antimony and tellurium.

(i) Number of electrons in outermost shell: The number of electrons present in the outermost shells does not change on moving down in a group, i.e remains the same. Hence, the valency also remains same within a group.
(ii) Number of shells: In going down a group the number of shells increases by one at each step and ultimately becomes equal to the period number to which the element belongs.
(iii) Valency: The valencies of all the elements of the same group are the same. The valency of an element with respect to oxygen is same in a group.
(iv) Metallic character: The metallic character of the elements increases in moving from top to bottom in a group.

(i) Number of electrons in outermost shell: The number of electrons present in the outermost shell increases from 1 to 8 as we proceed in a period.
(ii) Number of shells: As we move from left to right in a period the shells remains the same. The number of shells present in the elements corresponds to the period number. For example all the elements of 2^nd period have on 2 shells (K, L)
(iii) Valency: The valency of the elements increases from left to right in a period. With respect to hydrogen, the valency of period elements increases from 1 to 4 and then falls to one. With respect to oxygen, the valency increases from 1 to 7.
(iv) Metallic character: The metallic character of the elements decreases across a period.
For example: 3rd period  

Merits of Moseley's long form of periodic table:
(i) As this classification is based on atomic number, it relates the position of an element to its electronic configuration.
(ii) The elements having similar electronic configuration fall in a group. They also have similar physical and chemical properties.
(iii) The completion of each period is more logical. In a period as the atomic number increases, the energy shells are gradually filled up until an inert gas configuration is reached.
(iv) The position of zero group is also justified in the table as group 18.
(v) The table completely separates metals and non-metals.
(vi} The table separates two subgroups, lanthanides and actinides, dissimilar elements do not fall together.
(vii) The greatest advantage of this periodic table is that this can be divided into four blocks namely s, p, d and f-block elements.
(viii) This arrangement of elements is easier to remember, understand and reproduce.

(i) The long form of periodic table of the elements is constructed on the basis of modem periodic law. The arrangement resulted in repeating electronic configurations of atoms at regular intervals.
(ii) The elements placed in horizontal rows are called periods and in vertical columns are called groups.
(iii) According to IUPAC, the groups are numbered from 1 to 18.
(iv) There are 18 vertical columns which constitute 18 groups or families. All the members of a particular group have similar outer shell electronic configuration.
(v) There are 7 horizontal rows called periods.

The elements are shown in the above table along with its atomic number.
(vi) The atomic number also indicates the number of electrons in the atoms of an element.
(vii) This periodic table is important and useful because we can predict the properties of any element using periodic trend, even though that element may be unfamiliar to us.

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, but iodine (127) was placed in VII group.
Example: Tellurium (127.6) was placed in VI 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.
Example: ⁵⁹₂₇Co and ^58.7 ₂₈Ni
(iii) Isotopes did not find any place in Mendeleev'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) Elements with different nature were placed in one group, Example: Alkali metals and coinage metals were placed together:
(vii) Diagonal and horizontal relationships were not explained.

(a) Mendeleev's periodic law:
Mendeleev's periodic law states that the physical and chemical properties of elements are a periodic function of their atomic weights.
(b) Merits of Mendeleev's periodic table:
(i) The comparative studies of elements were made easier.
(ii) The table shows the relationship'in properties of elements in a group.
(iii) The table helped to correct the atomic masses of some elements later on. At the time of Mendeleev, the atomic weight of Au and Pt were known as 196.2 and 196.7 respectively. However, Mendeleev placed Au (196.2) after Pt (196.7) saying that atomic weight of Au is incorrect, which was later on found to be 197.
(iv) At the time of Mendeleev, about 70 elements were known and thus blank spaces were left for unknown elements which helped further discoveries.
(v) Both Gallium (Ga) in III group and Germanium. (Ge) in IV group, were unknown at that time by Mendeleev predicted their existence and properties. He referred the predicted elements as eka-aluminium and eka-silicon. After discovery of the actual elements, their properties were found to match closely to those predicted by Mendeleev.

Electronic Configuration of Aluminium
\(\underbrace { { Al }^{ 13 }{ 1s }^{ 2 } }_{ (n-2) } \underbrace { 2s^{ 2 }{ 2p }^{ 2 } }_{ (n-1) } \underbrace { { 3s }^{ 2 }{ 3p }^{ 1 } }_{ n } \)

∴ Effective nuclear charge = Z - S = 13 - 9.5
(Z_eff)_Al =3.5
Electronic Configuration of chlorine
\(\underbrace { { 1s }^{ 2 } }_{ (n-2) } \underbrace { 2s^{ 2 }{ 2p }^{ 2 } }_{ (n-1) } \underbrace { { 3s }^{ 2 }{ 3p }^{ 5 } }_{ n } \)

∴ Effective nuclear charge = Z - S = 17 - 10.9
(Z_eff)_cl = 6.1
(Z_eff)_cl > (Z_eff)_Al and hence r_clAl